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PREFACE

Bismillahirrahmanirrahim. Assalamu’alaikum Wr. Wb. Alhamdulillahirabbil’alamin, praise and our thanks gives presence of God who is praised and most high. that has gave us enjoy islam, enjoy faith and enjoy healthy that no predicted. Shalawat and greeting not forgets us submit to our big prophet, Mohammad saw that has brought us from stupidity era to science era. We render thanks to the whole party that has ed and helped us until we can finish this paper. Especially to both our old fellow, our lecturer Drs. Agung Purwanto, M.Si., friends and other parties that can not we mention one by one. “Nothing are perfect except God.”. Maybe that words that can depict this paper. Until we open our liver door to accept criticism and suggestion from all parties, in order to later, in paper hereinafter can be better next. And we apologize if existed insuffiency and mistake in this paper making. Because all that correctness is only God property, and all wrong ones is ours person. Akhirul kalam. Wassalamu’alaikum Wr. Wb.

Jakarta, October 2010 Chemistry 2009

COMPILERS Atomic Theory and Elements Periodic System 1. Ratna Purnama Sari 2. Sri Astriani 3. Endah Dianty Chemica Bonding and Molecular Geometry 1. Ulfah Choiriyah 2. Sandra Masduroh 3. Aditya Agam Hybridization 1. Fitri Hartanti 2. Sarah Yane Irene 3. Selline Yansu Groups IA and IB 1. Reni Andriani 2. Putriningtias Imansari Groups IIA and IIB 1. Sifa Fauziah Dwidara 2. Ranie Pujiastuti Groups IIIA and IIIB 1. Indah Budiarti 2. Fenty Kanisateja 3. Nurain Tri Rahayu Groups IVA and IVB 1. Fitri Hartanti 2. Sarah Yane Irene

3. Selline Yansu Groups VA and VB 1. Rully Putera S 2. Lilis Septiarini 3. Siti Lili A Groups VIA and VIB 1. Rizqi Meidani F 2. Mita Rahayu Groups VIIA and VIIB 1. Mega Puspita S 2. Christiandi Cahyo 3. Syarifah Nabilah F Groups VIIIA and VIIIB 1. Yulinar Eka P 2. Apriyanti Lantanide and Actinide 1. Taufik Triadi 2. Zenith Tarra F 3. Kartika Diah A

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LIST OF CONTENTS

Preface...................................................................................................... Compiler..................................................................................................... List of Contents.......................................................................................... Contents Atomic Theory and Elements Periodic System..................................1 Chemical Bonidng and Geometry Molecular...................................42 Hybridization...................................................................................80 Groups IA and IB...........................................................................107 Groups IIA and IIB.........................................................................147 Groups IIIA and IIIB.......................................................................167 Groups IVA and IVB.......................................................................207 Groups VA and VB.........................................................................245 Groups VIA and VIB............................................................................. Groups VIIA and VIIB........................................................................... Groups VIIIA and VIIIB......................................................................... Groups Lantanide and Actinide...........................................................

1. CHAPTER 1 SOME FUNDAMENTAL CONCEPTS Chemistry is the science that deals with the composition and properties of substance and the transformations they undergo. DIVISIONS OF CHEMISTRY General Chemistry is a broad survey of chemistry as a whole, with special emphasis on its basic principles and laws. It includes the properties and reactions of some of the most common elements and compounds. Organic Chemistry is the study of compounds of carbon, either as they are produced in plants and animals or as they are formed synthetically. Biochemistry is the study of the chemistry of living processes. All the chemical reactions taking place in the body are more specifically referred to as physiological chemistry. Analytical chemistry is concerned with the methods of determining the various constituents of matter as to what they are (qualitative analysis), or how much they are (quantitative analysis). Physical chemistry deals with the principles and laws that underlie chemical changes. Nuclear chemistry is the study of changes that take place in the nucleus of the atom. 1.1 Matter Matter is anything that occupies space and has mass. In ordinary chemical reactions, matter can neither be created nor destroyed (law of conservation of matter). In nuclear reaction, however, matter can be converted into energy, or vice versa.

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Physical States of Matter, matter exists in three physical states: solid, liquid, and gaseous, depending on temperature and pressure. Solids are rigid and have a definite volume and a definite form. Liquids have a definite volume but no definite form. They flow and assume the shape of the vessel which holds them. Gases have neither a definite volume nor a definite form. They diffuse into every part of the container in which they are placed. Classification of Matter. Matter can be in the form of an element, a compound, or mixture. Elements. An element is a substance that cannot be decomposed into simpler substance by ordinary chemical means. It may also be defined as a substance whose properties give it a definite place in the periodic table. There are 103 known chemical elements at the present time. They may be classified into metals and nonmetals. Examples of metals are iron, silver, and gold. Sulfur, oxygen and nitrogen are nonmetals. Compounds. A compounds is made up of two or more elements chemically combine in definite proportions by weight. Thus, the compound water is composed of 11.11 percent hydrogen and 88.89 percent oxygen by weight. A compound is homogeneous. Its properties are quite different from those of its constituent elements, and its constituent elements can be separated to oxygen in water illustrates the law of definite proportions, or the law of definite composition.

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Mixture. A mixture is made up of two or more substance that are not combined chemically. Its component parts retain their own properties and can be separated by mechanical means. For examples, cream of tartar baking powder is a mixture of sodium bicarbonate, cream of tartar, and starch. Changes in Matter. Matter undergoes changes, some of which are physical and others of which are chemical. Physical changes. A physical change is an alteration in the condition or state of substance. The chemical composition of the substance is not changed. Examples of physical changes are the chopping of wood, the breaking of glass, and the melting of ice. Chemical Changes. A chemical change is one in which a new substance is formed having a composition and properties different from those of the original substance. For example, iron on exposure to moist air becomes rust; sulfur, on burning, becomes sulfur dioxide. Properties of Matter. The distinguishing characteristics of a substance are referred to as its properties. There are two types of properties: physical chemical. Physical Properties. The physical properties of a substance are those associated with physical changes. They include characteristics such as color, odor, taste, density, crystalline form, boiling point, and melting point.

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Chemical Properties. Chemical properties are characteristics of elements and compounds which describe the manner in which these substance react with other substance. For example, sodium reacts readily with water to liberate hydrogen; water is decomposed into hydrogen and oxygen by electrolysis. 1.2 Energy Energy may be defined as the ability to do work. Matter always possesses energy in one form or another. All transformations of matter are accompanied by transformations of energy. Forms of Energy and Transformations. Energy can take many forms; heat, light, electrical, kinetic, chemical, and nuclear or atomic energy. Energy can be changed from one form to another but cannot be created or destroyed in reactions other than nuclear reactions (law of conservation of energy). For example, electricity is changed into light in a light bulb; heat from steam is changed into electricity in an electric generator.

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CHAPTER 2 ATOMIC THEORY Birth of Atomic Theory Modern chemistry is based on atomic theory. To understand the atomic theory, first you must learn the fundamental laws including the law of conservation of mass, comparative law remains, and the law of multiple comparisons. These laws are basic atomic theory and at the same time represent the conclusions drawn from the theory of atoms. However, the atomic theory itself is incomplete. Chemistry can be a consistent system since the atomic theory combined with the concept of molecule. In the past, the existence of an atom is only a hypothesis. In the early 20th century atomic theory finally proved. It also became clear that the atom consists of particles smaller. Current atomic theory is slowly growing in line with this development and become a skeleton material world. 2.1 Birth of chemical

Modern chemistry initiated by the French chemist Antoine Laurent Lavoisier (1743-1794). He discovered the law of conservation of mass in chemical reactions, and reveal the role of oxygen in combustion. Based on this principle, the chemistry developed in the right direction. Modern chemistry initiated by the French chemist Antoine Laurent Lavoisier (1743-1794). He discovered the law of conservation of mass in chemical reactions, and reveal the role of oxygen in combustion. Based on this principle, the chemistry developed in the right direction.

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Actually, oxygen is found independently by two chemists, the British chemist Joseph Priestley (1733-1804) and Swedish chemist Carl Wilhelm Scheele (1742-1786), at the end of the 18th century. Thus, only about two hundred years before the birth of modern chemistry. Thus, chemistry is a relatively young science compared to physics and mathematics, both have developed several thousand years. However, alchemy, metallurgy and pharmacy in ancient times can be considered as the root chemistry.

Many discoveries found by people

actively involved in all these areas contribute enormously to modern chemistry, although alchemy was based on an incorrect theory. Furthermore, prior to the 18th century, metallurgical and pharmaceutical industries are actually based on experience rather than theory. So, it seems unlikely this early points which later evolved into modern chemistry. Based on these things and the nature of modern chemistry is well organized and systematic methodology, real roots of modern chemistry may be found in ancient Greek philosophy. The road from the ancient Greek philosophy to modern atomic theory is not always smooth. In ancient Greece, there is a sharp dispute between the atomic theory and the denial of the existence of atoms. Actually, the atomic theory remains the world's unorthodox chemistry and science. Educated people are not interested in atomic theory until the 18th century. At the beginning of the 19th century, the British chemist John Dalton (1766-1844) birth anniversary of ancient Greek atomic theory. Even after his birth again, not all scientists accept the theory of atoms. Not until the early 20th century theory of ato, finally proven as fact, not just hypothetical. This is achieved by a skilled trial by the French chemist Jean

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Baptiste Perrin (1870-1942). So, it took long enough to establish the basis of modern chemistry. 2.1.1 Ancient atomic theory

As noted earlier, the roots of modern chemistry is the atomic theory developed by the ancient Greek philosopher. Atomic philosophy of ancient Greece is often associated with Democritos (roughly 460BC-roughly 370 BC). However, no posts Democritos living. Therefore, we must be the source of the long poem "De Rerum natura" written by the artists of the Roman Lucretius (about 1996 BC-about 55 BC). Presented by Lucretius atom has similarities with modern molecules. Grapes (wine) and olive oil, for example, has its own atoms. Atom is an abstract entity.

Atom has a distinctive shape with the corresponding

function with form. "Atom wine was round and smooth so it can through the esophagus with a smooth, while the atoms of quinine rough and it would be difficult through the esophagus." Molecule of modern structural theory states that there is a very close relationship between molecular structure and function. Although the philosophy articulated by Lucretius is not ed by evidence obtained from the experiment, this is the beginning of modern chemistry. In the long period from ancient times until the Middle Ages, the theory of the atom In heretikal (berlwanan with commonly accepted theory) because the theory of four elements (water, soil, air and fire) the proposed Aristotole ancient Greek philosopher (384 BC-322 BC) charge of.

When

otortas Aristotle began to decline in early modern centuries, many

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philosophers and scientists began to develop theories that influenced the Greek atomic theory.

Preview the material retained by the French

philosopher Rene Descartes (1596-1650), German philosopher Gottfried Wilhelm Freiherr von Leibniz (1646-1716), and British scientist Sir Isaac Newton (1642-1727) more or less influenced by the theory of atoms. 2.1.2 Dalton's atomic theory

At the beginning of the 19th century, atomic theory as a philosophical matter has been well developed by Dalton who developed the atomic theory based on the role of atoms in chemical reactions. The theory of atomic summarized as follows: i.

elementary particles that make up the elements are atoms. All the atoms of certain elements are identical.

ii.

the atomic mass of the same type will be identical but different from other types of elements atomic mass.

iii.

all atoms involved in chemical reactions.

The whole atom

would form a compound. The type and number of atoms in certain compounds remain.

The theoretical bases Dalton theory is primarily based on the law of conservation of mass and tetap1 comparative law, both have been found previously, and comparative law berganda2 developed by Dalton himself. A particular compound always contain the same element mass ratio.

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Democritos atoms can be regarded as a kind of miniature material. So the number of types of atoms will be equal to the amount of material. On the other hand, Dalton's atoms are constituents of matter, and many compounds can be formed by a limited number of atoms. So, there will be a limited number of types of atoms.

Dalton's atomic theory requires a

process of two or more atoms combine to form the material. This is the reason why atomic chemistry Dalton called atoms. Evidence of the existence of atoms When

Dalton

proposed

the atomic

theory,

theory

attracted

considerable attention. However, this theory fails to receive full . Some ers of Dalton made important efforts to persuade against this theory, but some opposition persists. Chemistry was not enough to prove the existence of atoms with the experiment. So the atomic theory remains a hypothesis.

Furthermore, science after the 18th century developed a

variety of experiments that make many scientists became skeptical of the atomic hypothesis. For example, such famous chemist Sir Humphry Davy (1778-1829) and Michael Faraday (1791-1867), both from England, both doubt on the theory of atoms. While atomic theory remains a hypothesis, various dibuta great progress in various fields of science. One of them is the rapid emergence of thermodynamics in the 19th century. Structural Chemistry at that time represented by the atomic theory is only a matter of academic with little possibility of practical application. But the thermodynamics derived from practical issues such as efficiency of steam engines seem more important. There was a sharp controversy between those ing atomic thermodynamics. The debate between Austrian physicist Ludwig Boltzmann

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(1844-1906) and the German chemist Friedrich Wilhelm Ostwald (1853-1932) with the Austrian physicist Ernst Mach (1838-1916) deserves note. This debate is bad, Boltzmann committed suicide. In the early 20th century, there were major changes in the interest of science.

A series of important discoveries, including radioactivity,

causing interest in the properties of atoms, and more generally, structural science. That there are atoms in the experiment was confirmed by sedimentation equilibrium experiments by Perrin. English botanist, Robert Brown (1773-1858) discovered takberaturan motion of colloidal particles and the movement called the motion Brown, in his honor. Swiss physicist Albert Einstein (1879-1955) developed a theory based on the theory of atomic motion. According to this theory, Brownian motion can be expressed by an equation that contains Avogadro's number. D = (RT / N). (1/6παη) ... (1.1) D is the motion of particles, R gas constant, T temperature, N Avogadro's number, α particle radius and η the viscosity of the solution. The core idea is as follows Perrin. Colloidal particles move randomly by Brownian motion and simultaneously settle down by the influence of gravity.

Equilibrium sedimentation equilibrium generated by these two

motion, random motion and sedimentation. Perrin carefully observe the distribution of colloidal particles, and with the help of equation 1.1 and its data, he got the Avogadro's number.

Surprising value that matches the

Avogadro's obtainment obtained by other methods are different.

This

agreement further proves the truth of atomic theory that became the basis of the theory of Brownian motion.

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Not to mention, Perrin can not observe atoms directly. What can be done at that time scientists, including Perrin, is to demonstrate that Avogadro's number obtained from a number of different methods based on the theory of identical atoms.

In other words they prove the theory of

atoms indirectly with logical consistency. Within the framework of modern chemistry, such methodology is still important. Even to this day still can not directly observe atoms as small particles with the naked eye or microscope optics. To observe directly with visible light, particle size must be larger than the wavelength of visible light. The wavelengths of visible light is in the range of 4.0 x 10 -7

-7

- 7.0 x 10

m, which amount 1,000 times bigger than the size of atoms. So clearly

outside the range of optical equipment to observe the atom. With the help of new tools such as electron microscopy (EM) or scanning tunneling microscope (STM), this impossibility can be overcome. Although the principle of observing the atom with this tool, in contrast to what is involved with observing moon or interest, we can say that we are now able to observe atoms directly. 2.1.3 JJ Thomson’s atomic theory

In early 1900 , J.J. Thomson proposed a new atomic model that include the presence of particles of electrons and protons . Since the experiments show that protons have a mass far greater than the electron , Thomson described the model atoms as a single large proton . In the particles of protons, electrons entering Thomson neutralize the positive charge of protons . According to Thomson, the atom consists of a sphere of positive charge is evenly distributed with the meeting charge . On the

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positive charge is spread electron with a negative charge , which represents the positive charge . A popular way to describe this model is to consider electrons as raisins ( plumb ) in the proton pudding cake , so this model is given a raisin cake model ( plumb - pudding model).

Although Thomson's atomic model is the first which incorporates the concept of the existence of protons and electrons are charged , Thomson's model was not able to through the observation of subsequent experiments . For the record, the protons used in the Thomson model of the proton beam is not found in more modern models . In fact it can be said Thomson's model has no protons , but a positively charged cells . Dalton atomic model of the influence can be seen clearly in the Thomson model . Dalton speculated that atoms were solid, and Thomson s this idea in his model of electrons and protons by grouping together.

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Catode tube x-ray

Thomson’s experiment to determine the ratio between and electron’s mass (e/m).

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2.1.4 Rutherford’s atomic theory

In 1910 , Ernest Rutherford experimented on the truth of these models by conducting an experiment now known as Rutherford scattering

( Rutherford

scattering experiment ) . Rutherford found a particle - α , a particle that is emitted by radioactive atoms , in the year 1909 . These particles have a positive charge , and the fact is we now know that α -particles such as atoms of helium released from the electron , giving the charge 2 +. In scattering experiments , the flow of α particles is directed onto the sheets of gold . This gold leaf is selected by Rutherford as it can be made very thin - only a few atoms thick gold. When the particle - α across the sheets of gold , Rutherford can measure how many α -particles will be scattered by the gold atom by observing the flash of light particles hit the screen scintilator - α . Under the Thomson atomic theory , hypotesys of Rutherfod - α particles will be deflected slightly , when the proton - α particles of gold declined positively charged high.

However, in reality , Rutherford scattering experiment showed results clearly reject the hypothesis and of course the atomic model of Thomson.

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Rutherfod found most alpha particles can penetrate to slabs of gold without deflected . Simultaneously, Rutherford also discovered that alpha particles deflected slightly , but with great surprise , Rutherford also discovered that some alpha particles deflected at a sharp angle back to the radioactive source . To explain the existence of most of the α -particles which penetrate the sheets of gold without deflected , Rutherford and then develop a model of atomic nuclei. In this model , Rutherford put a large proton ( such as experiments and previous models ) in the center of the atom . Rutherford theorized that there are protons around a large space that is empty of all particles except electrons are rare . This large open space give the reason for the alpha particles that are not unbending. The alpha particles are estimated to have diverted some of the protons close enough so that the deflected by electrostatic forces . While some of the alpha particles deflected back to the source is estimated to have undergone collisions with the core so that reflected back by electrostatic forces .

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2.1.5 Niels Bohr’s atomic theory In 1913 Niels Bohr atom model Bohr tried to explain through the concept that follow the orbits of electrons around the atomic nucleus contains protons and neutrons . According to Bohr , there are only a certain number of orbits , and the distinctions between the orbit of one another is the orbital distance from the nucleus. The existence of electrons in orbit either low or high depending entirely by the electron energy levels . So that the electrons in a low orbit will have less energy than electrons in higher orbits . Bohr's

orbiting

electrons

and

connecting

observations of gas through a spectrum of thought that some of the energy contained in the electrons can be changed and therefore electrons can change depending on the change of orbital energy. In the current situation through the use of low pressure gas , the electrons become de - excitation , and move to a lower orbit . In this change, the electrons lose some energy which is the second energy level difference orbit. This emitted energy can be viewed in the form of a photon of light wavelengths , based on the difference in energy levels both orbits . In short , Bohr submit : 1. Electrons in atoms move around the nucleus in certain paths , does not emit energy. electron trajectories are called skin or electron energy levels . 2. Electrons can move from one trajectory to another trajectory .

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3. Electron transfer from high to low energy levels with energy transmitting . Being the displacement of electrons from low to high levels of energy absorption with energy. 4. Electrons moving in an orbit is at a stationary state , meaning that electrons do not emit or absorb energy. Although the Bohr atom model sufficient to model the hydrogen spectrum , the model proved not enough to predict a more complex spectrum of elements

2.1.6 James Chadwick’s atomic theory

In 1932 , Rutherford atomic model be modified slightly by the discovery of the neutron by James Chadwick . Chadwick found that α particle bombardment of beryllium to produce neutrons , uncharged particles , but with a mass slightly larger than the proton mass . Thus ,

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contemporary atomic model is a model with a large nucleus which contains protons and neutrons surrounded by a thin cloud of electrons. The presence of neutrons also explains why the atomic mass is heavier than the total mass of protons and electrons. With a basic understanding of the fundamental parts of atoms like electrons , protons , and neutrons , then it can be enabled by the existence of more complex models and complete more than enough atoms that can explain the nature and characteristics of atoms and atomic compounds .

2.2

The birth of quantum mechanics

2.2.1 The wave nature of particles

In the first half of the 20th century, began to note that the electromagnetic waves, which were previously considered pure wave, behaves like a particle (photon). French physicist Louis Victor de Broglie (1892-1987) assumes that the opposite may also true, namely materials also behave like waves. Starting from the Einstein equation, E = with p is the momentum of the photon, c the speed of light and E is energy, he got the relationship: E = hν = ν = c / λ or hc / λ = E, then h / λ = p ... (2.12) De Broglie assumes each particle with momentum p = mv is accompanied by a wave (wave of matter) with wavelength λ is defined in equation (2.12) (1924). With the increasing size of the particle, its

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wavelength becomes shorter. So to macroscopic particles, particles, not possible to observe diffraction and other phenomena associated with waves. For microscopic particles, like electrons, the wavelength of the material can be observed. In fact, electron diffraction patterns observed (1927) and prove the theory of De Broglie. 2.2.2 Uncertainty principle

From what has been learned about the wave of the material, can be observed that care must be taken when the macroscopic world theory will be applied in the microscopic world.

German physicist Werner Karl

Heisenberg (1901-1976) stated it is impossible to determine accurately the position and momentum simultaneously very small particles such electrons. To observe the particle, one must radiate particles with light. Collisions between particles with the photons will change the position and momentum of particles. Heisenberg explained that the product of the uncertainty of the position

x and momentum uncertainties

p would be worth about

Planck's constant: λ p = h (2.13) This relationship is called the Heisenberg uncertainty principle. 2.2.3 Schrödinger equation

Austrian physicist Erwin Schrödinger (1887-1961) proposed the idea that the De Broglie equation can be applied not only to the free movement of particles, but also on movements such as electrons bound in atoms. With wide this idea, he formulated the system of wave mechanics. At the same

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time develop a system of Heisenberg's matrix mechanics. Then the second day of this system incorporated into quantum mechanics. In quantum mechanics, the system state described by wave functions. Schrödinger basing his theory on the idea that the system total energy, E can be estimated by solving the equation.

Because this equation has a

resemblance with the equation that expresses the wave in classical physics, then this equation called the Schrödinger wave equation. Wave equations of particles (eg electrons) that move in one direction (eg xdirection) is given by: (-H 2 / 8π 2 m) (d 2 Ψ / dx 2) + VΨ = EΨ ... (2.14) 2.2.4 Quantum Number Axial component of angular momentum are allowed only two values, +1 / 2 (h/2π) and -1 / 2 (h/2π). Spin magnetic quantum number associated with this value (m s = +1 / 2 or -1 / 2). Only the spin quantum number alone in an amount not rounded. Table 2.3 Numbers of quantum Name (quantum number) symbol Values allowed Main n 1, 2, 3, ... Azimuth l 0, 1, 2, 3, ... n Magnetic Magnetic spin

1 m (ml) 0, ± 1, ± 2 ,...± l ms +1 / 2, -1 / 2

Other symbols as given in Table 2.4 are commonly used instead. Hydrogen atom energy or hydrogen-like atom is determined only by the principal

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quantum number and energy equations that express identical to that already derived from the theory of Bohr. Table 2.4 Symbols azimuth quantum number value 0 1 2 3 4 symbol s

p d f

g

Electron wave functions called orbital. When the main quantum number n = 1, there is only one value of l, namely 0. In this case there is only one orbital, and a collection of the orbital quantum number for this is (n = 1, l = 0). When n = 2, there are two values l, 0 and 1, are allowed. In case there are four orbital by defined set of quantum numbers: (n = 2, l = 0), (n = 2, l = 1, m = -1), (n = 2, l = 1, m = 0) , (n = 2, l = 1, m = +1).

2.2.5 Configuration Electron Atom When atoms contained more than two electrons, the interaction between electrons must be considered, and difficult to solve the wave equation of this system is very complicated. With the assumption that each electron in the poly-electron atom will move in an electric field which is roughly symmetric orbital for each electron can be defined by three quantum

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numbers n, l and m and the number quantum spin m

s,

as in the case of

hydrogen-like atom . Hydrogen-like atomic energy is determined only by the principal quantum number n, but for poly-electron atoms are mainly determined by n and l. When the atoms have the same quantum number n, the larger l, the higher the energy. 2.2.6 Pauli Exclusion Principle According to the Pauli exclusion principle, only one electron in an atom occupy permitted circumstances defined by a particular set of four quantum numbers, or, at most two electrons can occupy an orbital which is defined by three quantum numbers n, l and m. The two electrons that must have a value of m s are different, in other words its spin antiparallel, and the pair of electrons is called with a pair of electrons. Groups of electrons with the same value of n is called with the skin or skin electron. Notation used for the outer electrons are given in Table 2.5. Table 2.5 Symbols electron shells. n 123 4 5 67 symbol K L M N O P Q Table 2.6 summarizes the maximum number of electrons in each skin, skin from K to N. When the atoms in the most stable state, the ground state, electrons will occupy the orbital with lowest energy, following the Pauli principle. Table 2.6 Maximum number of electrons that occupy each skin.

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skin

l

symbol

Total

n

total skin max

1 2

K L

3

M

4

N

0 0 1 0 1 2 0 1 2 3

1s 2s 2p 3s 3p 3d 4s 4p 4d 4F

electron 2 2 6 2 6 10 2 6 10 14

(2 = 2x12) (8 = 2x22) (18 = 2x32)

(32 = 2x42)

By increasing the orbital energy difference between the orbital energy becomes smaller, and sometimes the sequence into the outer shell inverted electrons . Configuration clearly change when the atomic number changes. This is the basic theory of periodic law. It should be added here, using the symbols given in Table 2.6, the electron configuration of atoms can published. For example, hydrogen atoms in the ground state has one electron skin Diu K and electron configuration (1s

1).

The carbon atom has 2 electrons in the skin of K and 4 electrons in the skin L. electron configuration is (1s 2 2s 2 2p 2).

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CHAPTER 3 PERIODIC TABLE Periodic Table One of the greatest intellectual achievements in chemistry is the periodic table of elements. The periodic table can be printed on one sheet of paper, but what is contained in it and what can be given to us very much and are invaluable. This table is the result of tireless efforts, which originated from Greek times, to know the true nature of matter. It can be said Shem chemical scripture. Value of the periodic system not only on the information known to the organization, but also its ability to predict the nature of the unknown. Actual efficacy of the periodic table is located here. 3.1 The proposals before Mendeleev The concept of elements is a very old concept, since the days of Greece, the Greek philosopher said, the material is formed of four elements: earth, water, fire and air. This view was gradually abandoned, and finally in the 17th century definition of elements given by the English chemist Robert Boyle (16,271,691) replaces the old definition earlier. Boyle stated that the element is a substance that can not be broken down into simpler substances.

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Lavoisier proposed a list of elements in his book "Traite de Chemie Elementire."

Although he entered the light and heat in the list, other

of the list is what we call the elements to date.

Besides, he

added to the list of elements that have not been detected, but he believed existed. For example, chlorine at that time has not been isolated, but he added that to the table as a radical from the acid muriatik. Similarly, sodium and potassium is also in the table. In the early 19th century, these elements were isolated by electrolysis, and slowly expanded the list of elements. In the mid-19th century, spectroscopic analysis, from method introduced detecting element and accelerate the accretion of this list. Although welcomed by chemists, emerging problems.

One of the questions was' Is a limited number of

elements? " and another question is' What is the nature of the elements expected to have a certain order? " The discovery of new elements catalysis discussions of this kind. When iodine was found in 1826, German chemist Johann Wolfgang Döbereiner (1780-1849) noted the similarity between this element with elements that have been known to chlorine and bromine. He also detects a trio of other similar elements. This is what is known as the theory triade Döbereiner.

Table 3.1 Triade Döbereiner

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lithium (Li) Sodium (Na)

calcium

Chlorine

sulfur (S)

(Ca) (Cl) strontium Bromine

(Sr) potassium barium (Ba)

(Br) iodine (I)

manganese

(Mn) selenium chromium (Cr) (Se) tellurium

(K)

Iron (Fe)

(Te)

Octave Newland In 1865, classifying the elements based on atomic mass increases, but from the properties of these elements he observed a repetition or periodic nature of the element. The nature of the elements to 8 similar properties to-1 elements, so on the nature of Element 9 has similar properties to the elements of the 2nd. Because of the repetition of such nature, then it is called the Law of Octaves. The following table law of octaves Do 1 H F Cl Co,Ni Br Pd Te Pt, Ir

Re 2 LI NA K Cu Rb Ag Cs Os

Mi 3 Be Mg Ca Zn Sr Cd Ba V

Fa 4 B Al Cr Y Ce,La U Ta Tl

Sol 5 C Si Ti In Zr Sn W Pb

La 6 N P Mn As Di, Mo Sb Nb Bi

Si 7 O S Fe Se Ro, Ru I Au Th

Meyer

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In 1864, Lothar Meyer conducted experiments that looked at the relationship between the increase in atomic mass with element properties. This is done by making the curve of atomic volume versus atomic mass of atoms,

showed

chart

below:

From these graphs, he noticed the regularity of the elements with similar properties. For example, lithium (Li), Sodium (Na), potassium (K), and Rubidium (Rb) is at the top. Even more important is that he saw number elements among these peaks is different. With it, repetition at the octave law does not apply.

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3.2 Prediction Mendeleev and the truth Many other elements grouping ideas put forward but does not satisfy the scientific community that time. However, the theory proposed by the Russian

chemist

Mendeleev

Dmitrij

Ivanovich

(1834-1907),

and

independently by the German chemist Julius Lothar Meyer (1830-1895) is different from other proposals and more persuasive. Both have the same views as follows: Mendeleev and Meyer's view 1.

List of elements that exist at that time might not yet complete.

2.

It is expected that the systematic nature of the elements varies. So the nature of the unknown element can be predicted.

28

At first the theory of Mendeleev failed to attract attention. However, in 1875, indicated that the new element gallium was discovered by French chemist Paul Emile Lecoq de Boisbaudran (18,381,912) was not the other is the existence of eka -aluminum and nature have been predicted by Mendeleev.

Thus, the significance of Mendeleev and Meyer's theory is

slowly accepted. Table 5.2 provides properties predicted by Mendeleev to the elements when it is not known eka-silikon and properties of germanium which was discovered by German chemist Clemens Alexander Winkler (18381904). Table 5.2 Prediction of the nature of eka-silicon elements by comparison with the nature of Mendeleev and later found. Nature Relative atomic mass Density Atomic volume Valence Heat type Meeting type dioxide Tetrakhlorida Boiling Point (° C)

eka-silicon 72 5,5 13 4 0,073 4,7 <100

germanium 72,32 5,47 13,22 4 0,076 4,703 86

Mendeleev published a table which can be regarded as the origin of the modern periodic table. In preparing the table, Mendeleev initially arrange the elements in order of their atomic masses, as its predecessor. However, he states preodic properties and sometimes rearranging elements, resulting in reverse order of atomic mass.

29

Furthermore, the situation is complicated because the procedures for determining the mass of the atom has not been standardized, and sometimes chemists may use a different atomic mass for the same element. This dilemma is slowly resolved after the International Chemical Congress (Congress was held in 1860 in Karlsruhe, . The purpose of this congress to discuss the problem of unification of atomic mass. Cannizzaro take this opportunity to introduce the theory of Avogadro.) First, which was attended by Mendeleev, but difficulties still exist. By basing on the valence in determining the atomic mass, Mendeleev bit much to solve the problem. As has been described Sisem Modern Periodic periodic system is perfected that had been developed by Mendeleev, Similarity properties of the elements with electron configuration of elements, but it also turns out 30

the elements in one group have the same valence electron. Here is a picture of modern periodic system

3.3 Periodic Table And Electron Configurations The periodic table is continuously growing element of the periodic table Mendeleev proposed. Meanwhile, there is a variety of problems. One important issue is how to handle the noble gases, transition elements and rare earth elements.

All these problems properly completed and make

periodic table is more valuable. Periodic table, chemical holy book, should be referred routinely. New class of noble gases easily inserted between the positive elements of highly reactive, alkali metals (group 1) and negative elements of highly reactive, halogen (group 7). Accommodated the transition metal element in the periodic table by inserting a long period despite reasonable not too clear. The real problem 31

is lantanoid.

Lantanoid treated as an element of "extra" and marginally

placed outside the main part of the periodic table.

However, this

procedure does not actually solve the main problem. First, why this extra element is not clear, even more to the puzzle is the question: is there a limit to the number of elements in the periodic table? Because there are elements that are very similar, very difficult to decide how many elements can exist in nature. Bohr theory and experiment Moseley generate theoretical solution of this problem. Elucidation of the periodic table of the first period until the third period can be explained by the theory of electron configuration described in chapter 4. The first period

(1

H and

2

He) associated with the

process of entering the 1s orbital. Similarly, the second period (from 3 Li to 10

Ne) associated with filling the orbital 1s, 2s and 2p, and 3rd period (from

11

Na to 18 Ar) associated with filling the orbital 1s, 2s, 2p, 3s and 3p. Long period begins the period of the 4th. The explanation for this is

because the d orbital of different shape drastically from the circle, and a 3d electron energy is even higher than 4s. Consequently, in the period-4, the electrons will fill the 4s orbital

(19

K and

20

Ca) immediately after filling the

3s and 3p orbital, skip the 3d orbital. Then the electrons begin to occupy the 3d orbital. This process is related to the ten elements from

21

Sc to

30

Zn. 4p orbital subsequent charging process associated with the six elements of the

31

Ga to

36

Kr.

This is the reason why the period-4 contains 18

elements rather than 8. 4F orbital electron energy is much higher than the 4d orbital electrons 4F and thus played no role in the period-4 elements. Period to period-5 is similar to the 4th. Electrons will fill the orbital 5s, 4d and 5P in this order. As a result the period of the 5th will have 18

32

elements. Orbital 4F not involved and this is what is the reason why the number of elements in the 5 is 18. The number of elements included in the 6 th period amounted to 32 because 7x2 = 14 elements involved relating to the filling orbital 4F. Initially the electrons fill the orbital 6s

(55

miraculous exceptions, the element from and then orbital 5d 4F.

(81

57

to

80

Lantanoid series (to

associated with orbital filling 4F. elements

Cs and

56

Ba). Although there are

Hg La relating to charging

71

Lu) rare earth elements

After this process, six main group

to 86 Tl Rn) to follow, this is related to charging 6p orbitas.

Period-7 to start with filling orbital 7s

(87

Fr and

88

Ra), followed by

filling 5f orbital produce aktinoid series of rare earth elements (from

89

Ac

until the element number 103). World element will be expanded further, but among the elements that exist naturally, the element with the largest atomic number is

92

U. U is the element after

92

artificial elements with a

very short half-life. It is difficult to predict the extension of the list of elements of this kind, but very possibly a new element will be very short half-life. In Table 3.3, summarized the relationship between the periodic table and electron configuration. Table 3.3 electron configuration of each period.

33

period

filled orbitals

number of elements

1 (short) 2 (short)

1s 2s, 2p

2 2+6=8

3 (short)

3s, 3p

2+6=8

4 (long)

3d, 4s, 4p

2 + 6 + 10 = 18

5 (long)

4d, 5s, 5P

2 + 6 + 10 = 18

6 (long)

4F, 5d, 6s, 6p

2 + 6 + 10 + 14 = 32

3.4 The nature of the periodic elements 3.4.1 The First Ionization Energy If these elements have been prepared in accordance with their atomic mass, the nature of the element or compound showed periodic, and this observation led to the discovery of the periodic law.

Electron

configuration of elements to determine not only the chemical properties of elements but also its physical properties. The periodic clearly shown for the atomic ionization energy is directly determined by the electron configuration. Ionization energy is defined as the heat of reaction required to remove electrons from neutral atoms, for example, for sodium: Na (g) → Na + (g) + e-(5.1) The first ionization energy, the energy required to move the first electron, shows a very clear periodic as shown in figure 5.1.

For any

period, the ionization energy increases with increasing atomic number and reach the maximum of the noble gases.

The same group in ionization

energy decreases with increasing atomic number. Trends like this can be 34

explained by the number of valence electrons, the nuclear charge, and the number of electrons inside. The second and third ionization energy is defined as the energy needed to move the second and third electron 3.4.2 Electron Affinity And Electronegative Electron affinity is defined as the heat of reaction when the electron is added to the neutral gas atoms, in the reaction. F (g) + E ° → F ° (g) (5.2) Positive values indicate exothermic reaction, endothermic reaction negative. Because not too many atoms that can be added electrons in the gas phase, the data is limited in number compared to the amount of data for the ionization energy.

Table 3.4

shows that the greater electron

affinity for the non-metals than for metals. Table 3.4 Affinity-electron atoms. H Li Na K

72,4 59, 54,0 48,2

The

amount

C O P S

122,5 141,8 72,4 200,7

electronegative

(electrons)

F Cl Br I

that

322,3 348,3 324,2 295,2

is

defined

by

electronegative, which is a measure of atomic electron binding. Chemists from the United States Robert Sanderson Mulliken (1896-1986) defines electronegative comparable to the arithmetic average of the ionization energy and electron affinity 35

Pauling electronegative defines the difference between the two atoms A and B as the difference in binding energy diatomic molecules AB, AA and BB. Assume D (AB), D (AA) and D (BB) is the binding energy respectively for AB, AA and BB. D (AB) is greater than the average geometry D (AA) and D (BB). This is due to hetero-diatomic molecule is more stable than homodiatomic molecules due to the contribution of ionic structures. As a result, Δ (AB), which is defined as follows, will be positive: (AB) = D (AB) - √ D (AA) D (BB)> 0 (5.3) (AB) will be greater with the growing ionic character. By using this value, Pauling electronegative defines x as a measure of atom attract the electron. |X

A-x B

| = √ D (AB) (5.4)

x A and x B are the atoms A and B. electronegative Whatever

electronegative

scale

is

chosen,

it

is

clear

that

electronegative increases from left to right and decreases from top to bottom.

electronegative very useful to understand the nature of the

chemical elements Dipole direction can be represented with arrows that point to a negative charge center with the beginning of the arrow centered at the center of positive charge. Dipole magnitude, rq, called the dipole moment. Dipole moment is a vector quantity and the magnitude is μ and has a direction.

36

The amount of dipole moment can be determined by experiment, but its direction can not be.

Dipole moment of a molecule (molecular dipole

moment) is the resultant vector dipole moment of existing bonds in the molecule.

If there is symmetry in the molecule, a large bond dipole

moment can eliminate one another so that the molecular dipole moment will be small or even zero. 3.4.3 The Oxidation Number Of Atoms There is a clear relationship between oxidation number (or oxidation) atom and its position in the periodic table. The oxidation number of atoms in covalent compounds is defined as an imaginary atomic charges to be acquired when the electron is shared equally shared between bonded atoms (when atoms are bonded together) or handed over all the atom the stronger the attraction (if different atoms bonded .) 3.5 Main Group Elements For main group elements, the oxidation state in most cases is the number of electrons that will be released or received to achieve the electron configuration management, ns 2 np 6 (except for the first period) or electron configuration nd

10

This is obvious for elements of low periods that are of groups 1, 2 and 13-18. For larger periods, the tendency has oxidation number associated with the electron with the electron configuration ns np electron is maintained and will be released. For example, tin Sn and Pb of lead, both class of 14, has a +2 oxidation state with the release of electrons np 2 but retain ns 2 electrons, in addition to oxidation state +4. The same reasons can be used to the fact that the phosphorus P and bismuth Bi, both

37

class 15 with electron configuration ns 2 np 3, has an oxidation number of +3 and +5. Generally, the importance of electron oxidation with ns 2 maintained will become increasingly important for larger periods. For nitrogen and phosphorus compounds, the dominant oxidation state +5, while for bismuth is dominant is +3 and +5 oxidation rather rare. Metallic element and semilogam (silicon germanium Si or Ge) rarely have a negative oxidation value, but for non-metals this phenomenon is common. In the hydrides of nitrogen and phosphorus, NH 3 and PH 3, the oxidation number of N and P is-3. The higher the period of the element, the element will lose this trait and bismuth Bi does not have the negative oxidation number. Among the elements of group 16, the dominant oxidation state-2 as in the case of oxygen O. This trend was again going down to the elements in the period higher. Suppose that oxygen has only a negative oxidation, but the S has a positive oxidation state +4 and +6 are also significant.

3.6 Trasition Elements Although the transition elements have multiple oxidation states, regularity can be recognized. The highest oxidation number of atoms that has five electrons ie the number of d orbitals associated with the current state of all the d electrons (in addition to electron s) issued. So, in the case of scandium with electron configuration (n-1) d 1 ns 2, oxidation 3. Manganese with the configuration (n-1) d 5 ns 2, will have maximum oxidation number +7.

38

If the amount exceeds 5 d electrons, the situation changed. For iron Fe with the electron configuration (n-1) d 6 ns 2, the main oxidation state is +2 and +3. Very rare oxidation state +6. The highest oxidation number of important transition metals like cobalt Co, Nickel Ni, copper Cu and zinc Zn lower than the oxidation number of atoms that lose all the electrons (n-1) d and ns his. Among the elements that exist within the same group, the higher the oxidation state is increasingly important for the elements in a larger period. The size of atoms and ions When Meyer atomic establish volume is defined as the volume of 1 mole of a particular element (atomic mass / density) against atomic number he got a saw tooth-shaped plot. This clearly is evidence that the atomic volume shows keperiodikan. Because a little difficult to determine the atomic volume of all elements that are identical to the standard, this correlation remains qualitative. However, Meyer's contribution in drawing attention to the existence of atomic size periodic noteworthy.

There are some commentators still double if you want to specify the size of atoms because the electron clouds do not have clear boundaries. For the size of metal atoms, we can determine the radius of the atom by dividing the two distances between atoms, as measured by X-ray diffraction analysis. It must be stated that this value depends on crystal shape (eg a simple cubic lattice or a face-centered cube, etc.) And this will produce a double interpretations that. The same problem is also in the determination of ionic radii are determined by X-ray diffraction analysis of ionic crystals. periodic general tendency of the radius of atoms and ions. For example, the radius of cation seperiode element will decrease with increasing atomic

39

number. This is logical because of the greater nuclear charge will attract electrons more strongly. For ionic radius, the larger the period the greater the radius of the ion.

Bibliography Butler, Ian S, John F Harrord. 1989. Inorganic Chemistry Principles and Application. California : The Benjamin or Cummings

40

Chen, S. Phillip. 1979. Inorganic, Organic, and Biological Chemistry. Second edition. United States of America : Harper & Row, Publisher, Inc F. Albert, Cotton etc. 1994. Basic Inorganic Chemistry. Canada : John Wiley and Sons J. Basset. 1965. Inorganic Chemistry a Concise text. Oxford Pergamon Press Therald Moeller. 1952. Inorganic Chemistry an Advance Textbook. New York : John Wiley and Sons http://wawanhar.wordpress.com/2010/01/05/konsep-dasar-atom/ 5:41 pm

41

CHEMICAL BONDING

Chemical bond refers to the forces holding atoms together to form molecules and solids. This force is of electric nature, and the attraction between electrons of one atom to the nucleus of another atom contribute to what is known as chemical bonds. Although electrons of one atom repell electrons of another, but the repulsion is relatively small. So is the repulsion between atomic nuclei. Various theories regarding chemical bond have been proposed over the past 300 years, during which our interpretation of the world has also changed. Some old concepts such as Lewis dot structure and valency are still rather useful in our understanding of the chemical properties of atoms and molecules, and new concepts involving quantum mechanics of chemical bonding interpret modern observations very well. While reading this page, you learn new concepts such as bondlength, bond energy, bond order, covalent bond, ionic bond, polar and non-polar bond etc. These concepts help you understand the material world at the molecular level. Chemical bonds between identical atoms such as those in H2, N2, and O2 are called covalent bonds, in which the bonding electrons are shared. In ionic compounds, such as NaCl, the ions gather and arrange in a systematic fashion to form a solid. The arrangement of (blue) Na+ and (green) Cl- ions in a solid is shown in on the right here. The attraction force between ions are called ionic bonding. Matals such as sodium, copper, gold, iron etc. have special properties such as being good electric conductors. Electrons in these solids move freely throughout the entire solid, and the forces holding atoms together are called metallic bonds. To some extend, metals are ions submerged is electrons. Though the periodic table has only 118 or so elements, there are obviously more substances in nature than 118 pure elements. This is because

42

atoms can react with one another to form new substances called compounds (see our Chemical Reactions module). Formed when two or more atoms chemically bond together, the resulting compound is unique both chemically and physically from its parent atoms. Let's look at an example. The element sodium is a silver-colored metal that reacts so violently with water that flames are produced when sodium gets wet. The element chlorine is a greenish-colored gas that is so poisonous that it was used as a weapon in World War I. When chemically bonded together, these two dangerous substances form the compound sodium chloride, a compound so safe that we eat it every day - common table salt!

In 1916, the American chemist Gilbert Newton Lewis proposed that chemical bonds are formed between atoms because electrons from the atoms interact with each other. Lewis had observed that many elements are most stable when they contain eight electrons in their valence shell. He suggested that atoms with fewer than eight valence electrons bond together to share electrons and complete their valence shells. While some of Lewis’s predictions have since been proven incorrect (he suggested that electrons occupy cube-shaped orbitals), his work established the basis of what is known today about chemical bonding. We now know that there are two main types of chemical bonding, ionic bonding and covalent bonding. Why are carbon dioxide and sodium chloride so different? Why can we divide compounds into two categories that display distinct physical properties? The answers come from an understanding of chemical bonds: the forces that attract atoms to each other in compounds. Bonding involves the interaction

43

between the valence electrons of atoms. Usually the formation of a bond between two atoms creates a compound that is more stable than either of the two atoms on their own. The different properties of ionic and covalent compounds result from the manner in which chemical bonds form between atoms in these compounds. Atoms can either exchange or share electrons. When two atoms exchange electrons, one atom loses its valence electron(s) and the other atom gains the electron(s). This kind of bonding usually occurs between a metal and a nonmetal. Metals have low ionization energies and non-metals have high electron affinities. That is, metals tend to lose electrons and non-metals tend to gain them. When atoms exchange electrons, they form an ionic bond. Atoms can also share electrons. This kind of bond forms between two non-metals. It can also form between a metal and a non-metal when the metal has a fairly high ionization energy. When atoms share electrons, they form a covalent bond. How can you determine whether the bonds that hold a compound together are ionic or covalent? Examining the physical properties of the compound is one method. This method is not always satisfactory, however. Often a compound has some ionic characteristics and some covalent characteristics. You saw this in the previous Thought Lab. For example, hydrogen chloride, also known as hydrochloric acid, has a low melting point and a low boiling point (It is a gas at room temperature). These properties might lead you to believe that hydrogen chloride is a covalent compound. Hydrogen chloride, however, is extremely soluble in water, and the water solution conducts electricity. These properties are characteristic of an ionic compound. Is there a clear, theoretical way to decide whether the bond between hydrogen and chlorine is ionic or covalent? The answer lies in a periodic trend.

Predicting Bond Type Using Electronegativity

44

You can use the differences between electronegativities to decide whether the bond between two atoms is ionic or covalent. The symbol ΔEN stands for the difference between two electronegativity values. When calculating the electronegativity difference, the smaller electronegativity is always

subtracted

from

the

larger

electronegativity,

so

that

the

electronegativity difference is always positive. How can the electronegativity difference help you predict the type of bond? By the end of this section, you will understand the aswer to this question. Consider three different substances: potassium fluoride, KF, oxygen, O2, and hydrochloric acid, HCl. Potassium fluoride is an ionic compound made up of a metal and a non-metal that have very different electronegativities. Potassium’s electronegativity is 0.82. Fluorine’s electronegativity is 3.98. Therefore, ΔEN for the bond between potassium and fluorine is 3.16. Now consider oxygen. This element exists as units of two atoms held together by covalent bonds. Each oxygen atom has an electronegativity of 3.44. The bond that holds the oxygen atoms together has an electronegativity difference of 0.00 because each atom in an oxygen molecule has an equal attraction for the bonding pair of electrons. Finally, consider hydrogen chloride, or hydrochloric acid. Hydrogen has an electronegativity of 2.20, and chlorine has an electronegativity of 3.16. Therefore, the electronegativity difference for the chemical bond in hydrochloric acid, HCl, is 0.96. Hydrogen chloride is a gas at room temperature, but its water solution conducts electricity. Is hydrogen chloride a covalent compound or an ionic compound? Its ΔEN can help you decide, as you will see below.

A Brief Past on Chemical Bond Concepts Various concepts or theories have been proposed to explain the formation of compounds. In particular, chemical bonds were proposed to explain why and how one element reacted with another element.

45

In 1852, E. Frankland proposed the concept of valence. He suggested that each element formed compounds with definite amounts of other elements due to a valence connection. Each element has a definite number of valance. Five years later, F.A. Kekule and others proposed a valence of 4 for carbon. Lines were used to represent valance, and this helped the development of organic chemistry. The structure of benzene was often quoted as an achievement in this development. More than 10 years later, J.H. van't Hoff and le Bel proposed the tetrahedral arrangement for the four valences around the carbon. These theory helped chemists to describe many organic compounds. In the mean time, chemical bonds were thaught to be electric nature. Since electrons have not been discovered as the negative charge carriers, they were thought to be involved in chemical bonds. Following the discoveries of electrons by J.J. Thomson and R. A. Millikan, G.N. Lewis proposed to use dots to represent valence electrons. His dots made the valence electrons visible to chemists, and he has been credited as the originator of modern bonding theory. He published a book, in 1923, called Valence and the Structure of Atoms and Molecules. X-ray diffractions by crystal allow us to calculate details of bondlength and bond angles. Using computers, we are able to generate images of molecules from the data provided by X-ray diffraction studies. These data prompted Linus Pauling to look at The Nature of Chemical Bond, a book that introduced many new concepts such as the resonance, electronegativity, ionic bond, and covalent bond. In England, N.V. Sidgwick and H.E. Powell paid their attention to the lone pairs in a molecule. They developed the valence bond theory, the VSEPR (Valence Shell Electron Pair Repulsion) theory. An excellent summary of VSEPR is given on this link. The application of quantum theory to chemical bonding gave birth to a molecular orbital theory.

46

In this and the few following modules, we will look at some of these concepts in detail.

Lewis Dot Structures For the elements in the 2nd and 3rd periods, the number of valence electrons range from 1 to 8. Lewis dot structure for them are as indicated: . .

.

.

.

.

..

..

Li Be .B. .C. .N: :O: :F : :Ne: `

`

`

`

`

``

Using dots, Lewis made the valence electron visible. The stability of noble gases is now associated with the 8 valence electrons around it. The stability of 8 valence electrons led him to conclude that all elements strive to acquire 8 electrons in the valence shell, and the chemical reaction takes place due to elements trying to get 8 electrons. This is the octet rule. For the hydrogen and helium atoms, 2 electrons instead of 8 are required. For example, the octet rule applies to the following molecules:

H

. H (2

:

H

.

.

H : O :H

electrons) H ''

F

. : :

''

.

.

H : N : H ''

.

H .

H : C : H '

'

. : . N ::: N : : O:

.

. . . O

:: : F:

. . F

.. . .

: : O :: C :: O:

H

47

To draw a Lewis dot structure, all the valence electrons are represented. A good way is to draw a type of dot for the valence electrons of one atom different from types in another. To do this on the computer screen using only type fonts is difficult, but you should draw a few by hand on paper. When a dash is used to represent a bond, it represents a pair of electrons. Thus, in the following representations, a dash represents two electrons, bonding or lone pairs.

_ :S=O: |

_ :O:H _|

_ :O:H

These structures satisfy

|-

the octet rule. Note the two ways of drawing

O

:O:S:O:

:O:O:O:

"|"

"|"

:O:H

:O:H

"

"

the

structures

of

H2SO4.

Ionic and Covalent Bonding: The Octet Rule The physical properties of covalent and ionic compounds. You learned how to distinguish between an ionic bond and a covalent bond based on the difference between the electronegativities of the atoms. By considering what happens to electrons when atoms form bonds, you will be able to explain some of the characteristic properties of ionic and covalent compounds.

48

The Octet Rule Why do atoms form bonds? When atoms are bonded together, they are often more stable. We know that noble gases are the most stable elements in the periodic table. What evidence do we have? The noble gases are extremely unreactive. They do not tend to form compounds. What do the noble gases have in common? They have a filled outer electron energy level. When an atom loses, gains, or shares electrons through bonding to achieve a filled outer electron energy level, the resulting compound is often very stable. According to the octet rule, atoms bond in order to achieve an electron configuration that is the same as the electron configuration of a noble gas. When two atoms or ions have the same electron configuration, they are said to be isoelectronic with one another. For example, Cl− is isoelectronic with Ar because both have 18 electrons and a filled outer energy level. This rule is called the octet rule because all the noble gases (except helium) have eight electrons in their filled outer energy level. (Recall that helium’s outer electron energy level contains only two electrons.) Elements in the 3rd and higher periods may have more than 8 valence electrons. A possible explanation for this is to say that these atoms have d-type atomic orbitals to accommodate more than 8 electrons. In the following molecules, the number of valence electrons in the central atoms are as indicated: S

Molecule

I

X

F6 Cl5 Cl3 eF4

No. of valence electrons

P

for

central atom

2

1

1

1

0

0

2

1

Draw the Lewis dot structures for the above molecules, and count the number of valence electrons for the central atoms. For H 2SO4, the S atom has

49

12 electrons in the structure shown on your right. Each dash represent a chemical bond, which has two electrons. There is a total of 6 bonds around the S atom, and therefore 12 electrons.

OH | O=S=O | OH

When B, Be, and some metals are the central atoms, they have less than 8 valence electrons. The following compounds do form, but the octet rule is not satisfied. These are electron defficient molecules.

B

Molecule

B

S

Cl2 F3 Cl3 nCl2

No. of valence electrons

B

for

4

6

6

6

central atom

Isoelectronic Molecules and Ions Counting the number of valence electrons often help us understand the formation of many molecules and ions. The charged molecule do exist under special circumstance.

50

The molecules of O2 are paramagnetic, and thus, they have unpaired electrons. The first dot structure does not agree with this observed fact, but the second one does. However, the second one does not obey the octet rule. Later, you will learn that the molecular orbital (MO) theory provides a good explanation for the electronic configuration for O2.

Ionic Bonding The electronegativity difference for the bond between sodium and chlorine is 2.1. Thus, the bond is an ionic bond. Sodium has a very low electronegativity, and chlorine has a very high electronegativity. Therefore, when sodium and chlorine interact, sodium transfers its valence electron to chlorine. Sodium becomes Na+ and chlorine becomes Cl−. How does the formation of an ionic bond between sodium and chlorine reflect the octet rule? Neutral sodium has one valence electron. When it loses this electron to chlorine, the resulting Na+ cation has an electron energy level that contains eight electrons. It is isoelectronic with the noble gas neon. On the other hand, chlorine has an outer electron energy level that contains seven

electrons.

When

chlorine

gains

sodium’s electron, it becomes an anion that is isoelectronic with the noble gas argon. You can represent the formation of an ionic bond using Lewis structures. Thus, in an ionic bond, electrons are transferred from one atom to another so that they form oppositely charged ions. The strong force of attraction between the oppositely charged ions is what holds them together.

51

In ionic bonding, electrons are completely transferred from one atom to another. In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely charged ions are attracted to each other by electrostatic forces, which are the basis of the ionic bond. Notice that when sodium loses its one valence electron it gets smaller in size, while chlorine grows larger when it gains an additional valence electron. This is typical of the relative sizes of ions to atoms. Positive ions tend to be smaller than their parent atoms while negative ions tend to be larger than their parent. After the reaction takes place, the charged Na + and Cl- ions are held together by electrostatic forces, thus forming an ionic bond. Ionic compounds share many features in common: 

Ionic bonds form between metals and nonmetals.



In naming simple ionic compounds, the metal is always first, the nonmetal second (e.g., sodium chloride).



Ionic compounds dissolve easily in water and other polar solvents.



In solution, ionic compounds easily conduct electricity.



Ionic compounds tend to form crystalline solids with high melting temperatures. This last feature, the fact that ionic compounds are solids, results from

the intermolecular forces (forces between molecules) in ionic solids. If we consider a solid crystal of sodium chloride, the solid is made up of many positively charged sodium ions (pictured below as small gray spheres) and an equal number of negatively charged chlorine ions (green spheres). Due to the interaction of the charged ions, the sodium and chlorine ions are arranged in an alternating fashion as demonstrated in the schematic. Each sodium ion is attracted equally to all of its neighboring chlorine ions, and likewise for the chlorine to sodium attraction. The concept of a single molecule does not apply to ionic crystals because the solid exists as one continuous system. Ionic solids form crystals with high melting points because of the strong forces between neighboring ions.

52

Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Sodium Chloride Crystal NaCl Crystal Schematic

Transferring Multiple Electrons In sodium chloride, NaCl, one electron is transferred from sodium to chlorine. In order to satisfy the octet rule, two or three electrons may be transferred from one atom to another. For example, consider what happens when magnesium and oxygen combine. The electronegativity difference for magnesium oxide is 3.4 −1.3, 2.1. Therefore, magnesium oxide is an ionic compound. Magnesium contains two electrons in its outer shell. Oxygen contains six electrons in its outer shell. In order to become isoelectronic with a noble gas, magnesium needs to lose two electrons and oxygen needs to gain two electrons. Hence, magnesium transfers its two valence electrons to oxygen. Magnesium becomes Mg+2, and oxygen becomes O2−.

Ionic Bonding That Involves More Than Two Ions Sometimes ionic compounds contain more than one atom of each element. For example, consider the compound that is formed from calcium and fluorine. Because the electronegativity difference between calcium and fluorine is 3.0, you know that a bond between calcium and fluorine is ionic. Calcium has two electrons in its outer energy level, so it needs to lose two electrons according to the octet rule. Fluorine has seven electrons in its outer energy level, so it needs to gain one electron, again according to the octet rule. How do the electrons of these elements interact so that each element

53

achieves a filled outer energy level? In an ionic bond, calcium tends to lose two electrons and fluorine tends to gain one electron. Therefore, one calcium atom bonds with two fluorine atoms. Calcium loses one of each of its valence electrons to each fluorine atom. Calcium becomes Ca+2, and fluorine becomes F−. They form the compound calcium fluoride, CaF2. In the following Practice Problems, you will predict the kind of ionic compound that will form from two elements.

Covalent Bonding The second major type of atomic bonding occurs when atoms share electrons. As opposed to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when two (or more) elements share electrons. Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons). This most commonly occurs when two nonmetals bond together. Because both of the nonmetals will want to gain electrons, the elements involved will share electrons in an effort to fill their valence shells. A good example of a covalent bond is that which occurs between two hydrogen atoms. Atoms of hydrogen (H) have one valence electron in their first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H 2. Because the hydrogen compound is a combination of equally matched atoms, the atoms will share each other's single electron, forming one covalent bond. In this way, both atoms share the stability of a full valence shell. Unlike

ionic

compounds,

covalent

molecules

exist

as

true

molecules. Because electrons are shared in covalent molecules, no full ionic charges are formed. Thus covalent molecules are not strongly attracted to one another. As a result, covalent molecules move about freely and tend to exist as liquids or gases at room temperature.

54

You have learned what happens when the electronegativity difference between two atoms is greater than 1.7. The atom with the lower electronegativity transfers its valence electron(s) to the atom with the higher electronegativity. The resulting ions have opposite charges. They are held together by a strong ionic bond. What happens when the electronegativity difference is very small? What happens when the electronegativity difference is zero? As an example, consider chlorine. Chlorine is a yellowish, noxious gas. What is it like at the atomic level? Each chlorine atom has seven electrons in its outer energy level. In order for chlorine to achieve the electron configuration of a noble gas according to the octet rule, it needs to gain one electron. When two chlorine atoms bond together, their electronegativity difference is zero. The electrons are equally attracted to each atom. Therefore, instead of transferring electrons, the two atoms each share one electron with each other. In other words, each atom contributes one electron to a covalent bond. A covalent bond consists of a pair of shared electrons. Thus, each chlorine atom achieves a filled outer electron energy level, satisfying the octet rule. When two atoms of the same element form a bond, they share their electrons equally in a pure covalent bond. Elements with atoms that bond to each other in this way are known as diatomic elements. When atoms such as carbon and hydrogen bond to each other, their electronegativities are so close that they share their electrons almost equally. Carbon and hydrogen have an electronegativity difference of only 2.6 −2.2, 0.4. One atom of carbon forms a covalent bond with four atoms of hydrogen. The compound methane, CH4, is formed. Each hydrogen atom shares one of its electrons with the carbon. The carbon shares one of its four valence electrons with each hydrogen. Thus, each hydrogen atom achieves a filled outer energy level, and so does carbon. (Recall that elements in the first period need only two electrons to fill their outer energy level.) When analyzing Lewis structures that show covalent bonds, count the shared electrons as if they belong to each of the bonding atoms. In

55

the following Practice Problems, you will represent covalent bonding using Lewis structures.

Multiple Covalent Bonds Atoms sometimes transfer more than one electron in ionic bonding. Similarly, in covalent bonding, atoms sometimes need to share two or three pairs of electrons, according to the octet rule. For example, consider the familiar diatomic element oxygen. Each oxygen atom has six electrons in its outer energy level. Therefore, each atom requires two additional electrons to achieve a stable octet. When two oxygen atoms form a bond, they share two pairs of electrons. This kind of covalent bond is called a double bond. Double bonds can form between different elements, as well. For example, consider what happens when carbon bonds to oxygen in carbon dioxide. To achieve a stable octet, carbon requires four electrons, and oxygen requires two electrons. Hence, two atoms of oxygen bond to one atom of carbon. Each oxygen forms a double bond with the carbon. When atoms share three pairs of electrons, they form a triple bond. Diatomic nitrogen contains a triple bond. Try the following problems to practise representing covalent bonding using Lewis structures. For every pair of electrons shared between two atoms, a single covalent bond is formed. Some atoms can share multiple pairs of electrons, forming multiple covalent bonds. For example, oxygen (which has six valence electrons) needs two electrons to complete its valence shell. When two oxygen atoms form the compound O2, they share two pairs of electrons, forming two covalent bonds. Lewis dot structures are a shorthand to represent the valence electrons of an atom. The structures are written as the element symbol surrounded by dots that represent the valence electrons. The Lewis structures for the elements in the first two periods of the periodic table are shown below.

56

Lewis Dot Structures

Lewis structures can also be used to show bonding between atoms. The bonding electrons are placed between the atoms and can be represented by a pair of dots or a dash (each dash represents one pair of electrons, or one bond). Lewis structures for H2 and O2 are shown below.

H2

H:H

H-H or

O2

Metallic Bonding In this chapter, you have seen that non-metals tend to form ionic bonds with metals. Non-metals tend to form covalent bonds with other nonmetals and with themselves. How do metals bond to each other? We know that elements that tend to form ionic bonds have very different electronegativities. Metals bonding to themselves or to other metals do not have electronegativity differences that are greater. Therefore, metals probably do not form ionic bonds with each other. Evidence bears this out. A pure metal, such as sodium, is soft enough to be cut with a butter knife. Other pure metals, such as copper or gold, can be drawn into wires or hammered into sheets. Ionic compounds, by contrast, are hard and brittle. Do metals form covalent bonds with each other? No. They do not have enough valence electrons to achieve stable octets by sharing electrons. Although metals do not form covalent bonds, however, they do share their electrons.

57

In metallic bonding, atoms release their electrons to a shared pool of electrons. You can think of a metal as a nonrigid arrangement of metal ions in a sea of free electrons. The force that holds metal atoms together is called a metallic bond. Unlike ionic or covalent bonding, metallic bonding does not have a particular orientation in space. Because the electrons are free to move, the metal ions are not rigidly held in a lattice formation. Therefore, when a hammer pounds metal, the atoms can slide past one another. This explains why metals can be easily hammered into sheets. Pure metals contain metallic bonds, as do alloys. An alloy is a homogeneous mixture of two or more metals. Different alloys can have different amounts of elements. Each alloy, however, has a uniform composition throughout. One example of an alloy is bronze. Bronze contains copper, tin, and lead, ed together with metallic bonds.

Polar and nonpolar covalent bonding There are, in fact, two subtypes of covalent bonds. The H2 molecule is a good example of the first type of covalent bond, the nonpolar bond. Because both atoms in the H2 molecule have an equal attraction (or affinity) for electrons, the bonding electrons are equally shared by the two atoms, and a nonpolar covalent bond is formed. Whenever two atoms of the same element bond together, a nonpolar bond is formed. A polar bond is formed when electrons are unequally shared between two atoms. Polar covalent bonding occurs because one atom has a stronger affinity for electrons than the other (yet not enough to pull the electrons away completely and form an ion). In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. A good example of a polar covalent bond is the hydrogenoxygen bond in the water molecule. Water molecules contain two hydrogen atoms (pictured in red) bonded to one oxygen atom (blue). Oxygen,

with

six

valence

electrons, needs two

additional electrons to complete its valence shell. Each H2O: a water molecule 58

hydrogen contains one electron. Thus oxygen shares the electrons from two hydrogen atoms to complete its own valence shell, and in return shares two of its own electrons with each hydrogen, completing the H valence shells. The primary difference between the H-O bond in water and the H-H bond is the degree of electron sharing. The large oxygen atom has a stronger affinity for electrons than the small hydrogen atoms. Because oxygen has a stronger pull on the bonding electrons, it preoccupies their time, and this leads to unequal sharing and the formation of a polar covalent bond. When two bonding atoms have an electronegativity difference that is greater than 0.5 but less than 1.7, they are considered to be a particular type of covalent bond called a polar covalent bond. In a polar covalent bond, the atoms have significantly different electronegativities. The electronegativity difference is not great enough, however, for the less electronegative atom to transfer its valence electrons to the other, more electronegative atom. The difference is great enough for the bonding electron pair to spend more time near the more electronegative atom than the less electronegative atom. For example, the bond between oxygen and hydrogen in water has an electronegativity difference of 1.24. Because this value falls between 0.5 and 1.7, the bond is a polar covalent bond. The oxygen attracts the electrons more strongly than the hydrogen. Therefore, the oxygen has a slightly negative charge and the hydrogen has a slightly positive charge. Since the hydrogen does not completely transfer its electron to the oxygen, their respective charges are not 1 and 1, but rather δ symbol δ

+

+

and δ -. The

(delta plus) stands for a partial positive charge. The symbol δ

-

(delta minus) stands for a partial negative charge. The dipole Because the valence electrons in the water molecule spend more time around the oxygen atom than the hydrogen atoms, the oxygen end of the molecule develops a partial negative charge (because of the negative charge on the electrons). For the same reason, the hydrogen end of the molecule develops a partial positive charge. Ions are not formed; however, the molecule develops a partial electrical charge across it called a dipole. The water dipole

59

is represented by the arrow in the pop-up animation (above) in which the head of the arrow points toward the electron dense (negative) end of the dipole and the cross resides near the electron poor (positive) end of the molecule.

MOLECULAR GEOMETRY Molecules are three-dimensional objects that occupy a threedimensional world, it is easy to forget this after seeing so many depictions of molecular structures on a two-dimensional page. In general, only the smallest molecules can be said to have a fixed geometrical shape; the icosahedral C 60 “soccer ball” is a rare exception. In most molecules, those parts ed by single bonds can rotate with respect to each other, giving rise to many 60

different geometric forms. However, the local ("coordination") geometry surrounding a given atom that is covalently bound to its neighbors is constant. Being able to understand and predict coordination geometry is an important part of chemistry and is the subject of this section. The Lewis electron-dot structures you have learned to draw have no geometrical significance other than depicting the order in which the various atoms are connected to one another. Nevertheless, a slight extension of the simple shared-electron pair concept is capable of rationalizing and predicting the geometry of the bonds around a given atom in a wide variety of situations. Electron-pair repulsion The valence shell electron pair repulsion (VSEPR) model that we describe here focuses on the bonding and nonbonding electron pairs present in the outermost (“valence”) shell of an atom that connects with two or more other atoms. Like all electrons, these occupy regions of space which we can visualize as electron clouds— regions of negative electric charge, also known as orbitals— whose precise character can be left to more detailed theories. The covalent model of chemical bonding assumes that the electron pairs responsible for bonding are concentrated into the region of apace

between

the

bonded

atoms. The fundamental idea of VSEPR regions

thoery of

is

that

negative

these electric

charge will repel each other, causing them (and thus the chemical bonds that they form) to stay as far apart as possible. Thus the two electron clouds contained in a simple triatomic molecule AX2 will extend out in opposite directions; an angular separation of 180° places the two bonding orbitals as far away from each other they can get. We therefore expect the two chemical bonds to extend in opposite directions, producing a linear molecule. If the central atom also contains one or more pairs of nonbonding electrons, these additional regions of negative charge will behave very much like those associated with the bonded atoms. The orbitals containing the 61

various bonding and nonbonding pairs in the valence shell will extend out from the central atom in directions that minimize their mutual repulsions. If the central atom possesses partially occupied d-orbitals, it may be able to accommodate five or six electron pairs, forming what is sometimes called an “expanded octet”. Digonal and trigonal coordination Linear molecules As we stated above, a simple triatomic molecule of the type AX2 has its two bonding orbitals 180° apart, producing a molecule that we describe as having linear geometry.

Examples of triatomic molecules for which VSEPR theory predicts a linear shape are BeCl2 (which, you will notice, doesn't possess enough electrons to conform to the octet rule) and CO 2. If you write out the electron dot formula for carbon dioxide, you will see that the C-O bonds are double bonds. This makes no difference to VSEPR theory; the central carbon atom is still ed to two other atoms, and the electron clouds that connect the two oxygen atoms are 180° apart. Trigonal molecules In an AX3 molecule such as BF3, there are three regions of electron density extending out from the central atom. The repulsion between these will be at a minimum when the angle between

any

two

is

(360° ÷ 3) = 120°.

This

requires that all four atoms be in the same plane; the resulting shape is called trigonal planar, or simply trigonal.

62

Tetrahedral coordination Methane, CH4, contains a carbon atom bonded to four hydrogens. What bond angle would lead to the greatest possible separation between the electron clouds associated with these bonds? In analogy with the preceding two cases,

where

the

bond

angles

were

360°/2=180° and 360°/3=120°, you might guess 360°/4=90°, if so, you would be wrong. The latter calculation would be correct if all the atoms were constrained to be in the same plane (we will see cases where this happens later), but here there is no such restriction. Consequently, the four equivalent bonds will point in four geometrically equivalent directions in three dimensions corresponding to the four corners of a tetrahedron centered on the carbon atom. The angle between any two bonds will be 109.5°. This is the most important coordination geometry in Chemistry: it is imperative that you be able to sketch at least a crude perspective view of a tetrahedral molecule. It is interesting to note that the tetrahedral coordination of carbon in most of its organic compounds was worked out in the nineteenth century on purely geometrical grounds and chemical evidence, long before direct methods of determining molecular shapes were developed. For example, it was noted that there is only one dichloromethane, CH2Cl2.

63

If the coordination around the carbon were square, then there would have to be two isomers of CH2Cl2, as shown in the pair of structures here. The distances between the two chlorine atoms would be different, giving rise to differences in physical properties would allow the two isomers to be distinguished and separated. The existence of only one kind of CH2Cl2 molecule means that all four positions surrounding the carbon atom are geometrically equivalent, which requires a tetrahedral coordination geometry. If you study the tetrahedral figure closely, you may be able to convince yourself that it represents the connectivity shown on both of the "square" structures at the top. A threedimensional ball-and-stick mechanical model would illustrate this very clearly. Tetrahedrally-coordinated carbon chains Carbon atoms are well known for their tendency to link together to form the millions of organic molecules that are known. We can work out the simpler hydrocarbon chains by looking at each central atom separately. Thus the hydrocarbon ethane is essentially two CH3 tetrahedra ed end-to-end. Similar alkane chains having the general formula H3C–(CH2)n–CH3 (or CnH2n+2) can be built up; a view of pentane, C5H12, is shown below.

64

Notice that these "straight chain hydrocarbons" (as they are often known) have a carbon "backbone" structure that is not really straight, as is illustrated by the zig-zag figure that is frequently used to denote hydrocarbon structures. Coordination geometry and molecular geometry Coordination number refers to the number of electron pairs that surround a given atom; we often refer to this atom as the central atom even if this atom is not really located at the geometrical center of the molecule. If all of the electron pairs surrounding the central atom are shared with neighboring atoms, then the coordination geometry is the same as the molecular geometry. The application of VSEPR theory then reduces to the simple problem of naming (and visualizing) the geometric shapes associated with various numbers of points surrounding a central point (the central atom) at the greatest possible angles. Both classes of geometry are named after the shapes of the imaginary geometric figures (mostly regular solid polygons) that would be centered on the central atom and would have an electron pair at each vertex. If one or more of the electron pairs surrounding the central atom is not shared with a neighboring atom (that is, if it is a lone pair), then the molecular geometry is simpler than the coordination geometry, and it can be worked out by inspecting a sketch of the coordination geometry figure. Tetrahedral coordination with lone pairs In the examples we have discussed so far, the shape of the molecule is defined by the coordination geometry; thus the carbon in methane is tetrahedrally coordinated, and there is a hydrogen at each corner of the tetrahedron, so the molecular shape is also tetrahedral. It is common practice to represent bonding patterns by "generic" formulas such as AX4, AX2E2, etc., in which "X" stands for bonding pairs and "E" denotes lone pairs. (This convention is known as the "AXE method") The bonding geometry will not be tetrahedral when the valence shell of the central atom contains nonbonding electrons, however. The reason is that the nonbonding electrons are also in orbitals that occupy space and repel the other orbitals. This means that in figuring the coordination number around the central atom, we must count both the bonded atoms and the nonbonding pairs. The water molecule: AX2E2 65

In the water molecule, the central atom is O, and the Lewis electron dot formula predicts that there will be two pairs of nonbonding electrons. The oxygen atom will therefore

be

tetrahedrally

coordinated,

meaning that it sits at the center of the tetrahedron as shown below. Two of the coordination positions are occupied by the shared electron-pairs that constitute the O–H bonds, and the other two by the non-bonding pairs. Thus although the oxygen atom is tetrahedrally

coordinated,

the

bonding

geometry (shape) of the H2O molecule is described as bent. There

is

an

important

difference

between bonding and non-bonding electron orbitals. Because a nonbonding orbital has no atomic nucleus at its far end to draw the electron cloud toward it, the charge in such an orbital will be concentrated closer to the central atom. As a consequence, nonbonding orbitals exert more repulsion on other orbitals than do bonding orbitals. Thus in H2O, the two nonbonding orbitals push the bonding orbitals closer together, making the H–O–H

angle

104.5°

instead

of

the

tetrahedral angle of 109.5°. Although the water molecule is electrically neutral, it is not electrically uniform; the non-bonding electrons create a higher concentration of negative charge (blue color) at the oxygen end, making the hydrogen side relatively positive (red). Ammonia: AX3E The electron-dot structure of NH3 places one pair of nonbonding electrons in the valence shell of the nitrogen atom. This means that there are three bonded atoms and one lone pair, for a coordination number of four 66

around the nitrogen, the same as occurs in H2O. We can therefore predict that the three hydrogen atom will lie at the corners of a tetrahedron centered on the nitrogen atom. The lone pair orbital will point toward the fourth corner of the tetrahedron, but since that position will be vacant, the NH3 molecule itself cannot be tetrahedral. Instead, it assumes a pyramidal shape. More precisely, the shape is that of a trigonal pyramid (i.e., a pyramid having a triangular base). The hydrogen atoms are all in the same plane, with the nitrogen above (or below, or to the side; molecules of course don’t know anything about “above” or “below”!) The fatter orbital containing the non-bonding electrons pushes the bonding orbitals together slightly, making the H–N–H bond angles about 107°. Central atoms with five bonds Compounds of the type AX5 are formed by some of the elements in Group 15 of the periodic table; PCl5 and AsF5 are examples. In what directions can five electron pairs arrange themselves in space so as to minimize their mutual repulsions? In the cases of coordination numbers 2, 3, 4, and 6, we could imagine that the electron pairs distributed themselves as far apart as possible on the surface of a sphere; for the two higher numbers, the resulting shapes correspond to the regular polyhedron having the same number of sides. The problem with coordination number 5 is that there is no such thing as a regular polyhedron with five vertices. In 1758, the great mathematian Euler proved that there are only five regular convex polyhedra, known as the platonic solids: tetrahedron (4 triangular faces), octahedron (6 triangular faces), icosahedron (20 triangular faces), cube (6 square faces), and dodecahedron (12 pentagonal faces). Chemical examples of all are known; the first icosahedral molecule, LaC 60 (in which the La atom has 20 nearest C neighbors) was prepared in 1986. Besides the five regular solids, there can be 15 semi-regular isogonal solids in which the faces have different shapes, but the vertex angles are all the same. These geometrical principles are quite important in modern structural chemistry.

67

The shape of PCl5 and similar molecules is a trigonal bipyramid. This consists simply of two triangular-base pyramids ed base-to-base. Three of the chlorine atoms are in the plane of the central phosphorus atom (equatorial positions), while the other two atoms are above

and

below

this

plane

(axial

positions). Equatorial and axial atoms have different geometrical relationships to their neighbors, and thus differ slightly in their chemical behavior. In 5-coordinated molecules containing lone pairs, these non-bonding orbitals (which you will recall are closer to the central atom and thus more likely to be repelled by other orbitals) will preferentially reside in the equatorial plane. This will place them at 90° angles with respect to no more than two axially-oriented bonding orbitals. Using this reasoning, we can predict that an AX 4E molecule (that is, a molecule in which the central atom A is coordinated to four other atoms “X” and to one nonbonding electron pair) such as SF4 will have a “see-saw” shape; substitution of more nonbonding pairs for bonded atoms reduces the triangular bipyramid coordination to even simpler molecular shapes, as shown below.

Octahedral coordination Just as four electron pairs experience the minimum repulsion when they are directed toward the corners of a tetrahedron, six electron pairs will try to point toward the corners of an octahedron. An octahedron is not as complex a shape as its name might imply; it is simply two square-based pyramids ed base to base. You should be able to sketch this shape as well as that of the tetrahedron.

68

The shaded plane shown in this octahedrally-coordinated molecule is only one of three equivalent planes defined by a four-fold symmetry axis. All the ligands are geometrically equivalent; there are no separate axial and equatorial positions in an AX6 molecule. At first, you might think that a coordination number of six is highly unusual; it certainly violates the octet rule, and there are only a few molecules (SF 6 is one) where the central atom is hexavalent. It turns out, however, that this is one of the most commonly encountered coordination numbers in inorganic chemistry. There are two main reasons for this: 

Many transition metal ions form coordinate covalent bonds with lonepair electron donor atoms such as N (in NH 3) and O (in H2O). Since transition elements can have an outer configuration of d10s2, up to six electron pairs can be accommodated around the central atom. A coordination number of 6 is therefore quite common in transition metal hydrates, such as Fe(H2O)63+.



Although the central atom of most molecules is bonded to fewer than six other atoms, there is often a sufficient number of lone pair electrons to bring the total number of electron pairs to six.

Octahedral coordination with lone pairs There are well known examples of 6-coordinate central atoms with 1, 2, and 3 lone pairs. Thus all three of the molecules whose shapes are depicted below possess octahedral coordination around the central atom. Note also that the orientation of the shaded planes shown in the two rightmost images are arbitrary; since all six vertices of an octahedron are identical, the planes could just as well be drawn in any of the three possible vertical orientations.

69

Summary of VSEPR The VSEPR model is an extraordinarily powerful one, considering its great simplicity. Its application to predicting molecular structures can be summarized as follows: 1. Electron pairs surrounding a central atom repel each other; this repulsion will be minimized if the orbitals containing these electron pairs point as far away from each other as possible. 2. The coordination geometry around the central atom corresponds to the polyhedron whose number of vertices is equal to the number of surrounding electron pairs (coordination number). Except for the special case of 5, and the trivial cases of 2 and 3, the shape will be one of the regular polyhedra. 3. If some of the electron pairs are nonbonding, the shape of the molecule will be simpler than that of the coordination polyhedron. 4. Orbitals that contain nonbonding electrons are more concentrated near the central atom, and therefore offer more repulsion than bonding pairs to other orbitals. While VSEPR theory is quite good at predicting the general shapes of most molecules, it cannot yield exact details. For example, it does not explain why the bond angle in H2O is 104.5°, but that in H2S is about 90°. This is not surprising, considering that the emphasis is on electronic repulsions, without regard to the detailed nature of the orbitals containing the electrons, and thus of the bonds themselves. At this point we are ready to explore the three dimensional structure of simple molecular (covalent) compounds and polyatomic ions. We will use a model called the Valence Shell Electron-Pair Repulsion (VSEPR) model that is based on the repulsive behavior of electron-pairs. This model is fairly powerful in its predictive capacity. To use the model we will have to memorize a collection of information. The table below contains several columns. We already have a concept of bonding pair of electrons and non-bonding pairs of electrons. Bonding pairs of electrons are those electrons shared by the central atom and any atom to which it is bonded. Non-bonding pairs of electrons are those pairs of electrons on an individual atom that are not shared with another atom. 70

In the table below the term bonding groups (second from the left column) is used in the column for the bonding pair of electrons. Groups is a more generic term. Group is used when a central atom has two terminal atoms bonded by single bonds and a terminal atom bonded with two pairs of electrons (a double bond). In this case there are three groups of electrons around the central atom and the molecualr geometry of the molecule is defined accordingly. The term electron-pair geometry is the name of the geometry of the electron-pairs on the central atom, whether they are bonding or non-bonding. Molecular geometry is the name of the geometry used to describe the shape of a molecule. The electron-pair geometry provides a guide to the bond angles of between a terminal-central-terminal atom in a compound. The molecular geometry is the shape of the molecule. So when asked to describe the shape of a molecule we must respond with a molecular geometry.

71

# of lone pair # of bonding groups (pair electrons on electrons) on 'central' atom 'central' atom

Electron -pair Geomet ry

Bo nd An gle

0

2

0

3

trigonal planar trigonal planar

1

2

trigonal planar

bent

less than 12 0

0

4

tetrahedral

tetrahedral

109.5

1

3

tetrahedral

2

2

linear

Molecu lar Geome try

tetrahedral

linear

trigonal pyrami dal

bent

180 120

less than 10 9. 5 less than 10 9. 5

trigonal bipyra midal

90, 120 an d 18 0 90, 120 an d 18 0

5

trigonal bipyram idal

1

4

trigonal bipyram idal

seesaw

2

3

trigonal bipyram idal

T-shaped

90 and 180

3

2

trigonal bipyram idal

linear

180

0

6

octahedral

octrahedral

90 and 180

1

5

octahedral

square pyrami 90 and 180 dal

2

4

octahedral

square planar 90 and 180

0

Note: for bent molecular geometry when the electron-pair geometry is trigonal planar the bond angle is slightly less than 120 degrees, around 118 degrees. For trigonal pyramidal geometry the bond angle is slightly less than 109.5 degrees, around 107 degrees. For bent molecular geometry when the electron-pair geometry is tetrahedral the bond angle is around 105 degrees.

72

If asked for the electron-pair geometry on the central atom we must respond with the electron-pair geometry. Notice that there are several examples with the same electron-pair geometry, but different molecular geometries. You should note that to determine the shape (molecular geometry) of a molecule you must write the Lewis structure and determine the number of bonding groups of electrons and the number of non-bonding pairs of electrons on the central atom, then use the associated name for that shape. The table below summarizes the molecular and electron-pair geometries for different combinations of bonding groups and nonbonding pairs of electrons on the central atom. Lets consider the Lewis structure for CCl4. We can draw the Lewis structure on a sheet of paper. The most convenient way is shown here. Notice

that

there are two

kinds

electron

groups in this

structure.

Bonding

electrons,

which

shared

by

a

of

are

pair of atoms

and

nonbonding

electrons,

which belong

to a particular

atom but do

not

participate

in bonding. In

CCl4

the

central carbon

atom

has

four bonding groups of electrons. Each chlorine atom has three nonbonding pairs of electrons. The arrangement of the atoms is correct in my structure. That is the carbon is the central atom and the four chlorine atoms are terminal. But this drawing does not tell us about the shape of the molecule. Lets look at what I mean. Here we have a ball and stick model of CCl4.

73

This is a nice representation of a two dimensional, flat structure. The Cl-C-Cl bond angles appear to be 90 degrees. However, the actual bond angles in this molecule are 109.5 degrees. What does this do to our geometry?

Lets rotate this molecule to see what has happened. We see the actual molecular geometry is not flat, but is tetrahedral. This ball and stick model does not adequately represent why the molecule has to have this 3-dimensional

74

arrangement. The shape we see is the only possible shape for a central carbon atom with four bonds. This geometry is a direct result of the repulsion experienced by the four groups of bonding electrons. The shape of this molecule is a result of the electrons in the four bonds positioning themselves so as to minimize the repulsive effects. This was seen in the 'balloon' example we used in class. When four balloons of the same size are tied together the natural arrangement is as a tetrahedron. With bond angles of 109.5 degrees. As indicated in Table I, any compound containing a central atom with four bonding groups (pairs) of electrons around it will exhibit this particular geometry. By recognizing that electrons repel each other it is possible to arrive at geometries which result from valence electrons taking up positions as far as possible from each other. The position of the atoms is dictated by the position of the bonding groups of electrons. These ideas have been explored and have resulted in a theory for molecular geometry known as Valence Shell Electron-Pair Repulsion Theory. VSEPR focuses on the positions taken by the groups of electrons on the central atom of a simple molecule. The positions can be predicted by imagining that all groups of electrons, whether they are bonding pairs of electrons (single bonds), nonbonding pairs or groups of electrons (multiple bonds) move as far apart as possible. Arriving at the geometry of a molecules requires writing a correct Lewis structure, determining the number of bonding groups and nonbonding groups on the central atom of the molecule and then recalling, from memory, the correct geometry based on the numbers of bonding and nonbonnding groups of electrons. Lone pairs of electrons are assumed to have a greater repulsive effect than bonding pairs. Because of the nonbonding pairs of electrons are spread over a larger volume of space compared to bonding electrons. Because nonbonding electrons are spread over more space they repel other electrons from a greater region of space. So it is more favorable, energetically, for nonbonding pairs of electrons to be as far away as possible from each other in space. So LP-LP repulsions are > LP-BP repulsions > BP-BP repulsions.

75

This can be used to explain the change in bond angles observed in going from methane to ammonia to water. To predict the geometry of a molecule a reasonable Lewis structure must be written for the molecule. The number of bonding and nonbonding pairs of electrons on the central atom are then determined. Let's progress, systematically, through the five basic electron-pair geometries and detail the variations in molecular geometries that can occur. Two Electron Pairs (Linear) The basic geometry for a molecule containing a central atom with two pairs of electrons is linear. BeF2 is an example. Another example of a linear compound is CO2. However, its Lewis structure contains two double bonds. We need to recognize that multiple bonds should be treated as a group of electron pairs when arriving at the molecular geometry. Three Electron Pairs (Trigonal Planar) The basic geometry for a molecule containing a central atom with three pairs of electrons is trigonal planar. BF3 is an example. If we replace a bonding pair with a lone pair, as in SO2, the geometry is described as bent or angular. Four Electron Pairs (Tetrahedral) The basic geometry for a molecule containing a central atom with four pairs of electrons is tetrahedral. An example of this geometry is CH4. As we replace bonding pairs with nonbonding pairs the molecular geometry become trigonal pyramidal (three bonding and one nonbonding), bent or angular (two bonding and two nonbonding) and linear (one bonding and three nonbonding). Notice that compounds with the same number of terminal atoms, BF 3 and NF3, do not necessarily have the same geometry. In this case BF3 has three bonding pairs and no nonbonding pairs with a geometry of trigonal planar, while NF3 has three bonding pairs and one nonbonding pair with a geometry of trigonal pyramidal. Also note that SO2 and H2O have a similar descriptor for their respective geometry. Although each molecule can be described as having a bent geometry the respective bond angles are different. For SO2 the O-S-O angle is near 120 degrees, actually slightly less than 120, about 118 degrees, for H2O the H-O-H angle is near 105 degrees. Five Electron Pairs (Trigonal Bipyramidal) 76

The basic geometry for a molecule containing a central atom with five pairs of electrons is trigonal bipyramidal. An example of this geometry is PCl5. As we replace bonding pairs with nonbonding pairs the molecular geometry changes to seesaw (four bonding and one nonbonding), T-shaped (three bonding and two nonbonding) and linear (two bonding and three nonbonding). This is an interesting system because of the two different types of terminal atoms in the structure, axial and equitorial. The equitorial terminal atoms are those in the trigonal plane. The axial atoms are those above and below the trigonal plane. When the first bonding pair of electrons is replaced with a nonbonding pair that occurs in the trigonal plane. the reason for this is due to the smaller replusions between the lone pair and the bonding pairs of electrons. If the lone pair replaced an axial atom the repulsions would be greater. So as the bonding pairs of electrons are replaced with nonbonding pairs the equitorial atoms are replaced. So as we move from trigonal bipyramidal to linear the nonbonding pairs of electrons occupy the equitorial plane, not the axial positions. Six Electron Pairs (Octahedral) The basic geometry for a molecule containing a central atom with six pairs of electrons is octahedral. An example of this geometry is SF 6. As we replace bonding pairs with nonbonding pairs the molecular geometry changes to square pyramidal(five bonding and one nonbonding) to square planar (four bonding and two nonbonding). There are no other combinations of bonding groups and nonbonding pairs of electrons when the electron-pair geometry is octahedral. The replacement of the first bonding group can occur in any position and always produces a square pyramidal molecular geometry. However the seond bonding group replaced is always opposite the first producing the square planar molecular geometry.

REFERENCES

77

Basset, J. 1965. Inorganic Chemistry: A Concise Text/ J. Basset. Oxford: Pergamon Press. Butler, Ian S. 1989. Inorganic Chemistry Principle and Aplication. California: The Benjamin Cummings. Cotton, F. Albert, etc. 1994. Basic Inorganic Chemistry Third Edition. Canada: John Wiley & Sons, Inc. Moeller, Therald. 1952. Inorhanic Chemistry: An Advance Textbook/Therald Moeller. New York: John Willey & Sons, Inc. Purcell, Keith F, etc. 1982. An Introduction to Inorganic Chemistry. Tokyo: Holt-Saunders Japan.

78

INTRODUCTION

Berilium khlorida BeCl2 and stanium (II) khlorida SnCl2 property of structure is looking like because having molecule formula is looking like. But, simply having structure first compound is linear is secondly curving. This thing is explainable with difference of orbital atomic of applied. If electrons fills orbital atomic of to follow principle Aufbau, electron will field orbital atomic of having low energy. Two electrons is permitted to fills orbital one. According to principle Pauli, there is no electron having an set of same correct quantum number of problem is arising is will is put down where electron four carbon atom. Has been contended that low electron configuration of atom is configuraton with number of couple electrons does is note is maximum and still permitted by order Pauli indium setting orbital is of with the same energy ( indium carbon case is orbital three of of 2p). Indium this case initially all electrons will bound the same spin quantum number ( namely, + 1/2 or - 1/2)

Berilium is atom with two valence electrons and electron configuration ( 1s2 2s2). That having ilium forms tying as atom divalen, orbital of 2s and 2p must form orbital couple are hybridization sp. Because second orbital hibrida sp forms angle of tying 180°, Becl2 thereby linear

Theoritical of VSEPR and also hybridization of atomic orbital will give conclusion of the same molecule structure and ion. Although the theory VSEPR only base on jerk between electron pairses, and the hybridization theory gives his theoretical justifikasi.

79

DISCUSSION

I.

Devi

nisi Hybridization

Hybridization is a concept coalesces it is orbital is atoms forms new hybrid orbital matching with qualitative explanation of atomic bond character. Orbital concept hybridization very useful in explaining form of molecular orbital from a molecule. The hybridization theory developed by chemistry Linus Pauling in explaining molecule structure like methane (CH4).

In hybridization case that is simple, this approach based on orbital of hydrogen atom. Orbital hybridization assumed as aliance from orbital of atom which overlap one other same with proportion varying.

Orbital mixture of sulfur and p at the same atom is equivalent with one orbital p. Result of his is visible in picture 1(b) constructive effect happened in direction of orbital phase of sulfur and p berimpit, whereas the mixing effect is destructive in direction which at the oposite. This effect yields big lobe at direction as of phase together with small lobe in opposite orientation. Improvement of a kind of this direction character related to the same atomic orbital mixture yields is orbital with higher level direction called as hybridization, and yielded to be orbital called as to be orbital hybridization or hybrid orbital. 80

Picture 1. The same atomic orbital mixing effect. ( a) change of direction to orbital mixture of p with different direction. ( b) Improvement of direction to orbital mixture of sulfur and orbital of p ( hybridization effect).

Hybrid orbital measures up to important following relating to chemical bond forming : 1.

Directivity becomes higher, and overlap with species coming from the

direction increases 2.

Distribution hybrid orbital electron to become asymetric, and

closeness of electron at direction of improvement becomes excelsior to become strong attractive force between this cores with near core.

a.

Hybridization sp3

Hybridization explains atoms tying from the aspect of approach an atom. For a coordinating carbon in tetrahedal ( like methane, CH4), hence carbon shall have is orbital having 81

correct symmetry with 4 hydrogen atom. Carbon ground state configuraton is 1s 2 2s2 2px1 2py1.

Valence bond theory predicts, based on at existence two orbital of p loaded by half, that C will form two covalent bonds, that is CH 2. But, methylene is a real reactive molecule, so that just valence bond theory insufficient to explain existence CH4.

Orbital of ground state cannot be applied for tying in CH4. Although excitation of electron 2s to orbital of 2p theoretically permits four tyings and as according to valence bond theory, this thing means there will be some tyings CH4 having different binding energy because of difference of orbital overlap level.

This idea has been refuted experimentally, every hydrogen at CH4 can be discharged from carbon with the same energy. To explain existence of this CH 4 molecule, hence the hybridization theory applied. Initial step of hybridization is excitation out of one ( or more) electron. Proton forming hydrogen atomic nucleus will draw one of carbon valence electron. This thing causes excitation, removes electron 2s to orbital of 2p. This thing increases atomic nucleus influence to valence electrons by increasing effective nuclear potential.

Hybridization process of Methane gaseous molecules (CH4), atom has atom configuraton H: 1s1 and atom configuraton C: 1s2 2s2 2Px1 2py1 2pz0.

82

In banding 4 atom H to become CH4, hence 1 electron ( orbital of 2s) from atom C will be promoted into orbital of 2pz, so that the electron configuration becomes: 1s 2 2s1 2px1 2py1 2pz1

Change happened covers orbital 1 of 2s and orbital 3 of 2p, hence called as hybridization sp3. If hybrid orbital sp3 makes four tyings? with other species, this orbital will yield molecule tetrahedral like CH4, SiH4, NH4 with angle of tetrahedral bond 109,47º. Strength of tying for relative orbital fourth of equivalent, like picture 2. Molecule structure tetrahedral enough stable, so that many molecules having this structure.

Picture 2. Form of molecule with hybridization sp3

Theoritical of orbital hybridization, methane valence electrons ought to have the same energy level, but the fotoelekron spectrum indicating that there are two ribbons, one by 12,7 83

eV ( one electron pairses) and one by 23 eV ( three electron pairses). This explainable unconsistency if assuming existence of additional orbital merger happened when orbital of sp 3 ts forces with orbital 4 of hydrogen.

Picture 3. Ethane Structure ( CH3-CH3)

Sigma bond (σ) formed by every tip of overlap. Molecule giration can around angle of tying.

b.

Hybridization sp2

Ethylene (C2H4) has double bond two among the carbons. Carbon will did hybridization sp2 because orbtial hybrid will only form sigma bond and one pi bonds like the one required for double bond two among carbons. Tying hidrogen-karbon has the same length and bond strength. 84

In hybridization sp2, orbital of 2s only t forces with two orbital of 2p forms orbital 3 of sp2 with one orbital of p remains. In ethylene, two carbon atoms forms a sigma bond with overlap with two orbital of sp2 other carbon and every carbon forms two covalent bonds with hydrogen with overlap s-sp2 which is angular 120°. Pi bond between vertical carbon atoms with molecule area and formed by overlap 2p-2p ( but, pi bond may happened and also no).

Hybridization process sp2, simply through phase as follows. Electron residing in at orbital of 2s promoted and made a move at orbital of 2Py.

So is formed hybrid orbital sp2, what can react with other atom by forming tying which approximately equal. Orbital of p which is not is hybridization to earns overlap yields second tying, tying π. Tying π overlap aside, happened at top and under from molecule. 85

Picture 3. Bonding At Ethene

c.

Hybridization sp

Chemical bond in compound like alkyne with triple bond is explained with hybridization sp. In modeling this, orbital of 2s only t forces with one orbital-p, yields two orbital of sp and leaves over two orbital p. Chemical bond in acetylene ( ethyne) consisted of overlap sp-sp between two carbon atoms forms sigma bond, and two addition pi bonds formed by overlap p-p. Every carbon also tying with hydrogen with angular s-sp overlap 180°.

Hybrid orbital sp yields linear molecule A-B-C or A-B-C-D ( like BeCl2, HgBr2, HCN, C2H2) in linear attributed to tying σ with angle of tying 180º. In HCN and C2H2, besides at tying CN

86

and CC is formed with hybridization sp two tying sets π because type overlap π with parallel direction with tying σ, and as a result triple bond C≡N and C≡C formed.

Picture 4. Bonding At Etuna

87

d.

Other Hybridization

Besides above hybrid orbital, other hybrid orbital type entangling orbital of d also important. As shown in Tabel this orbital 1 relates to forming various molecule structures.

Tables 1 Hybridization and molecule structure

88

A set of orbital of hybridization dsp3 at phosphorus atom ( PCl5)

II. plex Hybridization and Compound Geometry

Com

Valence bond theory developed by pauling. Based on this theory coordination compound formed from reaction of between acids Lewis ( central atom or ion) with alkaline Lewis ( ligand) through coordination covalent bond between both. In having central coordination compound or atom complex compound or ion is certain coordination number.

a.

Coordination Compound Geometry

Theoritical of VSEPR ( valence shell electron pair repulsion), outmost skin electron couples of incentre atom a molecule will stay on course which is each other berjauhan causing reject refuses between electron couples in each the tying mimimal. Based on at this principle, hence coordination compound geometry in general can be predicted based on the ligand amounts, be linear geometry, trigonal planar, tetrahedral, bipiramida trigonal, and oktahedral for complex with coordination number each 2, 3, 4, 5 and 6.

b.

Distortion Jahn-teller

Distortion Jahn-Teller is deviation of complex geometry (from oktahedral becomes tetragonal) which caused by existence of electron at orbital of d at his(its central ion. In this 89

case ligand viewed as negative charge, for the reason will get jerk by electron (also haves negative charge) found on orbital d. Even though only electrons at certain orbital which its(the jerks is effective so that distortion Jahn-Teller is observed. At tables of following summarized distortion yielded by orbital electrons of d at complex " oktahedral".

System

Structure predicted

Description

High spin

Distorsi tetragonal

Is not observed

d1, d6

Distorsi tetragonal

Is not observed

d2, d7

Is not distortion

Proven experimentally

d3, d8

Distortion tetragonal which is big

Proven experimentally

d4, d9

Is not distortion

Proven experimentally

d5, d10

Is not distortion

Proven experimentally

Low Spin

Distortion tetragonal which is big

Yields square complex

d6 d8

Coordination compound geometry with coordination number 2, 3, 4 and 6 ed to table 2 hereunder.

Tables 2. Complex Compound Geometry with a few Coordination number

Coordination number 2 3

Geometry Linear Flat triangle

Contoh [Ag(NH3)2]+, [Cu(CN)2][HgCl3]-, [AgBr(PPh3)2]

4 4

Tetrahedral Square

[FeCl4]2-, [Zn(NH3)4]2+ [Ni(CN)4]2-, [Pt(CN)4]290

5 6

Trigonal bipiramidal Octahedral

[CuCl5]3-, [Fe(CO)5] [CoF6]3-, [Fe(CN)6]3-

Based On Valence bond theory, geometry from complex compound closely related with orbital geometry from central atoms or ion applied in tying forming. If paid attention example at above tables, seen that the compound geometry or complex ion nothing that looks like orbital 3 geometry of p, or orbital 5 geometry d. thereby inferential that in forming of coordination covalent bond of orbital using central atom or ion of hibrida formed through hybridization process.

Hybridization is orbital forming process of hibrida with the same energy level from orbital of atom which the type and energy level differs in. Number of orbital of hibrida formed is equal to number of orbital of atom involving in hybridization. Hereunder some example of orbital hybridizations of central atom or ion along with orbital geometry of hibrida obtained:

Tables 3 Example Of Hybridizations In Complex Compound

Hibridisasi

Involving Atomic

Orbital Geometry of

Example

91

Orbital

Hibrida

sp

1 orbital s and 1

Linear

[Cu(CN)2]-, [Ag(NH3)2]+

sp2

orbital p 1 orbital s and 2

Flat triangle

[HgCl3]-, [AgBr(PPH3)2]

sp3

orbital p 1 orbital s and 3

Tetrahedral

[FeCl4]2-, [Zn(NH3)4]2+

dsp

orbital p 1 orbital d 1 orbita s

Bujursangkar

[Ni(CN)4]2-, [Cu(NH3)4]2+

2

dsp or sp d

and 2 orbital p 1 orbital d 1 orbita s trigonal bipiramidal

[CuCl5]3-, [Fe(CO)5]

d2sp3 or sp3d2

and 3 orbital p 2 orbital d 1 orbita s

[CoF6]3-, [Fe(CN)6]3-

3

3

oktahedral

and 3 orbital p

In explaining forming of tying at complex compound, orbital of hibrida from central atom or ion is depicted with box, radian or line. Forming of tying entangles some steps, covers promotion of electron; orbital forming of hibrida; and forming of tying between metals with ligand through overlap between orbital of hibrida empty metal orbitally ligand containing free electron pairs.

c.

Ligand Strength

Every ligand has strength of certain field. Based on strength research of ligand field is:

CN- > NO2- > NH3 > H2O > F- > OH- > Cl- > Br- > I-

Strength of magnetic field in molecule determined by whether there or is not there electron which couple.

92

 If all paired electronses hence will experience rejection in magnetic field, called as diamagnetic  If there are electron that is is not is couple hence will experience withdrawal by magnetic field, called as character paramagnetik. More and more many electrons that is is not is couple more and more strong the paramagnetik character

d.

Character Paramagnetik

Transition metal differs from metals from short period, because at ions forming, this metal doesn't discharge all the electronic valencies. Transition metal from length period that is first, forms ion by discharging two electrons from orbital of 4s by leaving orbital of 3d which tidal complete. This electrons according to law Hund is not the couple and spin at the oposite.

Because atom with electron is not couple have moment permanent magnet, hence this transition metal compounds haves the character of paramagnetik mean can be pulled by magnetic field. Paramagnetic susceptibilitl which is caused by electron that is is not is couple shown by formula:

µ = 4S ( S +1)

S is number of spins quanta. Because every electron has spin quanta ± ½ hence

µ = n( n + 2)

n= number of electrons is not couple 93

At hybridization entangling orbital of d, there is two kinds of possibility that hybridization. If in orbital hybridization of d entangled is orbital of d which beyond skin from orbital of sulfur and p having hybridization, hence complex formed conceived of external orbital complex, or orbital outer of complex. On the contrary, if in hybridization entangled is orbital of d in orbital skin of sulfur and p having hybridization, hence the complex is named orbital complex in or orbital inner of complex. Generally orbital complex in more stably is compared to external orbital complex, because energy involved in by orbital complex formation in smaller compared to energy involving in external orbital complex formation. Hybridization to orbital of d is residing in in orbital of sulfur and p is required smaller energy, because the energy level not too far.

Example :

 [Ni(CO)4]; has geometric structure of tetrahedral Ni28

: [Ar] 3d8 4s2

: [Ar] 3d8

4s2

4p0

 Electron at orbital of 4s experiences promotion to orbital of 3d, so that orbital empty 4s and can experience hybridization orbitally 4p forms orbital of hibrida sp3.

 Ni28

: [Ar] 3d8

4s 4p hybridization sp3

94

 Orbital of hibrida sp3 which has been formed then applied for tying with 4 ligand CO which rendering is each is free electron pairs

 [Ni(CO)4]

: [Ar] 3d10

sp3

 Because all paired electronses, hence compound haves the character of diamagnetic  [Fe(CN)6]3-; has geometrical form oktahedral

Fe26

: [Ar] 3d6 4s2

Fe3+

: [Ar] 3d5 4s0

: [ Ar] 3d5

4s1

4p0

 Two electron at orbital of d which initialy not couple are paired with other electron of the orbital of d, so that orbital 2 of d which initialy occupied by both the empty electrons and applicable to form orbital of hibridal d2sp3

Fe3+

: [Ar] hibridisasi d2sp3

95

 Because orbital of d applied in this hybridization comes from orbital of d is residing in disebelah in orbital of sulfur and p, hence complex orbitally a kind of this hibrida conceived of orbital complex in ( orbital inner of complex)

[Fe(CN)6]3-

: [Ar]

 Orbital of hibrida d2sp3 formed filled by electron pairs free of ligand CN-

 In complex there are one electrons that is is not is couple, so that complex had the character of paramagnetik.



[Ni(CN)4]2-, has geometrical form segiempat planar

Ni28

: [Ar] 3d8 4s2

: [Ar] 3d8

Ni2+

4s2

4p0

: [Ar] forms orbital of hibrida dsp3

 One of electron at orbital of d which is not is couple is paired with other electron, so that one of orbital of empty d and applicable to form orbital of hibrida dsp3

96

[Ni(CN4)]2-

: [Ar]

 All electrons in this complex couple so that complex had the character of diamagnetic

Most of more opting complex of orbital complex configuraton in, because energy required when hybridization to entangle orbital of d side in smaller compared to energy needed to entangles orbital of outer d. Nevertheless, if it is seen from measurement of the magnetic moment, some complexes simply stays in external orbital complex form.

Question that is often emerges is " When electrons of the orbital of d central ion paired and when is not paired?" in this case is paired or not of electrons is depend on the experiment fact. If from experiment it is obtained that a compound or complex ion haves the character of diamagnetic hence the central atom or ion has orbital of d or other orbital has been loaded is full or has orbital of d or other orbital which has not been loaded is full but all the electrons in a state of couple. In explaining forming of coordination covalent bond between ligands with central atom or ion is entangled excitation phase. This excitation tends to happened if the ligand is strong ligand like CN-, however factor influencing excitation is not only ligand type. Many other factors having an effect on between it is number of ligands, central ion type or atom and the complex geometry.

Until around year of 1943 valence bond theories is the only theory applied by the expert of inorganic chemistry in explaining geometry and magnetism of complex compound nevertheless this theory has weakness, that is:

1.

Cannot explain magnetism change symptom of complex compound

because temperature change. 2.

Cannot explain colour or spektra complex compound

97

3.

Cannot explain stability of complex compound

Limitation of the hibtidisation theory

be cannot explain magnetism character.

Existence of weakness from valence bond theory enables using of other theory of which can explain third of above fact. One of theory is crystal magnet field theory.

III.

Com

plex Compound Nomenclature

Nomenclature applied for complex ion to come from Inorganic Nomenclasture Committes of International Union of Pure Applied Chemistry ( C-IUPAC). For name of visible ligands at tables 4

Tables 4. Name Of Ligands

Ligands NH3 H2 O CO NO NH2-CH2-CH2-NH2

Name of Ligands Amino / amin Aquo Karbonil Nitrosil Etilen diamin (en)

Negative Ligands

Name of Ligands

CN-

Siano

F-

Fluoro

Cl-

Kloro

Br-

Bromo 98

I-

Iodo

OH-

Hidrokso

SCN-

Tiosiano

S2O32-

Tiosulfato

SO42-

Sulfato

CO32-

Karbonato

NO2-

Nitro

O2-

okso

If there is ligand multiple that is conspecific or not of a kind, the numbers expressed as prefix in, three, tetra, and so. For ligands which is difficult, for example ethylen is amen ( shortened en), if the numbers multiple, called as with prefix:

Ethylen bus amen for ( en)= 2 Tris ethylen amen for ( en)= 3 Tetrakis ethylen amen for ( en)=4

a.

Complex Ion Haves Positif Charge

99

Giving of complex compound ion which haves positif charge originally is mentioned name of its(the ligand is then is followed name of its(the metal atom ( central atom). Oxidation number from central atom is written with number romawi in parenthesis.

Example: [Ag(NH3)2]+ = ion diamin perak (1) [Fe(NH3)6]3+ = ion keksamin besi (III) [Fe(H2O)6]2+ = ion heksaquo besi (II)

If the ligand multiple hence ligand sequence is mentioned as according to initial letter alphabet from ligand and followed name of its(the metal atom ( central atom)

Example: [Fe(H2O)4Cl2]+ = ion tetraquo dikloro besi (III) [Co(NH3)4(CN)Br]+ = ion tetraamin bromo siano kobalt (III)

b.

Complex Ion Haves Negative Charge

Giving of name of complex ion which haves negative charge, originally is called as name of the ligand is then is followed name of its central atom is added with suffix at. The oxidation number is written down with number romawi in parenthesis.

Example:

100

[Ag(CN)2]- = ion disiano argemtat (I) [Fe(Cl)6]3- = ion heksakloro ferrat (III)

If the ligand multiple hence ligand sequence is mentioned as according to initial letter alphabet from ligand and followed name of the metal atom ( central atom).

Example: [Cr(H2O)2I4]- = ion diaquo tetraiodo kromat (III) [Co(NH3)2(CN)4]2- = ion diamin tetrasiano kobatlat (II)

c.

Neutral Complex Compound ( nonionic)

Giving of name of neutral complex ion equal to giving of name of complex ion loading negative.

Example: [Zn(H2O)2Br] = diaquo dibromo zinkat(III) [Co(NH3)3(NO2)3] = triamin trinitro kobaltat (III)

d.

Complex Salts

Giving of the name is originally is mentioned the cation is then is followed the anion.

101

Example: K4[Fe(CN)6] = kalium heksasiano ferrat (II) Na3[Co(Br06] = natrium heksabromo kobaltat (III) [Cu(NH3)4]SO4 = tetramin tembaga (II) sulfat [Cr(NH3)5Cl]Cl2 = pentamin kloro krom (III) Klorida

102

CONCLUSION

One of theory applied to explain tying forming, geometry and magnetic of complex compound is valence bond theory. Based on this orbital theory of atom or or central ion before receiving electron pairs free of ligand will experience hybridization. Hybridization is orbital forming of hibrida which the energy level is same from orbital of atom that is conspecific and the energy level differs in. Orbital of hibrida will be filled by free electron couples or electron π is coming from ligand. Complex compound geometry formed is explicable based on hybridization involving in orbital forming of the hibrida.

Magnetic from predictable complex compound based on whether there or is not there electron that is is not couple at orbital of central atom or ion. If there is electron which is not is couple hence complex compound had the character of paramagnetik, and when there are no electron which is not is couple hence complex compound had the character of diamagnetic.

103

REFERENCES

Chang, Raymond. 2003. Edition Core Concepts Base Chemistry Third Volume 1. Jakarta : Erlangga

Effendi. 1998. Coordination Chemistry. Malang: FMIPA IKIP Malang

Effendi. 2003. The theory VSEPR and Molecule Polar. Malang: Bayu Media Publishing.

Kachi, Ohno. 2004. Quantum Chemistry. Tokyo : Iwanami Shoten Publishers 104

Yashito, Takeuchi. Chemical Deliverer Online Textbook. Tokyo : Iwanami Shoten Publisher

http://www.chem-is-try.org/hibridisasi-situs kimia http:// www. chwm-is-try.org/ valence bond theory http://www.chem-is-try.org/ orbital hybridization http://www.chem-is-try.org//complex compound http://suyantakimiafmipaugm.wordpress.com/2010/05/17/bab-iii-stereokoimia-senyawakompleks/

105

Preliminary

Alkali metal The alkali metals are a series of chemical elements forming Group 1 (IUPAC style) of the periodic table: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). (Hydrogen, although nominally also a member of Group 1, very rarely exhibits behavior comparable to the alkali metals). The alkali metals provide one of the best examples of group trends in properties in the periodic table, with well characterized homologous behavior down the group.

Properties The alkali metals are all highly reactive and are never found in elemental forms in nature. Because of this, they are usually stored immersed in mineral oil or kerosene (paraffin oil). They also tarnish easily and have low melting points and densities. Physically, the alkali metals are mostly silver-colored, except for metallic caesium, which has a golden tint. These elements are all soft metals of low density. Chemically, all of the alkali metals react aggressively with the halogens to form ionic salts. They all react with water to form strongly alkaline hydroxides. The vigor of reaction increases down the group. All of the atoms of alkali metals have one electron in their outmost electron shells, hence their only way for achieving the equivalent of filled outmost electron shells is to give up one electron to an element with high electronegativity, and hence to become singly charged positive ions, i.e. cations. When it comes to their nuclear physics, the elements potassium and rubidium are naturally weakly radioactive because they each contain a long half-life radioactive isotope. The element hydrogen, with its solitary one electron per atom, is usually placed at the top of Group 1 of the periodic table for convenience, but hydrogen is not counted as an alkali metal. Under typical conditions, pure hydrogen exists as a diatomic gas consisting of two atoms per molecule. The removal of the single electron of hydrogen requires considerably more energy than removal of the outer electron from the atoms of the alkali metals. As in the halogens, only one additional electron is required to fill in the outermost shell of the hydrogen atom, so hydrogen can in some circumstances behave like a halogen, forming the negative hydride ion. Binary compounds of hydrogen with the alkali metals and some transition metals have been produced in the laboratory, but these are only laboratory curiosities without much practical use. Under extremely high 106

pressures and low temperatures, such as those found at the cores of the planets Jupiter and Saturn, hydrogen does become a metallic element, and it behaves like an alkali metal. (See metallic hydrogen.) The alkali metals have the lowest ionization potentials in their periods of the periodic table, because the removal of their single electrons from their outmost electron shells gives them the stable electron configuration of inert gases. Another way of stating this is that they all have a high electropositivity. The "second ionization potential" of all of the alkali metals is very high, since removing any electron from an atom having a noble gas configuration is difficult to do.

Series of alkali metals, stored in mineral oil ("natrium" is sodium.) All of the alkali metals are notable for their vigorous reactions with water, and these reactions become increasingly vigorous when going down their column in the periodic table towards the heaviest alkali metals, such as caesium. Their chemical reactions with water are as follows: Alkali metal + water → Alkali metal hydroxide + hydrogen gas For a typical example (M represents an alkali metal): 2 M (s) + 2 H2O (l) → 2 MOH (aq) + H2 (g)

Trends Like in other columns of the periodic table, the of the alkali metal family show patterns in their electron configurations, especially their outmost electron shells. This causes similar patterns in their chemical properties: Z

Element

No. of electrons/shell

1

Hydrogen 1

3

Lithium

2, 1

11

Sodium

2, 8, 1 107

19

Potassium 2, 8, 8, 1

37

Rubidium 2, 8, 18, 8, 1

55

Caesium

87

Francium 2, 8, 18, 32, 18, 8, 1

2, 8, 18, 18, 8, 1

The alkali metals show a number of trends when moving down the group - for instance: decreasing electronegativity, increasing reactivity, and decreasing melting and boiling point. Their densities generally increase, with the notable exception that potassium is less dense than sodium, and the possible exception of francium being less dense than caesium. (The highly radioactive element francium only exists in microscopic quantities.) Alkali metal Lithium Sodium Potassium Rubidium Caesium Francium

Standard Melting atomic weight point (K) (u) 6.941 453 22.990 370 39.098 336 85.468 312 132.905 301 (223) 295

Boiling Density point (K) (g·cm−3)

Electronegativity (Pauling)

1615 1156 1032 961 944 950

0.98 0.93 0.82 0.82 0.79 0.70

0.534 0.968 0.89 1.532 1.93 1.87

Group IB element IUPAC has not recommended a specific format for the periodic table, so different conventions are permitted and are often used for group IB. The following d-block transition metals are always considered of group IB: • •

scandium (Sc) yttrium (Y)

When defining the remainder of Group IB, four different conventions may be encountered: •

Some tables [1] include lanthanum (La) and actinium (Ac), (the beginnings of the lanthanide and actinide series of elements, respectively) as the remaining of Group IB. In their most commonly encountered tripositive ion forms, these elements do not possess any partially filled f orbitals, thus resulting in more d-block-like behavior. 108

•

Some tables [2] include lutetium (Lu) and lawrencium (Lr) as the remaining of Group IB. These elements terminate the lanthanide and actinide series, respectively. Since the f-shell is nominally full in the ground state electron configuration for both of these metals, they behave most like d-block metals out of all the lanthanides and actinides, and thus exhibit the most similarities in properties with Sc and Y. For Lr, this behavior is expected, but it has not been observed because sufficient quantities are not available. (See also Periodic table (wide) and Periodic table (extended).)

Some tables [3] refer to all lanthanides and actinides by a marker in Group IB. A third and fourth alternative are suggested by this arrangement: •

The third alternative is to regard all 30 lanthanide and actinide elements as included in Group IB. Lanthanides, as electropositive trivalent metals, all have a closely related chemistry, and all show many similarities to Sc and Y.

•

The fourth alternative is to include none of the lanthanides and actinides in Group IB. The lanthanides possess additional properties characteristic of their partially-filled f orbitals which are not common to Sc and Y. Furthermore, the actinides show a much wider variety of chemistry (for instance, in range of oxidation states) within their series than the lanthanides, and comparisons to Sc and Y are even less useful.

Elements of group IA 109

Hydrogen Hydrogen is the lightest element. It is by far the most abundant element in the universe and makes up about about 90% of the universe by weight. Hydrogen as water (H2O) is absolutely essential to life and it is present in all organic compounds. Hydrogen is the lightest gas. Hydrogen gas was used in lighter-than-air balloons for transport but is far too dangerous because of the fire risk (Hindenburg). It burns in air to form only water as waste product and if hydrogen could be made on sufficient scale from other than fossil fuels then there might be a possibility of a hydrogen economy. Note that while normally shown at the top of the Group 1 elements in the periodic table, the term "alkaline metal" refers only to Group 1 elements from lithium onwards. Table: basic information about and classifications of hydrogen. • Name: Hydrogen • Group in periodic table: 1 • Symbol: H • Group name: (none) • Atomic number: 1 • Period in periodic table: 1 • Atomic weight: 1.00794 (7) [see notes g • Block in periodic table: s-block m r] • Colour: colourless • Standard state: gas at 298 K • Classification: Non-metallic • CAS Registry ID: 1333-74-0 The lifting agent for the ill fated Hindenberg ballooon was hydrogen rather than the safer helium. The image below is the scene probably in a way you have not seen it before. This is a "ray-traced" image reproduced with the permission of Johannes Ewers, the artist, who won first place with this image in the March/April 1999 Internet Raytracing Competition. For details of ray-tracing you can't beat the POV-Ray site.

Isolation Isolation: in the laboratory, small amounts of hydrogen gas may be made by the reaction of calcium hydride with water. CaH2 + 2H2O → Ca(OH)2 + 2H2 This is quite efficient in the sense that 50% of the hydrogen produced comes from water. Another very convenient laboratory scale experiment follows Boyle's early synthesis, the reaction of iron filings with dilute sulphuric acid. 110

Fe + H2SO4 → FeSO4 + H2 There are many industrial methods for the production of hydrogen and that used will depend upon local factors such as the quantity required and the raw materials to hand. Two processes in use involve heating coke with steam in the water gas shift reaction or hydrocarbons such as methane with steam. CH4 + H2O (1100°C) → CO + 3H2 C(coke) + H2O (1000°C) → CO + H2 In both these cases, further hydrogen may be made by ing the CO and steam over hot (400°C) iron oxide or cobalt oxide. CO + H2O → CO2 + H2 Table: valence shell orbital radii for hydrogen. Orbital Radius [/pm]

Radius [/AU]

s orbital

52.9

1.00003

p orbital

no data

no data

d orbital

no data

no data

f orbital

no data

no data

Effective Nuclear Charges 111

The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats. Table: effective nuclear charges for hydrogen 1s no data 2s

no data

2p

no data

3s

no data

3p

no data

3d

no data

4s

no data

4p

no data

4d

no data

5s

no data

5p

no data

5d

no data

6s

no data

6p

no data

4f

no data

7s

Electron binding energies This table contains electron binding energies for hydrogen. Label Orbital eV [literature reference] K

1s

13.6 [1]

112

Notes All values of electron binding energies are given in eV. The binding energies are quoted relative to the vacuum level for rare gases and H2, N2, O2, F2, and Cl2 molecules; relative to the Fermi level for metals; and relative to the top of the valence band for semiconductors.

Lithium Lithium is a Group 1 (IA) element containing just a single valence electron (1s22s1). Group 1 elements are called "alkali metals". Lithium is a solid only about half as dense as water and lithium metal is the least dense metal. A freshly cut chunk of lithium is silvery, but tarnishes in a minute or so in air to give a grey surface. Its chemistry is dominated by its tendency to lose an electron to form Li +. It is the first element within the second period. Lithium is mixed (alloyed) with aluminium and magnesium for light-weight alloys, and is also used in batteries, some greases, some glasses, and in medicine. Table: basic information about and classifications of lithium. • Name: Lithium • Group in periodic table: 1 • Symbol: Li • Group name: Alkali metal • Atomic number: 3 • Period in periodic table: 2 • Atomic weight: [ 6.941 (2)] [see notes g • Block in periodic table: s-block m r] 113

•

Standard state: solid at 298 K

•

Colour: silvery white/grey

•

CAS Registry ID: 7439-93-2

•

Classification: Metallic

Isolation Isolation: lithium would not normally be made in the laboratory as it is so readily available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative lithium ion Li+. The ore spodumene, LiAl(SiO3)2, is the most important commercial ore containing lithium. The α form is first converted into the softer β form by heating to around 1100°C. This is mixed carefully with hot sulphuric acid and extracted into water to form lithium sulphate, Li2SO4, solution. The sulphate is washed with sodium carbonate, Na2CO3, to form a precipitate of the relatively insoluble lithium carbonate, Li2CO3. Li2SO4 + Na2CO3 → Na2SO4 + Li2CO3 (solid) Reaction of lithium carbonate with HCl then provides lithium chloride, LiCl. Li2CO3 + 2HCl → 2LiCl + CO2 +H2O Lithium chloride has a high melting point (> 600°C) meaning that it sould be expensive to melt it in order to carry out the electrolysis. However a mixture of LiCl (55%) and KCl (45%) melts at about 430°C and so much less energy and so expense is required for the electrolysis. cathode: Li+(l) + e- → Li (l) anode: Cl-(l) → 1/2Cl2 (g) + eTable: valence shell orbital radii for lithium. Orbital Radius [/pm]

Radius [/AU]

s orbital

164.1

3.10010

p orbital

no data

no data 114

d orbital

no data

no data

f orbital

no data

no data

Effective Nuclear Charges The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats. Table: effective nuclear charges for lithium 1s 2.69 2s

1.28

2p

no data

3s

no data

3p

no data

3d

no data

4s

no data

4p

no data

4d

no data

5s

no data

5p

no data

5d

no data

6s

no data

6p

no data

4f

no data

7s Electron binding energies This table contains electron binding energies for lithium. Label Orbital eV [literature reference] K

1s

54.7 [2]

Sodium Sodium is a Group 1 element (or IA in older labelling styles). Group 1 elements are often referred to as the "alkali metals". The chemistry of sodium is dominated by the +1 ion Na+. Sodium salts impart a characteristic orange/yellow colour to flames and orange street lighting is orange because of the presence of sodium in the lamp.

115

Soap is generally a sodium salt of fatty acids. The importance of common salt to animal nutrition has been recognized since prehistoric times. The most common compound is sodium chloride, (table salt). Table: basic information about and classifications of sodium. • Name: Sodium • Group in periodic table: 1 • Symbol: Na • Group name: Alkali metal • Atomic number: 11 • Period in periodic table: 3 • Atomic weight: 22.98976928 (2) • Block in periodic table: s-block • Standard state: solid at 298 K • Colour: silvery white •

CAS Registry ID: 7440-23-5

•

Classification: Metallic

The result of adding different metal salts to a burning reaction mixture of potassium chlorate and sucrose. The red colour originates from strontium sulphate. The orange/yellow colour originates from sodium chloride. The green colour originates from barium chlorate and the blue colour originates from copper (I) chloride. The lilac colour that should be evident from the potassium chlorate is washed out by the other colours, all of which are more intense (only to be demonstrated by a professionally qualified chemist following a legally satisfactory hazard asessment). Improperly done, this reaction is dangerous!

116

The picture above shows the colour arising from adding common salt (NaCl) to a burning mixture of potassium chlorate and sucrose. Isolation Isolation: sodium would not normally be made in the laboratory as it is so readily available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative sodium ion Na+. Sodium is present as salt (sodium chloride, NaCl) in huge quantities in underground deposits (salt mines) and seawater and other natural waters. It is easily recovered as a solid by drying. Sodium chloride has a high melting point (> 800°C) meaning that it sould be expensive to melt it in order to carry out the electrolysis. However a mixture of NaCl (40%) and calcium chloride, CaCl2 (60%) melts at about 580°C and so much less energy and so expense is required for the electrolysis. cathode: Na+(l) + e- → Na (l) anode: Cl-(l) → 1/2Cl2 (g) + eThe electrolysis is carried out as a melt in a "Downs cell". In practice, the electrolysis process produces calcium metal as well but this is solidified in a collection pipe and returned back to the melt. Table: valence shell orbital radii for sodium. Orbital Radius [/pm]

Radius [/AU]

s orbital

179.4

3.39059

p orbital

no data

no data

d orbital

no data

no data

f orbital

no data

no data

Effective Nuclear Charges The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats. Table: effective nuclear charges for sodium 1s 10.63 2s

6.57

2p

6.80 117

3s

2.51

3p

no data

3d

no data

4s

no data

4p

no data

4d

no data

5s

no data

5p

no data

5d

no data

6s

no data

6p

no data

4f

no data

7s

Electron binding energies This table contains electron binding energies for sodium. Label Orbital eV [literature reference] K

1s

1070.8 [3]

L

I

2s

63.5 [3]

L

II

2p1/2

30.4 [3]

L

III

2p3/2

30.5 [2]

Potassium Potassium is a metal and is the seventh most abundant and makes up about 1.5 % by weight of the earth's crust. Potassium is an essential constituent for plant growth and it is found in most soils. It is also a vital element in the human diet. Potassium is never found free in nature, but is obtained by electrolysis of the chloride or hydroxide, much in the same manner as prepared by Davy. It is one of the most reactive and electropositive of metals and, apart from lithium, it is the least dense known metal. It is soft and easily cut with a knife. It is silvery in appearance immediately after a fresh surface is exposed. It oxidises very rapidly in air and must be stored under argon or under a suitable mineral oil. As do all the other metals of the alkali group, it decomposes in water with the evolution of hydrogen. It usually catches fire during the reaction with water. Potassium and its salts impart a lilac colour to flames.

118

Table: basic information about and classifications of potassium. • Name: Potassium • Group in periodic table: 1 • Symbol: K • Group name: Alkali metal • Atomic number: 19 • Period in periodic table: 4 • Atomic weight: 39.0983 (1) • Block in periodic table: s-block • Standard state: solid at 298 K • Colour: silvery white •

CAS Registry ID: 7440-09-7

•

Classification: Metallic

The reaction between potassium metal and water (only to be demonstrated by a professionally qualified chemist).

The picture above shows the colour arising from a burning mixture of potassium chlorate (KClO3) and sucrose (only to be demonstrated by a professionally qualified chemist). Isolation Isolation: potassium would not normally be made in the laboratory as it is so readily available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative potassium ion K+. Potassium is not made by the same method as sodium as might have been expected. This is because the potassium metal, once formed by electrolysis of liquid potassium chloride (KCl), is too soluble in the molten salt. 119

cathode: K+(l) + e- → K (l) anode: Cl-(l) → 1/2Cl2 (g) + eInstead, it is made by the reaction of metallic sodium with molten potassium chloride at 850°C. Na + KCl ⇌ K + NaCl This is an equilibrium reaction and under these conditions the potassium is highly volatile and removed from the system in a form relatively free from sodium impurities, allowing the reaction to proceed. Table: valence shell orbital radii for potassium. Orbital Radius [/pm]

Radius [/AU]

s orbital

230.0

4.34528

p orbital

no data

no data

d orbital

no data

no data

f orbital

no data

no data

Effective Nuclear Charges The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats. Table: effective nuclear charges for potassium 1s 18.49 2s

13.01

2p

15.03

3s

8.68

3p

7.73

3d

no data

4s

3.50

4p

no data

4d

no data

5s

no data

5p

no data

5d

no data

6s

no data

6p

no data

4f

no data

7s

Electron binding energies 120

This table contains electron binding energies for potassium. Label Orbital eV [literature reference] K

1s

3608.4 [2]

L

I

2s

378.6 [2]

L

II

2p1/2

297.3 [2]

L

III

2p3/2

294.6 [2]

M

I

3s

34.8 [2]

M

II

3p1/2

18.3 [2]

M

III

3p3/2

18.3 [2]

Rubidium Rubidium can be liquid at ambient temperature, but only on a hot day given that its melting point is about 40°C. It is a soft, silvery-white metallic element of the alkali metals group (Group 1). It is one of the most most electropositive and alkaline elements. It ignites spontaneously in air and reacts violently with water, setting fire to the liberated hydrogen. As so with all the other alkali metals, it forms amalgams with mercury. It alloys with gold, caesium, sodium, and potassium. It colours a flame yellowish-violet. Table: basic information about and classifications • Name: Rubidium • Symbol: Rb • Atomic number: 37 • Atomic weight: 85.4678 (3) [see note g] • Standard state: solid at 298 K •

CAS Registry ID: 7440-17-7

of rubidium. • Group in periodic table: 1 • Group name: Alkali metal • Period in periodic table: 5 • Block in periodic table: s-block • Colour: silvery white •

Classification: Metallic

121

Image adapted with permission from Prof James Marshall's (U. North Texas, USA) Walking Tour of the elements CD. Isolation Isolation: rubidium would not normally be made in the laboratory as it is available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative rubidium ion Rb+. Rubidium is not made by the same method as sodium as might have been expected. This is because the rubidium metal, once formed by electrolysis of liquid rubidium chloride (RbCl), is too soluble in the molten salt. cathode: Rb+(l) + e- → Rb (l) anode: Cl-(l) → 1/2Cl2 (g) + eInstead, it is made by the reaction of metallic sodium with hot molten rubidium chloride. Na + RbCl ⇌ Rb + NaCl This is an equilibrium reaction and under these conditions the rubidium is highly volatile and removed from the system in a form relatively free from sodium impurities, allowing the reaction to proceed. The following are calculated values of valence shell orbital radii, Rmax Table: valence shell orbital radii for rubidium. Orbital Radius [/pm]

Radius [/AU]

s orbital

249.0

4.70616

p orbital

no data

no data

d orbital

no data

no data

f orbital

no data

no data 122

Effective Nuclear Charges The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats. Table: effective nuclear charges for rubidium 1s 36.21 2s

27.16

2p

33.04

3s

21.84

3p

21.30

3d

21.68

4s

12.39

4p

10.88

4d

no data

5s

4.98

5p

no data

5d

no data

6s

no data

6p

no data

4f

no data

7s

Electron binding energies This table contains electron binding energies for rubidium. Label Orbital eV [literature reference] K

1s

15200 [1]

L

I

2s

2065 [1]

L

II

2p1/2

1864 [1]

L

III

2p3/2

1804 [1]

M

I

3s

326.7 [2]

M

II

3p1/2

248.7 [2]

M

III

3p3/2

239.1 [2]

M

IV

3d3/2

113 [2]

M

V

3d5/2

112 [2]

N

I

4s

30.5 [2] 123

N

II

4p1/2

16.3 [2]

N

III

4p3/2

15.3 [2]

Caesium Caesium is known as cesium in the USA. The metal is characterised by a spectrum containing two bright lines in the blue (ing for its name). It is silvery gold, soft, and ductile. It is the most electropositive and most alkaline element. Caesium, gallium, and mercury are the only three metals that are liquid at or around room temperature. Caesium reacts explosively with cold water, and reacts with ice at temperatures above -116°C. Caesium hydroxide is a strong base and attacks glass. Table: basic information about and classifications of caesium. • Name: Caesium • Group in periodic table: 1 • Symbol: Cs • Group name: Alkali metal • Atomic number: 55 • Period in periodic table: 6 • Atomic weight: 132.9054519 (2) • Block in periodic table: s• Standard state: solid at 298 K (but melts only block slightly above this temperature) • Colour: silvery gold •

CAS Registry ID: 7440-46-2

•

Classification: Metallic

124

Image adapted with permission from Prof James Marshall's (U. North Texas, USA) Walking Tour of the elements CD. Isolation Isolation: caesium (cesium in USA) would not normally be made in the laboratory as it is available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative caesium ion Cs+. Caesium is not made by the same method as sodium as might have been expected. This is because the caesium metal, once formed by electrolysis of liquid caesium chloride (CsCl), is too soluble in the molten salt. cathode: Cs+(l) + e- → Cs (l) anode: Cl-(l) → 1/2Cl2 (g) + eInstead, it is made by the reaction of metallic sodium with hot molten caesium chloride. Na + CsCl ⇌ Cs + NaCl This is an equilbrium reaction and under these conditions the caesium is highly volatile and removed from the system in a form relatively free from sodium impurities, allowing the reaction to proceed. It can be purified by distillation. The following are calculated values of valence shell orbital radii, Rmax Table: valence shell orbital radii for caesium. Orbital Radius [/pm]

Radius [/AU]

s orbital

282.4

5.33704

p orbital

no data

no data

d orbital

no data

no data

f orbital

no data

no data

125

Effective Nuclear Charges The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats. Table: effective nuclear charges for caesium 1s 53.90 2s

40.51

2p

50.82

3s

36.38

3p

36.58

3d

40.98

4s

27.04

4p

25.86

4d

22.84

5s

15.44

5p

13.65

5d

no data

6s

6.36

6p

no data

4f

no data

7s

Electron binding energies This table contains electron binding energies for caesium. Label Orbital eV [literature reference] K

1s

35985 [1]

L

I

2s

5714 [1]

L

II

2p1/2

5359 [1]

L

III

2p3/2

5012 [1]

M

I

3s

1211 [2, values derived from reference 1]

M

II

3p1/2

1071 [2]

M

III

3p3/2

1003 [2]

M

IV

3d3/2

740.5 [2]

M

V

3d5/2

726.6 [2]

126

N

I

4s

232.3 [2]

N

II

4p1/2

172.4 [2]

N

III

4p3/2

161.3 [2]

N

IV

4d3/2

79.8 [2]

N

V

4d5/2

77.5 [2]

N

VI

4f5/2

-

N

VII

4f7/2

-

O

I

5s

22.7

O

II

5p1/2

14.2 [2]

O

III

5p3/2

12.1 [2]

Francium Francium occurs as a result of α disintegration of actinium. Francium is found in uranium minerals, and can be made artificially by bombarding thorium with protons. It is the most unstable of the first 101 elements. The longest lived isotope, 223Fr, a daughter of 227Ac, has a half-life of 22 minutes. This is the only isotope of francium occurring in nature, but at most there is only 20-30 g of the element present in the earth's crust at any one time. No weighable quantity of the element has been prepared or isolated. There are about 20 known isotopes. Table: basic information about and classifications of francium. • Name: Francium • Group in periodic table: 1 • Symbol: Fr • Group name: Alkali metal • Atomic number: 87 • Period in periodic table: 7 • Atomic weight: [ 223 ] • Block in periodic table: s-block • Standard state: solid at 298 K • Colour: metallic •

CAS Registry ID: 7440-73-5

•

Classification: Metallic

127

This sample of uraninite contains some francium because of a steady-state decay chain. An estimate suggests there is about 10 -20 grammes of francium (about 1 atom!) at any one time. Image adapted with permission from Prof James Marshall's (U. North Texas, USA) Walking Tour of the elements CD. Isolation Isolation: francium is vanishingly rare and is found only as very small traces in some uranium minerals. It has never been isolated as the pure element. As it is so radioactive, any amount formed would decompose to other elements. Actinium decays by β decay most of the time but about 1% of the decay is by α decay. The "daughter" element of this reaction, which used to be called actinium-K, is now recognized as 22387Fr - the longest-lived isotope of actinium with a half life of about 22 minutes. The following are calculated values of valence shell orbital radii, Rmax Table: valence shell orbital radii for francium. Orbital Radius [/pm]

Radius [/AU]

s orbital

298.6

5.64197

p orbital

no data

no data

d orbital

no data

no data

f orbital

no data

no data

Effective Nuclear Charges The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats. 128

Table: effective nuclear charges for francium 1s no data 2s

no data

2p

no data

3s

no data

3p

no data

3d

no data

4s

no data

4p

no data

4d

no data

5s

no data

5p

no data

5d

no data

6s

no data

6p

no data

4f

no data

7s

Electron binding energies This table contains electron binding energies for francium. Label Orbital eV [literature reference] K

1s

101137 [1]

L

I

2s

18639 [1]

L

II

2p1/2

17907 [1]

L

III

2p3/2

15031 [1]

M

I

3s

4652 [1]

M

II

3p1/2

4327 [1]

M

III

3p3/2

3663 [1]

M

IV

3d3/2

3136 [1]

M

V

3d5/2

3000 [1]

N

I

4s

1153 [2]

N

II

4p1/2

980 [2]

N

III

4p3/2

810 [2]

N

IV

4d3/2

603 [2]

N

V

4d5/2

577 [2]

N

VI

4f5/2

268 [2] 129

N

VII

4f7/2

268 [2]

O

I

5s

234 [2]

O

II

5p1/2

182 [2]

O

III

5p3/2

140 [2]

O

IV

5d3/2

58 [2]

O

V

5d5/2

58 [2]

P

I

6s

34 [1]

P

II

6p1/2

15 [1]

P

III

6p3/2

15 [1]

Elements of group IB

Scandium Scandium is a silvery-white metal which develops a slightly yellowish or pinkish cast upon exposure to air. It is relatively soft, and resembles yttrium and the rare-earth metals more than it resembles aluminium or titanium. Scandium reacts rapidly with many acids. Scandium is apparently a much more abundant element in the sun and certain stars than on earth. Table: basic information about and classifications of scandium. • Name: Scandium • Group in periodic table: 3 • Symbol: Sc • Group name: (none) • Atomic number: 21 • Period in periodic table: 4 • Atomic weight: 44.955912 (6) • Block in periodic table: d-block • Standard state: solid at 298 K • Colour: silvery white •

CAS Registry ID: 7440-20-2

•

Classification: Metallic 130

Image adapted with permission from Prof James Marshall's (U. North Texas, USA) Walking Tour of the elements CD. Isolation Isolation: preparation of metallic samples of scandium is not normally necessary given that it is commercially avaialable. In practice littel scandium is produced. The mineral thortveitite contains 35-40% Sc2O3 is used to produce scandium metal but another important source is as a byproduct from uranium ore processing, even though these only contain 0.02% Sc2O3. Brief description: yttrium has a silvery-metallic lustre. Yttrium turnings ignite in air. Yttrium is found in most rare-earth minerals. Moon rocks contain yttrium and yttrium is used as a "phosphor" to produce the red colour in television screens. Scandium: orbital properties Valence shell orbital radii The following are calculated values of valence shell orbital radii, Rmax Table: valence shell orbital radii for scandium. Orbital Radius [/pm]

Radius [/AU]

s orbital

171.6

3.24169

p orbital

no data

no data

d orbital

59.2

1.11908

f orbital

no data

no data

131

Effective Nuclear Charges The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats. Table: effective nuclear charges for scandium 1s 20.46 2s

14.57

2p

17.05

3s

10.34

3p

9.41

3d

7.12

4s

4.63

4p

no data

4d

no data

5s

no data

5p

no data

5d

no data

6s

no data

6p

no data

4f

no data

7s

132

Electron binding energies This table contains electron binding energies for scandium. Label Orbital eV [literature reference] K

1s

4492 [1]

L

I

2s

498 [2]

L

II

2p1/2

403.6 [2]

L

III

2p3/2

398.7 [2]

M

I

3s

51.1 [2]

M

II

3p1/2

28.3 [2]

M

III

3p3/2

28.3 [2]

133

Yttrium Table: basic information about and classifications of yttrium. • Name: Yttrium • Group in periodic table: 3 • Symbol: Y • Group name: (none) • Atomic number: 39 • Period in periodic table: 5 • Atomic weight: 88.90585 (2) • Block in periodic table: d-block • Standard state: solid at 298 K • Colour: silvery white •

CAS Registry ID: 7440-65-5

•

Classification: Metallic

134

This sample is from The Elements Collection, an attractive and safely packaged collection of the 92 naturally occurring elements that is available for sale. Isolation Isolation: yttrium metal is available commercially so it is not normally necesary to make it in the laboratory. Yttrium is found in lathanoid minerals and the extraction of the yttrium and the lanthanoid metals from the ores is highly complex. Initially, the metals are extractedas salts from the ores by extraction with sulphuric acid (H 2SO4), hydrochloric acid (HCl), and sodium hydroxide (NaOH). Modern purification techniques for these lanthanoid salt mixtures involve selective complexation techniques, solvent extractions, and ion exchange chromatography. Pure yttrium is available through the reduction of YF3 with calcium metal. 2YF3 + 3Ca → 2Y + 3CaF2

Yttrium: orbital properties Valence shell orbital radii The following are calculated values of valence shell orbital radii, Rmax Table: valence shell orbital radii for yttrium. Orbital Radius [/pm]

Radius [/AU]

s orbital

189.1

3.57338

p orbital

no data

no data

d orbital

93.9

1.77384

f orbital

no data

no data

135

Effective Nuclear Charges The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats. Table: effective nuclear charges for yttrium 1s 38.18 2s

28.62

2p

35.00

3s

23.55

3p

23.09

3d

25.40

4s

14.26

4p

12.75

4d

15.96

5s

6.26

5p

no data

5d

no data

6s

no data

6p

no data

4f

no data

7s

Electron binding energies This table contains electron binding energies for yttrium. Label Orbital eV [literature reference] K

1s

17038 [1]

L

I

2s

2373 [1]

L

II

2p1/2

2156 [1]

L

III

2p3/2

2080 [1]

M

I

3s

392 [2, values derived from reference 1]

M

II

3p1/2

310.6 [2]

M

III

3p3/2

298.8 [2]

M

IV

3d3/2

157.7 [3]

M

V

3d5/2

155.8 [3]

N

I

4s

43.8 [2]

N

II

4p1/2

24.4 [2]

N

III

4p3/2

23.1 [2] 136

Lutetium Lutetium: the essentials Brief description: pure metal lutetium has been isolated only in recent years and is one of the more difficult to prepare. It can be prepared by the reduction of anhydrous LuCl3 or LuF3 by an alkali or alkaline earth metal. The metal is silvery white and relatively stable in air. It is a rare earth metal and perhaps the most expensive of all rare elements. It is found in small amounts with all rare earth metals, and is very difficult to separate from other rare elements. Table: basic information about and classifications • Name: Lutetium • Symbol: Lu • Atomic number: 71 • Atomic weight: 174.9668 (1) [see note g] • Standard state: solid at 298 K •

CAS Registry ID: 7439-94-3

of lutetium. • Group in periodic table: 3 • Group name: (none) • Period in periodic table: 6 • Block in periodic table: d-block • Colour: silvery white •

Classification: Metallic

137

This sample is from The Elements Collection, an attractive and safely packaged collection of the 92 naturally occurring elements that is available for sale. Isolation Isolation: lutetium metal is available commercially so it is not normally necessary to make it in the laboratory, which is just as well as it is difficult to isolate as the pure metal. This is largely because of the way it is found in nature. The lanthanoids are found in nature in a number of minerals. The most important are xenotime, monazite, and bastnaesite. The first two are orthophosphate minerals LnPO4 (Ln deonotes a mixture of all the lanthanoids except promethium which is vanishingly rare) and the third is a fluoride carbonate LnCO3F. Lanthanoids with even atomic numbers are more common. The most comon lanthanoids in these minerals are, in order, cerium, lanthanum, neodymium, and praseodymium. Monazite also contains thorium and ytrrium which makes handling difficult since thorium and its decomposition products are radioactive. For many purposes it is not particularly necessary to separate the metals, but if separation into individual metals is required, the process is complex. Initially, the metals are extracted as salts from the ores by extraction with sulphuric acid (H 2SO4), hydrochloric acid (HCl), and sodium hydroxide (NaOH). Modern purification techniques for these lanthanoid salt mixtures are ingenious and involve selective complexation techniques, solvent extractions, and ion exchange chromatography. Pure lutetium is available through the reduction of LuF3 with calcium metal. 2LuF3 + 3Ca → 2Lu + 3CaF2 This would work for the other calcium halides as well but the product CaF 2 is easier to handle under the reaction conditions (heat to 50°C above the melting point of the element in an argon atmosphere). Excess calcium is removed from the reaction mixture under vacuum.

Lutetium: orbital properties

138

Valence shell orbital radii The following are calculated values of valence shell orbital radii, Rmax Table: valence shell orbital radii for lutetium. Orbital Radius [/pm]

Radius [/AU]

s orbital

186.7

3.52829

p orbital

no data

no data

d orbital

95.0

1.79586

f orbital

24.6

0.465285

Effective Nuclear Charges The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats. Table: effective nuclear charges for lutetium 1s 69.62 2s

52.45

2p

66.61

3s

49.53

3p

50.17

3d

57.42

4s

38.27

4p

37.19

4d

35.29

5s

20.96

5p

18.68

5d

20.11

6s

8.80

6p

no data

4f

30.93

7s

Electron binding energies This table contains electron binding energies for lutetium. Label Orbital eV [literature reference] K L

I

1s

63314 [1]

2s

10870 [1]

139

L

II

2p1/2

10349 [1]

L

III

2p3/2

9244 [1]

M

I

3s

2491 [1]

M

II

3p1/2

2264 [1]

M

III

3p3/2

2024 [1]

M

IV

3d3/2

1639 [1]

M

V

3d5/2

1589 [1]

N

I

4s

506.8 [2]

N

II

4p1/2

412.4 [2]

N

III

4p3/2

359.2 [2]

N

IV

4d3/2

206.1 [2]

N

V

4d5/2

196.3 [2]

N

VI

4f5/2

8.9 [2]

N

VII

4f7/2

7.5 [2]

O

I

5s

57.3 [2]

O

II

5p1/2

33.6 [2]

O

III

5p3/2

26.7 [2]

Lawrencium Lawrencium: the essentials Brief description: lawrencium is a synthetic "rare earth metal" which does not occur in the environment. Table: basic information about and classifications of lawrencium. • Name: Lawrencium • Group in periodic table: 3 • Symbol: Lr • Group name: (none) 140

• • •

Atomic number: 103 Atomic weight: [ 262 ] Standard state: presumably a solid at 298 K

• • •

Period in periodic table: 7 Block in periodic table: d-block Colour: unknown, but probably metallic and silvery white or grey in appearance

•

CAS Registry ID: 22537-19-5

•

Classification: Metallic

Lawrencium: orbital properties Valence shell orbital radii The following are calculated values of valence shell orbital radii, Rmax Table: valence shell orbital radii for lawrencium. Orbital Radius [/pm]

Radius [/AU]

s orbital

203.1

3.83753

p orbital

no data

no data

d orbital

104.0

1.96493

f orbital

38.9

0.735260

Effective Nuclear Charges The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats. Table: effective nuclear charges for lawrencium 1s no data 2s

no data

2p

no data

3s

no data

3p

no data

3d

no data

4s

no data

4p

no data

4d

no data

5s

no data

5p

no data

5d

no data

6s

no data

6p

no data

4f

no data

7s

141

Reference 1. 2. 3. 4. 5.

6. 7.

D.R. Lide, (Ed.) in Chemical Rubber Company handbook of chemistry and physics, CRC Press, Boca Raton, Florida, USA, 81st edition, 2000. E. Clementi and D.L.Raimondi, J. Chem. Phys. 1963, 38, 2686. E. Clementi, D.L.Raimondi, and W.P. Reinhardt, J. Chem. Phys. 1967, 47, 1300. J. A. Bearden and A. F. Burr, "Reevaluation of X-Ray Atomic Energy Levels," Rev. Mod. Phys., 1967, 39, 125. J.B. Mann, Atomic Structure Calculations II. Hartree-Fock wave functions and radial expectation values: hydrogen to lawrencium, LA-3691, Los Alamos Scientific Laboratory, USA, 1968. J. C. Fuggle and N. Mårtensson, "Core-Level Binding Energies in Metals," J. Electron Spectrosc. Relat. Phenom., 1980, 21, 275. Gwyn Williams WWW table of values 142

8.

M. Cardona and L. Ley, Eds., Photoemission in Solids I: General Principles (Springer-Verlag, Berlin) with additional corrections, 1978.

143

CHAPTER I IIA ELEMENTS Some of the physical properties usually associated with metals—hardness, high m.p. and b.p.—are noticeably lacking in these metals, but they all have a metallic appearance and are good electrical conductors. Group II metals are more dense, are harder and have higher m.p. and b.p. than the corresponding Group I metals. bonding generally depends on the ratio (number of electrons available for bonding)/(atomic size). The greater this ratio is, the stronger are the bonds between the metal atoms. In the pre-transition metals, this ratio is small and at a minimum in Group I with only one bonding electron. Metallic bond strength is greater in Group II but there are still only two bonding electrons available, hence the metals are still relatively soft and have low melting and boiling points. Hardness, m.p. and b.p. all decrease steadily down Group I, the metallic bond strength decreasing with increasing atomic radius. These changes are not so well marked in Group II but note that beryllium and, to a lesser extent, magnesium are hard metals, as a result of their small atomic size; this property, when coupled with their low density, makes them of some technological importance. Group II (alkaline earths) are known as s-block elements because their valence (bonding) electrons are in s orbitals. PHYSICAL PROPERTIES OF IIA ELEMENTS Element

Atomic

Outer

Density

m.p.

b.p.

Hardness

Be

number 4

electrons 2

(g/ 1.86

(K) 1553

(K) 3243

(Brinell) -

Mg

12

3

1.75

924

1380

30-40

Ca

20

4

1.55

1124

1760

23

Sr

38

5

2.6

1073

1639

20.

Ba

56

6

3.59

998

1910

-

Element

lonisation

Metallic

Ionic

Heat of

)

Hydrat ion 144

energy*

radius

radius

vaporisation energy of

(kJ/mol)

(nm)

(nm)

at 298 K

gaseous ion

Be

2657

0.112

0.031

(kJ/mol) 326

(kJ/mol) 2494

Mg

2187

0.160

0.065

149

1921

Ca

1735

0.197

0.099

177

1577

Sr

1613

0.215

0.113

164

1443

Ba

1467

0.221

0.135

178

1305

Atomic Radius Increases down each group electrons are in shells further from the nucleus Ionic Size Increases down the group The size of positive ions is less than the original atom because the nuclear charge exceeds the electronic charge. Melting Points Decrease down each group

metallic bonding gets weaker due to

increased size Each atom contributes two electrons to the delocalised cloud. Melting points tend not to give a decent trend as different crystalline structures affect the melting point.

CHEMICAL PROPERTIES OF THE ELEMENTS Overall Reactivity increases down the Group due to the ease of cation formation Oxygen • Water •

react with increasing vigour down the group react with increasing vigour down the group

OXIDES OF GROUP II ELEMENTS Properties • ionic solids; EXC. beryllium oxide which has covalent character Hydroxides • basic strength also increases down group • this is because the solubility increases • the metal ions get larger so charge density decreases • there is a lower attraction between the OH¯ ions and larger dipositive ions 145

• the ions will split away from each other more easily • there will be a greater concentration of OH¯ ions in water Uses of hydroxides Ca(OH) • used in agriculture to neutralise acid soils CARBONATES Properties • insoluble in water • undergo thermal decomposition to oxide and carbon dioxide We note first that the elements are all electropositive, having relatively low ionisation energies, and are, in consequence, very reactive. The enthalpy change required for the process M(metal) -»M + (g) for Group I, or M(metal) -> M2+(g) for Group II is at a maximum at the top of each group, and it is, therefore, not surprising to find that lithium, beryllium and, to some extent, magnesium do form some covalent compounds. Most solid compounds of Group 1 and II elements, however, have ionic structures and the properties associated with such structures—high m.p. and b.p., solubility in water rather than in organic solvents and electrical conductance when molten. IONS IN SOLUTION The hydration energies (strictly, hydration enthalpies) fall, as expected, as we descend either Group, and are larger for Group II than for Group I ions. The solubilities of the salts of Groups I and II are determined by a balance between lattice energy, hydration energy and the entropy change in going from solid to solution, and only a few generalisations are possible. Thus high charge and low ionic radii tend to produce insolubility (for example salts of lithium, beryllium and magnesium, especially those with doubly charged anions such as carbonate COa~). At the other end of the scale, low charge and large radii also produce low solubility (for example salts of potassium, rubidium and caesium containing large anions such as the tetraphenylborate anion (p. 136). In between, solubility is the rule for all Group I salts, and for most Group II salts containing singly-charged negative ions; for many Group II salts with doublyor triply-charged anions (for example COj", SOj", PO^ )

146

insolubility is often observed. The decreasing tendency to form salts with water of crystallization (as a group is descended) is again in line with the falling hydration energy. For example, both sodium sulphate and carbonate form hydrates but neither of the corresponding potassium salts do; the sulphates of Group II elements show a similar trend MgSO4 , 7H2O, CaSO4 . 2H2O, BaSO4. For the most part, however, the chemistry of the Group I and II elements is that of the metal and the ions M + for Group I and M2* for Group II. As already noted the two head elements, lithium and beryllium, tend to form covalent compounds; the beryllium ion Be2 + , because of its very small radius and double charge, has also some peculiar properties in solution, which are examined later OCCURRENCE AND EXTRACTION Of the Group II metals (beryllium to barium) beryllium, the rarest, occurs as the aluminatesilicate, beryl \magnesium is found as the carbonate and (with calcium) as the double carbonate dolomite', calcium, strontium and barium all occur as carbonates, calcium carbonate being very plentiful as limestone. The general characteristics of all these elements generally preclude their extraction by any method involving aqueous solution. For the lighter, less volatile metals (Li, Na, Be, Mg, Ca) electrolysis of a fused salt (usually the chloride), or of a mixture of salts, is used. The heavier, more volatile metals in each group can all be similarly obtained by electrolysis, but it is usually more convenient to take advantage of their volatility and obtain them from their oxides or chlorides by displacement, i.e. by general reactions such as 3M2O + 2Mm -* M2 mO3 4- 6M| MCI + M1 ~» M!C1 + M| Thus potassium is obtained by heating potassium chloride with sodium, and barium by reduction of barium oxide with aluminium. Sodium is important in many technical processes and is therefore prepared in considerable quantity. Almost all of it is now made by electrolysis of the fused sodium chloride, using the Downs cell. The graphite anode is cylindrical and is surrounded by the steel gauze diaphragm and the

147

concentric cylindrical cathode (also of steel). The electrolyte is usually a mixture of sodium chloride and calcium chloride; the latter is added to reduce the m.p. of the sodium chloride to approximately 800 K. (Some calcium is therefore liberated with the sodium.) The gap between anode and cathode is kept as small as possible to reduce resistance: the heat developed by the current maintains the temperature of the cell. Chlorine is set free at the anode surface, rises into the nickel cone and can be collected. Sodium, liberated at the cathode, is prevented by the diaphragm from ing into the anode region; the molten sodium collects under the circular hood and rises up the pipe, being assisted if necessary by the puddle-rod. The calcium, being almost immiscible with sodium and much more dense, can readily be separated from the molten sodium. The graphite anode wears away and must be renewed from time to time. Element

Be

Mg

Ca

Reaction

Does not react with with water,

Very slowl with water, readily with

React with cold water. vigour of reaction increasing.

Conditions

Sr

Ba

steam Basic properties

Be(OH)2

MgO

amphoteric

insoluble

Slightly soluble

M

soluble

of product Base strength increasing

USES Beryllium is added to copper to produce an alloy with greatly increased wear resistance; it is used for current-carrying springs and non-sparking safety tools. It is also used as a neutron and reflector in nuclear reactors. Much magnesium 148

is used to prepare light metal alloys; other uses include the extraction of titanium and in the removal of oxygen and sulphur from steels; calcium finds a similar use. BIOLOGICAL IMPORTANCE Sodium and potassium ions are found in all animal cells and, usually, the concentration of potassium ions inside the cell is greater than that of sodium. In many cells, this concentration difference is maintained by a 'sodium pump', a process for which the energy is supplied by the hydrolysis of adenosine triphosphate (ATP). Diffusion of excess potassium ions outwards through the cell wall gives the inside of the cell a net negative charge (due to the anions present) and a potential difference is established across the cell wall. In a nerve cell, a momentary change in the permeability of the cell wall to sodium ions can reverse the sign of this potential difference, and this produces the electrical impulse associated with the action of the nerve. The ability of living organisms to differentiate between the chemically similar sodium and potassium ions must depend upon some difference between these two ions in aqueous solution. Essentially, this difference is one of size of the hydrated ions, which in turn means a difference in the force of electrostatic (coulombic) attraction between the hydrated cation and a negatively-charged site in the cell membrane; thus a site may be able to accept the smaller ion Na+(aq) and reject the larger K+ (aq). This same mechanism of selectivity operates in other 'ion-selection' processes, notably in ionexchange resins. All organisms seem to have an absolute need for magnesium. In plants, the magnesium complex chlorophyll is the prime agent in photosynthesis. In animals, magnesium functions as an enzyme activator; the enzyme which catalyses the ATP hydrolysis mentioned above is an important example. Calcium plays an important part in structure-building in living organisms, perhaps mainly because of its ability to link together phosphate-containing materials. Calcium ions in the cell play a vital part in muscle contraction. THE HYDRIDES All Group I and II elements, except beryllium, form hydrides by direct combination with hydrogen. The hydrides of the metals except those of beryllium and magnesium, are white mainly ionic solids, all Group I hydrides having the sodium chloride lattice 149

structure. All the hydrides are stable in dry air but react with water, the vigour of the reaction increasing with the molecular weight of the hydride for any particular group. Group II metals also form halides by direct combination. The trends in heat of formation and m.p., however, whilst following the general pattern of the corresponding Group I compounds, are not so regular. * Lithium bromide and iodide probably have some degree of covalency but this does not affect the general conclusion. As a consequence of the high ionisation energy of beryllium its halides are essentially covalent, with comparatively low m.p., the melts being non-conducting and (except beryllium fluoride) dissolving in many organic solvents. The lower in Group II form essentially ionic halides, with magnesium having intermediate properties, and both magnesium bromide and iodide dissolve in organic solvents. The lattice energies of the Group II fluorides are generally greater than those for the corresponding Group I fluorides; consequently all but beryllium fluoride are insoluble. (The solubility of beryllium fluoride is explained by the high hydration energy of the beryllium ion, cf. LiF.) The high hydration energy of the Be2+ ion* results in hydrolysis in neutral or alkaline aqueous solution; in this reaction the beryllium halides closely resemble the aluminium halides. The magnesium ion having a high hydration energy also shows hydrolysis but to a lesser extent, The chloride forms several hydrates which decompose on heating to give a basic salt. Other Group II halides are essentially ionic and therefore have relatively high m.p., the melts acting as conductors, and they are soluble in water but not in organic solvents. a. Beryllium

Beryllium was discovered by Louis-Nicholas Vauquelin in 1798. Vauquelin found beryllia (BeO) in emeralds and in the mineral beryl (beryllium aluminum cyclosilicate). Beryllium was first isolated by Friederich Wöhler in 1828. Wöhler reacted potassium with beryllium chloride in a platinum crucible yielding potassium chloride and beryllium.

150

Unlike most metals, beryllium is virtually transparent to x-rays and hence it is used in radiation windows for x-ray tubes. Beryllium alloys are used in the aerospace industry as light-weight materials for high performance aircraft, satellites and spacecraft. Beryllium is also used in nuclear reactors as a reflector and absorber of neutrons, a shield and a . Beryllium is used as an alloying agent in producing beryllium copper, which is extensively used for springs, electrical s, spot-welding electrodes, and nonsparking tools. It is applied as a structural material for high-speed aircraft, missiles, spacecraft, and communication satellites. Other uses include windshield frame, brake discs, beams, and other structural components of the space shuttle. Because beryllium is relatively transparent to X-rays, ultra-thin Be-foil is finding use in X-ray lithography for reproduction of micro-miniature integrated circuits. Beryllium is used in nuclear reactors as a reflector or for it has a low thermal neutron absorption cross section. It is used in gyroscopes, computer parts, and instruments where lightness, stiffness, and dimensional stability are required. The oxide has a very high melting point and is also used in nuclear work and ceramic applications.

Beryllium is found in some 30 mineral species, the most important of which are bertrandite, beryl, chrysoberyl, and phenacite. Aquamarine and emerald are precious forms of beryl. Beryl and bertrandite are the most important commercial sources of the element and its compounds. Most of the metal is now prepared by reducing beryllium fluoride with magnesium metal. Beryllium metal did not become readily available to industry until 1957. The metal, steel gray in color, has many desirable properties. As one of the lightest of all metals, it has one of the highest melting points of the light metals. Its modulus of elasticity is about one third greater than that of steel. It resists attack by 151

concentrated nitric acid, has excellent thermal conductivity, and is nonmagnetic. It has a high permeability to X-rays and when bombarded by alpha particles, as from radium or polonium, neutrons are produced in the amount of about 30 neutrons/million alpha particles. At ordinary temperatures, beryllium resists oxidation in air, although its ability to scratch glass is probably due to the formation of a thin layer of the oxide. Beryllium and its salts are toxic and should be handled with the greatest of care. Beryllium and its compounds should not be tasted to the sweetish nature of beryllium (as did early experimenters). The metal, its alloys, and its salts can be handled if certain work codes are observed, but no attempt should be made to work with beryllium before becoming familiar with proper safeguards. b. Magnesium

Although it is the eighth most abundant element in the universe and the seventh most abundant element in the earth's crust, magnesium is never found free in nature. Magnesium was first isolated by Sir Humphry Davy, an English chemist, through the electrolysis of a mixture of magnesium oxide (MgO) and mercuric oxide (HgO) in 1808. Today, magnesium can be extracted from the minerals dolomite (CaCO3·MgCO3) and carnallite (KCl·MgCl2·6H2O), but is most often obtained from seawater. Every cubic kilometer of seawater contains about 1.3 billion kilograms of magnesium (12 billion pounds per cubic mile). 152

Magnesium burns with a brilliant white light and is used in pyrotechnics, flares and photographic flashbulbs. Magnesium is the lightest metal that can be used to build things, although its use as a structural material is limited since it burns at relatively low temperatures. Magnesium is frequently alloyed with aluminum, which makes aluminum easier to roll, extrude and weld. Magnesium-aluminum alloys are used where strong, lightweight materials are required, such as in airplanes, missiles and rockets. Cameras, horseshoes, baseball catchers' masks and snowshoes are other items that are made from magnesium alloys. Magnesium oxide (MgO), also known as magnesia, is the second most abundant compound in the earth's crust. Magnesium oxide is used in some antacids, in making crucibles and insulating materials, in refining some metals from their ores and in some types of cements. When combined with water (H 2O), magnesia forms magnesium hydroxide (Mg(OH)2), better known as milk of magnesia, which is commonly used as an antacid and as a laxative. Hydrated magnesium sulphate (MgSO4·7H2O), better known as Epsom salt, was discovered in 1618 by a farmer in Epsom, England, when his cows refused to drink the water from a certain mineral well. He tasted the water and found that it tasted very bitter. He also noticed that it helped heal scratches and rashes on his skin. Epsom salt is still used today to treat minor skin abrasions. Other magnesium compounds include magnesium carbonate (MgCO3) and magnesium fluoride (MgF2). Magnesium carbonate is used to make some types of paints and inks and is added to table salt to prevent caking. A thin film of magnesium fluoride is applied to optical lenses to help reduce glare and reflections.

c. Calsium Calcium is essential for human nutrition. Animals skeletons get their rigidity primarily from calcium phosphate. The eggs of birds and shells of mollusks are comprised of

153

calcium carbonate. Calcium is also necessary for plant growth. Calcium is used as a reducing agent when preparing metals from their halogen and oxygen compounds; as a reagent in purification of inert gases; to fix atmospheric nitrogen; as a scavenger and decarbonizer in metallurgy; and for making alloys. Calcium compounds are used in making lime, bricks, cement, glass, paint, paper, sugar, glazes, as well as for many other uses. Sources: The Romans prepared lime (called calx) in the first century, but the metal was not discovered until 1808. Berzelius and Pontin prepared calcium amalgam by electrolyzing lime in mercury. Davy isolated the impure metal. The metal may be prepared by electrolysis of CaCl2 at a temperature slightly above its melting point. Calcium is the fifth most abundant element in the earth's crust, making up 3.22% of the earth, air, and oceans. Natural forms of calcium include limestone (CaCO3), gypsum (CaSO4·2H2O), and fluorite (CaF2). Apatite is the fluorophosphate or chlorophosphate of calcium. d. Stronsium Strontium was recognized as distinct from barium in 1790 by Adair Crawford in a mineral sample from a mine near Strontian, Scotland. The element took its name from the Scottish town. The metal was first isolated by Sir Humphry Davy in 1808, by electrolysis. Strontium is a soft, silvery metal. When cut it quickly turns a yellowish color due to the formation of strontium oxide (strontia, SrO) . Finely powdered strontium metal is sufficiently reactive to ignite spontaneously in air. It reacts with water quickly (but not violently like the Group 1 metals) to produce strontium hydroxide and hydrogen gas. Strontium and its compounds burn with a crimson flame and are used in fireworks. Uses: Strontium is used for producing glass (cathode ray tubes) for color televisions. It is also used in producing ferrite ceramic magnets and in refining zinc. The world's most accurate atomic clock, accurate to one second in 200 million years, has been developed using strontium atoms. Strontium salts are used in flares and fireworks for a crimson color.

Strontium chloride is used in toothpaste for sensitive teeth.

Strontium oxide is used to improve the quality of pottery glazes. The isotope

90

Sr is 154

one of the best long-lived, high-energy beta emitters known. It is used in cancer therapy.

e. Barium Uses: Barium is used as a "flashed getter" in vacuum tubes to remove the last traces of gases. Barium is an important element in yttrium barium copper oxide (YBCO) superconductors. An alloy of barium with nickel is used in sparkplug wire. Hardness: 1.25 mohs Harmful effects: Barium compounds that are water or acid soluble are highly poisonous. Barium powder can ignite spontaneously in air. Characteristics: Barium is a metallic element chemically resembling calcium but more reactive. Barium oxidizes very easily in air and is also highly reactive with water or alcohol. Barium is most commonly found as the mineral barite (BaSO4) and witherite (BaCO3) f. Radium Radium was discovered in 1898 by Marie S. Curie and her husband Pierre in pitchblende (mainly uranium dioxide UO2). If pitchblende contains 50 percent uranium oxides, about eight tons of it is needed to extract 1 gram of radium. Already that year the Curies had discovered polonium, a radioactive element whose properties they said were similar to bismuth's. Now, in radium - in the form of radium bromide - they had discovered a further radioactive element whose chemistry was very similar to that of group II metal barium. Metallic radium was first isolated in 1910 by Marie S. Curie and Andre Debierne by the electrolysis of a solution of pure radium chloride. The element's name comes from the Latin word 'radius', meaning ray, after the rays emitted by this radioactive element. In the discovery of radioactivity, chemists realized that one of alchemy's dreams - the transmutation of elements - was possible. Harmful effects: Radium is highly radioactive and hence carcinogenic. Microscopic quantities of radium in the environment can lead to some accumulation of radium in bone tissue. Radium, like calcium, is a group II element and our bodies treat it in a similar way. Characteristics: Radium is a silvery-white metal. It is highly radioactive and its decay product, radon 155

gas, is also radioactive. One result of radium's intense radioactivity is that the metal and its compounds glow in the dark. When it is exposed to air, it reacts with nitrogen to quickly form a black coating of radium nitride. Radium's chemistry is similar to that of the other alkali earth metals. It reacts very vigorously with water to form hydrogen gas and radium hydroxide. It reacts with even more vigorously with hydrochloric acid to form radium chloride. Uses: Radium was used in the production of luminous paints, but this is now considered too dangerous. Radium chloride was used medicinally to produce radon gas for cancer treatment. Safer treatments are now available.

156

CHAPTER II IIB ELEMENTS II (ZINC), CIUM, MERCURY These elements formed Group IIB of Mendeleef 's original periodic table. As we have seen in Chapter 13, zinc does not show very marked 'transition-metal' characteristics. The other two elements in this group, cium and mercury, lie at the ends of the second and third transition series (Y-Cd, La-Hg) and, although they resemble zinc in some respects in showing a predominantly 4- 2 oxidation state, they also show rather more transition-metal characteristics. Additionally, mercury has characteristics, some of which relate it quite closely to its immediate predecessors in the third transition series, platinum and gold, and some of which are decidedly peculiar to mercury. PROPERTIES OF THE ELEMENTS Elem

Atomi

ent

Outer electrons

Atomi

Density

c

c

(g/

numb

radius

er

(nm)

m.p. b.p. lonisation

) (K)

(K)

energy*

atomisation

(kJ/mol)

(kJ/mol)

1st Zn

30

[Ar]3

Cd

48

[Kr]4

Hg

80

[Ar]4

4

5

0.133

7.13

693

0.149

8.65

594

0,152

13,53

234

Heat of

2nd

118

906

1

1734

103

816

8

1630

630

1007 1809

131 286 61

157

The table shows that all three elements have remarkably low melting points and boiling points—an indication of the weak metallic ^bonding, especially notable in mercury. The low heat of atomisation of the latter element compensates to some extent its higher ionisation energies, so that, in practice, all the elements of this group can form cations M2 + in aqueous solution or in hydrated salts; anhydrous mercury(II) compounds are generally covalent. ZINC Isotopes: There are 21 known isotopes of zinc, 5 stable and 16 unstable. Natural zinc contains the 5 stable isotopes. Properties: Zinc has a melting point of 419.58°C, boiling point of 907°C, specific gravity of 7.133 (25°C), with a valence of 2. Zinc is a lustous blue-white metal. It is brittle at low temperatures, but becomes malleable at 100-150°C. It is a fair electrical conductor. Zinc burns in air at high red heat, evolving white clouds of zinc oxide.

Uses: Zinc is used to form numerous alloys, including brass, bronze, nickel silver, soft solder, Geman silver, spring brass, and aluminum solder. Zinc is used to make die castings for use in the electrical, automotive, and hardware industries. The alloy Prestal, consisting of 78% zinc and 22% aluminum, is nearly as strong as steel yet exhibits superplasticity. Zinc is used to galvanize other metals to prevent corrosion. Zinc oxide is used in paints, rubbers, cosmetics, plastics, inks, soap, batteries, pharmaceuticals, and many other products. Other zinc compounds are also widely used, such as zinc sulfide (luminous dials and fluorescent lights) and ZrZn2 (ferromagnetic materials). Zinc is an essential element for humans and other animal nutrition. Zinc-deficient animals require 50% more food to gain the same weight as animals with sufficient zinc. Zinc metal is not considered toxic, but if fresh zinc oxide is inhaled it can cause a disorder referred to as zinc chills or oxide shakes. Sources: The primary ores of zinc are sphalerite or blende (zinc sulfide), smithsonite (zinc carbonate), calamine (zinc silicate), and franklinite (zinc, iron, and manganese oxides). An old method of producing zinc was by reducing calamine with charcoal. More recently, it has been obtained by roasting the ores to form zinc oxide and then reducing the oxide with carbon or coal, followed by distillation of the metal.

158

CIUM THE ELEMENT Cium is usually found in zinc ores and is extracted from them along with zinc (p. 416); it may be separated from the zinc by distillation (cium is more volatile than zinc) or by electrolytic deposition. Cium is a soft metal, which forms a protective coating in air, and bums only on strong heating to give the brown oxide CdO. It dissolves in acids with evolution of hydrogen. It is used as a protective agent, particularly for iron, and is more resistant to corrosion by sea water than, for example, zinc or nickel. In its chemistry, cium exhibits exclusively the oxidation state -f 2 in both ionic and covalent compounds. The hydroxide is soluble in acids to give

cium(II)

salts,

and

slightly

soluble

in

concentrated

alkali

where

hydroxociates are probably formed; it is therefore slightly amphoteric. It is also soluble in ammonia to give ammines, for example [Cd(NH3)4]2+. Of the halides, ciumill) chloride is soluble in water, but besides [Cd(H2O)J2 + ions, complex species [CdCl]*, [CdQ3]~ and the undissociated chloride [CdCl2] exist in the solution, and addition of chloride ion increases the concentrations of these chloro-complexes at the expense of Cd2+(aq) ions. Solid cium(II) iodide CdI2 has a layer lattice' — a structure intermediate between one containing Cd2* and I~ ions and one containing CdI2 molecules — and this on vaporisation gives linear, covalent I — Cd — I molecules. In solution, iodo-complexes exist, Cium(ll} sulphide, CdS, is a canary-yellow solid, precipitated by addition of hydrogen sulphide (or sulphide ion) to an acid solution of a cium(II) salt; presence of chloride ion may reduce the concentration of Cd2+ (aq) sufficiently to prevent precipitation. Complexes of cium include, besides those already mentioned, a tetracyanociate [Cd(CN)4]2~ which in neutral solution is sufficiently unstable to allow precipitation of cium(II) sulphide by hydrogen sulphide. Octahedral [CdCl6]4" ions are known in the solid state, as, for example, K4CdCl6. 159

MERCURY THE ELEMENT Mercury has been known for many centuries, perhaps because its extraction is easy; it has an almost unique appearance, it readily displaces gold from its ores and it forms amalgams with many other metals—all properties which caused the alchemists to regard it as one of the "fundamental' substances. It occurs chiefly as cinnabar, the red sulphide HgS, from which it is readily extracted either by roasting (to give the metal and sulphur dioxide) or by heating with calcium oxide; the metal distils off and can be purifyied by vacuum distillation. Mercury has a large relative atomic mass, but, like zinc and cium, the bonds in the metal are not strong. These two factors together may for the very low melting point and boiling point of mercury. The low boiling point means that mercury has an appreciable vapour pressure at room temperature; 1 m3 of air in equilibrium with the metal contains 14 mg of vapour, and the latter is highly toxic. Exposure of mercury metal to any reagent which produces volatile mercury compounds enhances the toxicity. The metal is slowly oxidised by air at its boiling point, to give red mercury(II) oxide; it is attacked by the halogens (which cannot therefore be collected over mercury) and by nitric acid. (The reactivity of mercury towards acids is further considered on pp. 436, 438.) It forms amalgams— liquid or solid—with many other metals; these find uses as reducing agents (for example with sodium, zinc) and as dental fillings (for example with silver, tin or copper). USES Mercury is extensively used in various pieces of scientific apparatus, such as thermometers, barometers, high vacuum pumps, mercury lamps, standard cells (for example the Weston cell), and so on. The metal is used as the cathode in the KellnerSolvay cell (p. 130). Mercury compounds (for example mercury(II) chloride) are used in medicine because of their antiseptic character. The artificial red mercury(II) sulphide is the artist's 'vermilion1. Mercury(II) sulphate is a catalyst in the manufacture of ethanal from ethyne: C2H2 + H2O ^^ CH3. CHO COMPOUNDS OF MERCURY The chemistry of mercury compounds is complicated by the 160

equilibrium The relevant redox potentials are : Hg2+(aq) 4- 2e~ -> Hg(I) : E^ = 0.85 V + 2e~ -> 2Hg(I) : E^ = 0.79 V Hence mercury is a poor reducing agent; it is unlikely to be attacked by acids unless these have oxidising properties (for example nitric acid), or unless the acid anion has the power to form complexes with one or both mercury cations Hg2+ or Hgf +, so altering the E^ values. Nitric acid attacks mercury, oxidising it to Hg2+(aq) when the acid is concentrated and in excess, and to Hg2+(aq) when mercury is in excess and the acid dilute. Hydriodic acid HI(aq) attacks mercury, because mercury(II) readily forms iodo-complexes. Oxidation state +1 The mercury(I) ion has the structure so that each mercury atom is losing one electron and sharing one electron, i.e. is 'using' two valency electrons. The existence of Hg| + has been established by experiments in solution and by X-ray diffraction analysis of crystals of mercury(I) chloride, Hg2Cl2 where the mercury ions are in pairs with the chloride ions adjacent, i.e. *Hg—Hg+. Cl~. (It is now known that mercury can also form species Hg^ up to Hgg+ ; cium also gives Cd^+, and other polymetallic cations, for example Bi^ are known.) The ion Hg|+(aq) tends to disproportionate, especially if the concentration of Hg2 +(aq) is reduced, for example by precipitation or by complex formation. However, the equilibrium can be moved to the left by using excess of mercury, or by avoiding aqueous solution. Thus, heating a mixture of mercury and solid mercury(II) chloride gives mercury(I) chloride, which sublimes off: Hg + HgCl2 -> Hg2Cl2 The product, commonly called calomel, is a white solid, insoluble in water; in its reactions (as expected) it shows a tendency to produce mercury(II) and mercury. Thus under the action of light, the substance darkens because mercury is formed; addition of aqueous ammonia produces the substance H2N—Hg—Hg—Cl, but this also darkens on standing, giving H2N—Hg—Cl and a black deposit of mercury. Mercury(I) ions can be produced in solution by dissolving excess mercury in dilute nitric acid: 6Hg + 8H+ + 2NO3~ -» 3Hg|+ + 2NO + 4H2O From the acid solution white hydrated mercury(I) nitrate Hg2(NO3)2.2H2O 161

can be crystallised out; this contains the ion [H2O-Hg-Hg-H2O]2 + which is acidic (due to hydrolysis) in aqueous solution. Addition of chloride ion precipitates mercury(I) chloride. Oxidation state + 2 Mercury(II) oxide, HgO, occurs in both yellow and red forms; the yellow form is precipitated by addition of hydroxide ion to a solution containing mercury(II) ions, and becomes red on heating. Mercury(II) oxide loses oxygen on heating. Mercury(II) chloride is obtained in solution by dissolving mercury(II) oxide in hydrochloric acid; the white solid is obtained as a sublimate by heating mercury(II) sulphate and solid sodium chloride: HgSO4 + 2NaCl -» HgCl2 + Na2SO4 The aqueous solution has a low conductivity, indicating that mercury(II) chloride dissolves essentially as molecules Cl—Hg—Cl and these linear molecules are found in the solid and vapour. A solution of mercury(II) chloride is readily reduced, for example by tin(II) chloride, to give first white insoluble mercury(I) chloride and then a black metallic deposit of mercury. The complexes formed from mercury(II) chloride are considered below. Mercury(H) iodide, HgI2, is coloured either red or yellow, and is precipitated (yellow, turning red) by adding the stoichiometric amount of iodide ion to a solution containing mercury(II): Hg2+ +2r-»HgI2 Addition of excess iodide gives a complex (see below). Mercury(II) sulphate and nitrate are each obtained by dissolving mercury in the appropriate hot concentrated acid; the sulphate is used as a catalyst. MercuryiH) sulphide, HgS, again appears in two forms, red (found naturally as cinnabar) and black, as precipitated by hydrogen sulphide from a solution containing Hg(II) ions. Complexes Mercury (I) forms few complexes, one example is the linear [H2O- Hg-Hg—H2O]2 + found in the mercury(I) nitrate dehydrate (above, p. 437). In contrast, mercury(II) forms a wide variety of complexes, with some peculiarities: (a) octahedral complexes are rare, (b) complexes with nitrogen as the donor atom are common, (c) complexes are more readily formed with iodine than with other halogen ligands. Mercury(II)

162

halides, HgX2, can be regarded as neutral, 2- co-ordinate linear complexes X—Hg- X. X is readily replaced; addition of ammonia to a solution of mercury(II) chloride gives a white precipitate NH2—Hg—Cl; in the presence of concentrated ammonium chloride, the same reagents yield the diamminomercury(II) cation, [NH3—Hg—NH3]2+, which precipitates as [Hg(NH3)2]Cl2. In presence of excess chloride ion, mercury(II) chloride gives complexes [HgCl3]~ and [HgCl4]2~, but the corresponding iodo-complex [HgI4]2", from mercury(II) iodide and excess iodide, is more stable. (It is rare for iodocomplexes to form at all and very rare to find them with stabilities greater than those of thloro-complexes.) In both solid HgI2 and the complex [HgI4]2~the mercury is tetrahedrally 4-co-ordinated. The [HgI4]2" ion has a characteristic reaction with ammonia—a trace produces a yellow colour and more ammonia gives a brown precipitate. (An alkaline solution containing [HgI4]2~ ions is therefore used as a test for ammonia; it is sometimes called Messier's reagent.) Insoluble salts of the anion [HgI4]2~ are known, for example Cu2[HgI4] (red)

163

CHAPTER II MATTER ELEMENTS OF GROUP III A BORON (B)

Physical Information 1. Atomic Number

:5

2. Relative Atomic Mass (12C=12.000)

: 10.81

3. Melting Point/K

: 2573

4. Boiling Point/K

: 3931

5. Density/kg m-3

: 2340 (293K)

6. Liquid range

: 1851 K

7. Ground State Electron Configuration

: [He]2s2 2p1

8. Electron Affinity (M-M-)/kJ mol-1

: -15

Discovery Boron compounds have been known for thousands of years, but the element was not isolated until 1808 by Sir Humphry Davy, Joseph-Louis Gay-Lussac (1778-1850) and Louis Jaques Thenard (1777-1857) in London. This was accomplished through the reaction of boric acid (H3BO3) with potassium. Appearance The element is a grey powder, but is not found free in nature. Source

164

The element is not found free in nature, but occurs as orthoboric acid usually found in certain volcanic spring waters and as borates in boron and colemantie. Ulexite, another boron mineral, is interesting as it is nature's own version of "fiber optics." Important sources of boron are ore rasorite (kernite) and tincal (borax ore). Both of these ores are found in the Mojave Desert. Tincal is the most important source of boron from the Mojave. Extensive borax deposits are also found in Turkey. Boron exists naturally as 19.78%

B isotope and 80.22%

10

B isotope. High-purity

11

crystalline boron may be prepared by the vapor phase reduction of boron trichloride or tribromide with hydrogen on electrically heated filaments. The impure or amorphous, boron, a brownish-black powder, can be obtained by heating the trioxide with magnesium powder. Boron of 99.9999% purity has been produced and is available commercially. Elemental boron has an energy band gap of 1.50 to 1.56 eV, which is higher than that of either silicon or germanium. Uses Amorphous boron is used in pyrotechnic flares to provide a distinctive green colour, and in rockets as an igniter. The most important compounds of boron are boric (or boracic) acid, widely used as a mild antiseptic, and borax which serves as a cleansing flux in welding and as a water softener in washing powders. Boron compounds are also extensively used in the manufacture of borosilicate glasses. Other boron compounds show promise in treatingarthritis. The isotope boron 10 is used as a control for nuclear reactors, as a shield for nuclear radiation, and in instruments used for detecting neutrons. Demand is increasing for boron filaments, a high-strength, low-density material chiefly employed for advanced aerospace structures. Boron used as a deoxidiser and flux in metallurgy, in the manufacture of BoroSilicate Glass, in the bearings of turbines the form of metallic borides where hardness and resistance to corrosion is required, and in the nuclear industry as a for neutrons. The following uses for boron are gathered from a number of sources as well as from anecdotal comments. I'd be delighted to receive corrections as well as additional referenced uses (please use the mechanism to add uses). •

amorphous boron is used in pyrotechnic flares (distinctive green colour), and rockets (as an igniter) 165

•

boric, or boracic, acid, is used as a mild antiseptic

•

borax, Na2B4O7.10H2O, is a cleansing flux in welding

•

borax, Na2B4O7.10H2O is a water softener in washing powders

•

boron compounds are used in production of enamels for covering steel of refrigerators, washing machines, etc.

•

boron compounds are extensively used in the manufacture of enamels and borosilicate glasses

•

boron compounds show promise in treating arthritis

•

10

B is used as a control for nuclear reactors, as a shield for nuclear radiation, and in

instruments used for detecting neutrons •

boron nitride is as hard as diamond. It behaves like an electrical insulator, but conducts heat like a metal. It also has lubricating properties similar to graphite

•

the hydrides are sometimes used as rocket fuels

•

boron filaments, a high-strength, lightweight material, are used for advanced aerospace structures, .

•

lightweight compounds used for aerospace structures

•

boron filaments used in fibre optics research

•

Boric Acid is also used in North America for the control of cockroaches, silverfish, ants, fleas, and other insects.

Biological Role Elemental boron is not considered a poison, and indeed is essential to plants, but assimilation of its compounds has a cumulative toxic effect. General Information Elemental boron has an energy band gap of 1.50 to 1 .56 eV, which is higher than that of either silicon or germanium. It has interesting optical characteristics, transmitting portions of the infrared only. It is a poor conductor of electricity at room temperature, but a good conductor at high temperatures. Boron in its crystalline form is a very hard solid with a quaesi-metallic sheen, and in its amorphous form is a brown powder. Reactions •

Boron burns with a brilliant flame in oxygen to form boron trioxide.

•

Boron burns in air when heated to give a mixture of boron trioxide and Boron Nitride.

166

•

Boron is relatively inert and must be in a highly divided state to react with acids or alkalis.

•

Boron is oxidised by nitric acid to boric acid.

•

Boron reacts with fused sodium hydroxide to form sodium borate and hydrogen.

Reaction of boron with air The behaviour of boron to air depends upon the crystallinity of the sample, temperature, particle size, and purity. By and large, boron does not react with air at room temperature. At higher temperatures, boron does burn to form boron (III) oxide, B2O3. 4B + 3O2(g) → 2B2O3(s) Reaction of boron with water Boron does not react with water under normal conditions. Reaction of boron with the halogens Boron reacts vigorously with the halogens fluorine, F2, chlorine, Cl2, bromine, Br2 to to form the trihalides boron(III) fluoride, BF3, boron(III) chloride, BCl3, and boron(III) bromide, BBr3 respectively. 2B(s) + 3F2(g) → 2BF3(g) 2B(s) + 3Cl2(g) → 2BCl3(g) 2B(s) + 3Br2(g) → 2BF3(l) Reaction of boron with acids Crystalline boron does not react with boiling hydrochloric acid, HCl, or boiling hydrofluoric acid, HF. Powdered boron oxidizes slowly when treated with concentrated nitric acid, HNO3.

Detection and Analysis

167

Boron is detected by converting the material under analysis to Borax by heating with concentrated nitric acid and then heating with concentrated sulphuric acid and ethanol to form ethyl borate, which burns with a green flame. Handling Elemental boron and the borates are not considered to be toxic, and they do not require special care in handling. However, some of the more exotic boron hydrogen compounds are definitely toxic and do require care. Isolation It is not normally necessary to make boron in the laboratory and it would normally be purchased as it is available commercially. The most common sources of boron are tourmaline, borax [Na2B4O5(OH)4.8H2O], and kernite [Na2B4O5(OH)4.2H2O]. It is difficult to obtain pure. It can be made through the magnesium reduction of the oxide, B2O3. The oxide is made by melting boric acid, B(OH)3, which in turn is obtained from borax. B2O3 + 3Mg → 2B + 3MgO Samm amounts of high purity boron are available through the thermal decomposition of compounds such as BBr3 with hydrogen gas using a heated tantalum wire. Results are better with hot wires at tmeperatures over 1000°C.

ALUMUNIUM (Al)

Physical Information 1. Atomic Number

: 13 168

2. Relative Atomic Mass (12C=12.000)

: 26.982

3. Melting Point/K

: 933

4. Boiling Point/K

: 2740

5. Density/kg m-3

: 2698 (293K)

6. Ground State Electron Configuration

: [Ne ]3s2 3p1

7. Electron Affinity (M-M-)/kJ mol-1

: -44

Discovery (L. alumen: alum) The ancient Greeks and Romans used alum as an astringent and as a mordant in dyeing. In 1761 de Morveau proposed the name alumine for the base in alum, and Lavoisier, in 1787, thought this to be the oxide of a still undiscovered metal. Wohler is generally credited with having isolated the metal in 1827, although an impure form was prepared by Oersted two years earlier. In 1807, Davy proposed the name aluminium for the metal, undiscovered at that time, and later agreed to change it to aluminum. Shortly thereafter, the name aluminum was adopted to conform with the "ium" ending of most elements. Aluminium was also the accepted spelling in the U.S. until 1925, at which time the American Chemical Society decided to use the name aluminum thereafter in their publications. Appearance Aluminium is a hard and strong, silvery-white metal. An oxide film prevents it from reacting with air and water. Occurrence Aluminium is highly reactive and does not occur in the free state. However, it is widely distributed and it is third in abundance on earth after Oxygen and Silicon. Aluminium exists primarily as Alumino-Silicates (i.e. as Felspar, NaAlSi3O8, or KAlSi3O8, or CaAl2Si2O8), in igneous rocks and as Clays, H4Al2Si2O9, in sedimentary rocks. Aluminium has three principal ores •

Gibbsite or Hydrargillite, Al2O3.3H2O,

•

Bauxite, Al2O3.2H2O, Diaspore, Al2O3.H2O, and

•

Cryolite, AlF3.3HF. Aluminium also occurs in the form of its Aluminium Oxide, Al2O3, in the semiprecious

stones. The colouration in these gems is caused by trace quantities of impurities : 169

•

Emerald, Sapphire (coloured blue by Cobalt Oxide, CoO), and Ruby (coloured blue by Chromium Oxide, Cr2O3).

Source The method of obtaining aluminum metal by the electrolysis of alumina dissolved in cryolite was discovered in 1886 by Hall in the U.S. and at about the same time by Heroult in . Cryolite, a natural ore found in Greenland, is no longer widely used in commercial production, but has been replaced by an artificial mixture of sodium, aluminum, and calcium fluorides. Aluminum can now be produced from clay, but the process is not economically feasible at present. Aluminum is the most abundant metal to be found in the earth's crust (8.1%), but is never found free in nature. In addition to the minerals mentioned above, it is also found in granite and in many other common minerals. Most commercially produced aluminium is obtained by the Bayer process of refining bauxite. In this process the bauxite is refined to pure aluminium oxide, which is mixed with cryolite and then electrolytically reduced to pure aluminium. Uses Aluminium is used in an enormous variety of products, due to its particular properties. It has low density, is non-toxic, has a high thermal conductivity, has excellent corrosion resistance, and can be easily cast, machined and formed. It is also non-magnetic and nonsparking. It is the second most malleable metal and the sixth most ductile. It is therefore extensively used for kitchen utensils, outside building decoration, and in any area where a strong, light, easily constructed material is needed. The electrical conductivity of aluminium is about 60% that of copper per unit area of cross-section, but it is nevertheless used in electrical transmission lines because of its low density. Alloys of aluminium with copper, manganese, magnesium and silicon are of vital importance in the construction of aeroplanes and rockets. Aluminium, when evaporated in a vacuum, forms a highly reflective coating for both light and heat which does not deteriorate as does a silver coating. These aluminium coatings are used for telescope mirrors, in decorative paper, packages and toys, and have many other uses.

170

The following uses for aluminium are gathered from a number of sources as well as from anecdotal comments. I'd be delighted to receive corrections as well as additional referenced uses (please use the mechanism to add uses). •

cans and foils

•

kitchen utensils

•

outside building decoration

•

industrial applications where a strong, light, easily constructed material is needed

•

although its electrical conductivity is only about 60% that of copper per area of cross section, it is used in electrical transmission lines because of its lightness and price

•

alloys are of vital importance in the construction of modern aircraft and rockets

•

aluminium, evaporated in a vacuum, forms a highly reflective coating for both visible light and radiant heat. These coatings soon form a thin layer of the protective oxide and do not deteriorate as do silver coatings. These coatings are used for telescope mirrors, decorative paper, packages, toys, and in many other uses

•

the oxide, alumina, occurs naturally as ruby, sapphire, corundum, and emery, and is used in glass making and refractories. Synthetic ruby and sapphire are used in the construction of lasers

Biological Role Aluminium has no known biological role. It can be accumulated in the body from daily intake, and at one time was suggested as a potential causative factor in Alzheimer’s disease (senile dementia). General Information The ancient Greeks and Romans used alum (potassium aluminium sulfate) in medicine as an astringent, and in dyeing as a mordant. Sir Humphry Davy proposed the name aluminum for the element, which was undiscovered at the time, and later agreed to change it to aluminium. Aluminium oxide, alumina, occurs naturally as corundum and emery, and is used in glass-making and refractories. The precious stones ruby and sapphire contain aluminium with very small amounts of specific impurities. Reactions 171

Aluminium reacts rapidly with the Oxygen in air to form Aluminium Oxide, Al2O3, which forms a tough layer on the surface of the metal, thereby preventing any further reaction. This aluminium oxide layer can be thickened by a process known as anodising, which involves using the aluminium object as the anode in an electrolytic cell. Reaction of aluminium with air Aluminium is a silvery white metal. The surface of aluminium metal is covered with a thin layer of oxide that helps protect the metal from attack by air. So, normally, aulumium metal does not react with air. If the oxide layer is damaged, the aluminium metal is exposed to attack. Aluminium will burn in oxygen with a brilliant white flame to form the trioxide alumnium(III) oxide, Al2O3. 4Al(s) + 3O2(l) → 2Al2O3(s) Reaction of aluminium with water Aluminium is a silvery white metal. The surface of aluminium metal is covered with a thin layer of oxide that helps protect the metal from attack by air. So, normally, aulumium metal does not react with air. If the oxide layer is damaged, the aluminium metal is exposed to attack, even by water. Reaction of aluminium with the halogens Aluminium metal reacts vigorously with all the halogens to form aluminium halides. So, it reacts with chlorine, Cl2, bromine, I2, and iodine, I2, to form respectively aluminium(III) chloride, AlCl3, aluminium(III) bromide, AlBr3, and aluminium(III) iodide, AlI3. 2Al(s) + 3Cl2(l) → 2AlCl3(s) 2Al(s) + 3Br2(l) → Al2Br6(s) 2Al(s) + 3I2(l) → Al2I6(s) Reaction of aluminium with acids

172

Aluminium metal dissolves readily in dilute sulphuric acid to form solutions containing the aquated Al(III) ion together with hydrogen gas, H2. The corresponding reactions with dilute hydrochloric acid also give the aquated Al(III) ion. Concentrated nitric acid ivates aluminium metal. 2Al(s) + 3H2SO4(aq) → 2Al3+(aq) + 2SO42-(aq) + 3H2(g) 2Al(s) + 6HCl(aq) → 2Al3+(aq) + 6Cl-(aq) + 3H2(g) Reaction of aluminium with bases Aluminium dissolves in sodium hydroxide with the evolution of hydrogen gas, H2, and the formation of aluminates of the type [Al(OH)4]-. 2Al(s) + 2NaOH(aq) + 6H2O → 2Na+(aq) + 2[Al(OH)4]- + 3H2(g)

Isolation Isolation: aluminium would not mormally be made in the laboratory as it is so readily available commercially. Aluminium is mined in huge scales as bauxite (typically Al 2O3.2H2O). Bauxite contains Fe2O3, SiO2, and other impurities. In order to isolate pure aluminium, these impurities must be removed from the bauxite. This is done by the Bayer process. This involves treatment with sodium hydroxide (NaOH) solution, which results in a solution of sodium aluminate and sodium silicate. The iron remains behind as a solid. When CO 2 is blown through the resulting solution, the sodium silicate stays in solution while the aluminium is precipitated out as aluminium hydroxide. The hydroxide can be filtered off, washed, and heated to form pure alumina, Al2O3. The next stage is formation of pure aluminium. This is obtained from the pure Al2O3 by an electrolytic method. Electrolysis is necessary as aluminium is so electropositive. It seems these days that electrolysis of the hot oxide in a carbon lined steel cell acting as the cathode with carbon anodes is most common.

173

GALIUM (Ga)

Physical Information 1. Atomic Number

: 31

2. Relative Atomic Mass (12C=12.000)

: 69.723

3. Melting Point/K

: 303

4. Boiling Point/K

: 2676

5. Density/kg m-3

: 5907 (293K)

6. Ground State Electron Configuration

: [Ar] 3d10 4s2 4p1

7. Electron Affinity (M-M-)/kJ mol-1

: -36

Discovery (L. Gallia: ; also from Latin, gallus, a translation of "Lecoq", a cock) Predicted and described by Mendeleev as ekaaluminum. Gallium was an element whose existence was predicted by Mendeleev in 1871. He predicted that the then unknown element gallium should resemble aluminium in its properties. He suggested therefore the name ekaaluminium (symbol Ea). His predictions for the properties of gallium are remarkably close to the reality. Gallium was discovered spectroscopically by Paul-Emile Lecoq de Boisbaudran in 1875, who in the same year obtained the free metal by electrolysis of a solution of the hydroxide Ga(OH) 3 in KOH. Appearance Gallium is a silvery, glass-like, soft metal. Source Gallium is present in trace amounts in the minerals diaspore, sphalerite, germanite, bauxite and coal. The free metal can be obtained by electrolysis of a solution of gallium(III) hydroxide in potassium hydroxide. 174

Reactions Reaction with air: mild, → Ga2O3 Reaction with 6 M HCl: mild, →H2, GaCl3 Reaction with 6 M NaOH: mild, → H2, [Ga(OH4)]2Uses Gallium readily alloys with most metals, and is used especially in low-melting alloys. It has a high boiling point, which makes it ideal for recording temperatures that would vaporise a thermometer. It has found recent use in doping semiconductors and producing solid-state devices such as transistors. The following uses for gallium are gathered from a number of sources as well as from anecdotal comments. I'd be delighted to receive corrections as well as additional referenced uses (please use the mechanism to add uses). •

gallium wets glass or porcelain, and forms a brilliant mirror when it is painted on glass

•

used for doping semiconductors and producing solid-state devices such as transistors

•

gallium arsenide converts electricity into coherent light

•

alloying

•

90 tons of gallium (2 or 3 years of world production) is used to detect solar neutrinos by the use of the reaction: nu + 71Ga > 71Ge + e-. The rate, although very low (less than 1 interaction per day in 30 tonnes of Ga) makes gallium unique for this purpose. Two experiments are running : - GALLEX using 30 tons in the Gran Sasso underground laboratory (Italy) and SAGE with 60 tons in the Baksan laboratory in Caucasus (Russia). [thanks Michel]

Biological Role Gallium has no known biological role. It is non-toxic. General Information Gallium reacts with acids and alkalis. It has the longest liquid range of all elements, and can be liquid near room temperatures - it can melt in the hand. It also expands as it

175

freezes, which is unusual for a metal, by 3.1%. Gallium wets glass or porcelain, and forms a brilliant mirror when painted on glass. Isolation Isolation: gallium is normally a byproduct of the manufacture of aluminium. The purification of bauxite by the Bayer process results in concentration of gallium in the alkaline solutions from an aluminium:gallum ratio from 5000 to 300. Electrolysis using a mercury electrode gives a further concentration and further electrolysis using a stainless steel cathode of the resulting sodium gallate affords liquid gallium metal. Very pure gallium requires a number of further processes ending with zone refining to make very pure gallium metal.

INDIUM (In)

176

Physical Information Atomic Number

: 49

Relative Atomic Mass (12C=12.000)

: 114.82

Melting Point/K

: 429

Boiling Point/K

: 2353

Density/kg m-3

: 7310 (298K)

Ground State Electron Configuration

: [Kr]4d105s25p1

Electron Affinity (M-M-)/kJ mol

: -34

-1

Discovery (from the brilliant indigo line in its spectrum) Discovered by Reich and Richter, who later isolated the metal. Until 1924, a gram or so constituted the world's supply of this element in isolated form. It is probably about as abundant as silver. About 4 million troy ounces of indium are now produced annually in the Free World. Canada is presently producing more than 1,000,000 troy ounces annually. Appearance Indium is a very soft, silvery-white metal with a brilliant lustre. Source Indium is often associated with zinc minerals and iron, lead and copper ores. It is commercially produced from the zinc minerals, usually as a by-product.

Uses Indium has semiconductor uses in transistors, thermistors and photoconductors. It is also used to make low-temperature alloys; for example, an alloy of 24% indium-76% gallium is liquid at room temperature. Indium can also be plated on to metal and evaporated on to glass to give a mirror with better resistance to corrosion than silver. A tiny long-lived indium battery has been devised to power new electronic watches. The following uses for indium are gathered from a number of sources as well as from anecdotal comments. I'd be delighted to receive corrections as well as additional referenced uses (please use the mechanism to add uses). •

used in making bearing alloys, germanium transistors, rectifiers, thermistors, and photoconductors 177

•

it can be plated onto metal and evaporated onto glass, forming a mirror as good as those made with silver, but with more resistance to atmospheric corrosion

•

photocells

•

used to make low-melting alloys, alloyed with gallium

•

indium is used in solders

Biological Role Indium has no known biological role but has been shown to cause birth defects in unborn children. It has low toxicity. General Information Indium is stable in air and with water, but reacts with acids. Isolation Isolation: indium would not normally be made in the laboratory as it is commercially available. Indium is a byproduct of the formation of lead and zinc. Indium metal is isolated by the electrolysis of indium salts in water. Further processes are required to make very pure indium for electronics purposes.

THALIUM (Tl)

Physical Information Atomic Number

: 81

Relative Atomic Mass (12C=12.000)

: 204.38

Melting Point/K

: 576.7

Boiling Point/K

: 1730

Density/kg m-3

: 11850 (293K)

Ground State Electron Configuration

: [Xe] 4f14 5d10 6s2 6p1 178

Electron Affinity (M-M-)/kJ mol-1

: -30

Discovery (Gr. thallos: a green shoot or twig) Thallium was discovered spectroscopically in 1861 by Crookes. The element was named after the beautiful green spectral line, which identified the element. The metal was isolated both by Crookes and by Lamy in 1862 at about the same time. Appearance Thallium is a soft, silvery metal, but it soon develops a bluish-grey tinge as the oxide forms if it is exposed to the air. Thallium is similar to lead in many of its physical properties. Source Thallium is found in several ores, one of which is pyrites, used in the production of sulphuric acid. The commercial source of thallium is as a by-product of pyrites roasting in sulphuric acid production. It can also be obtained from the smelting of lead and zinc ores. Thallium is also present in manganese nodules found on the ocean floor. Uses Thallium sulfate has been widely employed as a rodenticide and ant killer. It is odorless and tasteless, giving no warning of its presence. Its use, however, has been prohibited in the U.S. since 1975 as a household insecticide and rodenticide. The electrical conductivity of thallium sulfide changes with exposure to infrared light, and this compound is used in photocells. Thallium bromide-iodide crystals have been used as infrared optical materials. Thallium has been used, with sulfur or selenium and arsenic, to produce low melting glasses with become fluid between 125 and 150C. These glasses have properties at room temperatures similar to ordinary glasses and are said to be durable and insoluble in water. Thallium oxide has been used to produce glasses with a high index of refraction, and is used in the manufacture of photo cells. Thallium has been used in treating ringworm and other skin infections; however, its use has been limited because of the narrow margin between toxicity and therapeutic benefits. The following uses for thallium are gathered from a number of sources as well as from anecdotal comments. I'd be delighted to receive corrections as well as additional referenced uses (please use the mechanism to add uses). 179

•

the sulphate was widely used as a rodenticide and ant killer. It is odourless and tasteless, giving no warning of its presence

•

the electrical conductivity of thallium sulphide changes with exposure to infrared light, and so it is used in photocells

•

thallium bromide-iodide crystals are used as infrared detectors

•

used, with sulphur or selenium and arsenic, to produce low melting glasses which become fluid between 125 and 150°C

•

originally used in treating ringworm and other skin infections. Its use was limited because of the narrow margin between toxicity and therapeutic benefits

•

A mercury-thallium alloy, which forms a eutectic at 8.5% thallium, freezes at -60°C, some 20° below the freezing point of mercury

Biological Role Thallium has no known biological role. It is very toxic and teratogenic. of the metal with the skin is dangerous, and there is evidence that the vapour is both teratogenic and carcinogenic. General Information Thallium is soft, malleable and can be cut with a knife. It tarnishes readily in moist air and reacts with steam to form the hydroxide. It is attacked by all acids, most rapidly nitric acid. Isolation Isolation: thallium metal would not normally be made in the laboratory as it is available commercially. Crude thallium is present as a component in flue dust along with arsenic, cium, indium, germanium, lead, nickel, selenium, tellurium, and zinc. This is done by dissolving in dilute acid, precipitating out lead sulphate, and then adding HCl to precipitate thallium chloride, TlCl. Further purification can be achieve by electrolysis of soluble thallium salts. Reactions Reaction of thallium with air

180

Freshy cut thallium tarnishes slowly to give a grey oxide film that protects the remaining metal from further oxidation. When heated strongly to red heat in air, poisonous thallium(I) oxide is formed. 2Tl(s) + O2(g) → Tl2O(s) Reaction of thallium with water Thallium seems not to react with air-free water. Thallium metal tarnishes slowly in moist air or dissolves in water to give poisonous thallium(I) hydroxide. 2Tl(s) + 2H2O(l) → 2TlOH(aq) + H2(g) Reaction of thallium with the halogens Thallium metal reacts vigorously with fluorine, F2, chlorine, Cl2, and bromine, Br2, to form the dihalides thallium(III) fluoride, TlF3, thallium(III) chloride, TlCl3, tand hallium(III) bromide, TlBr3, respectively. All these compounds are poisonous. 2Tl(s) + 3F2(g) → 2TlF3(s) [] 2Tl(s) + 3Cl2(g) → 2TlCl3(s) [] 2Tl(s) + 3Br2(l) → 2TlBr3(s) [] Reaction of thallium with acids Thallium dissolves only slowly in sulphuric acid, H2SO4, or hydrochloric acid, HCl, because the poisonous thallium(I) salts produced are not very soluble.

181

ELEMENTS OF GROUP III B SCANDIUM (Sc)

Physic, chemical, and atomic properties 1. Name

: Scandium

2. Symbol

: Sc

3. Atomic number

: 21

4. Atomic weight

: 44.9559

5. Atomic radius

: 160.6 pm

6. Standard state

: Solid at 298 K

7. Melting point

: 15410C

8. Boiling Point

: 28360C

9. Liquid range

: 1289 K

10. Group in periodic table

:3

11. Period in periodic table

:4

12. Block in periodic table

: d-block

13.Electron configuration

: [Ar]4s23d1 182

14. Oxidation States

:3

15. Colour

: silvery white

16. classification

: Metallic

Electronegativities The most used definition of electronegativity is that an element's electronegativity is the power of an atom when in a molecule to attract electron density to itself. The electronegativity depends upon a number of factors and in particuler as the other atoms in the molecule. The first scale of electronegativity was developed by Linus Pauling and on his scale scandium has a value of 1.36 on a scale running from from about 0.7 (an estimate for francium) to 2.20 (for hydrogen) to 3.98 (fluorine). Electronegativity has no units but "Pauling units" are often used when indicating values mapped on to the Pauling scale. On the interactive plot below you may find the "Ball chart" and "Shaded table" styles most useful. There are a number of ways to produce a set of numbers representing electronegativity and five are given in the table above. The Pauling scale is perhaps the most famous and suffices for many purposes. Electronic configuration The following represents the electronic configuration and its associated term symbol for the ground state neutral gaseous atom. The configuration associated with scandium in its compounds is not necessarily the same. •

Ground state electron configuration: [Ar].3d1.4s2

•

Shell structure: 2.8.9.2

•

Term symbol:

D3/2

2

183

One measure of size is the element-element distance within the element. The bond length in ScSc is: 321.2 pm. It is not always easy to make sensible comparisons between the elements however as some bonds are quite short because of multiple bonding (for instance the O=O distance in O2 is short because of the the double bond connecting the two atoms. There are several other ways ways to define radius for atoms and ions. Follow the appropriate hyperlinks for literature references and definitions of each type of radius. All values of radii are given in picometres (pm). Conversion factors are: •

1 pm = 1 x 10-12 metre (meter)

•

100 pm = 1 Ångstrom

•

1000 pm = 1 nanometre (nm, nanometer)

Discovery (L. Scandia: Scandinavia) On the basis of the Periodic System, Mendeleev predicted the existence of ekaboron, which would have an atomic weight between 40 of calcium and 48 of titanium. Scandium was discovered by Lars Fredrik Nilson at Sweden. Origin of name from the Latin word “Scandia” meaning “Scandinavia”. The element was discovered by Nilson in 1878 in the minerals euxenite and gadolinite, which had not yet been found anywhere except in Scandinavia. He and his coworkers were actually looking for rare earth metals. By processing 10 kg of euxenite and other residues of rare-earth minerals, Nilson was able to prepare about 2g of highly pure scandium oxide (scandia, Sc2O3) of high purity. Later scientists pointed out that Nilson's scandium was identical with Mendeleev's ekaboron. In 1871 Mendeleev predicted that an element should exist that would resemble boron in its properties. He therefore called it ekaboron, (symbol Eb). Per Theodor Cleave found scandium oxide at about the same time. He noted that the new element was the element ekaboron predicted by Mendeleev in 1871. Appearance Scandium is a silver-white metal which develops a slightly yellowish or pinkish cast upon exposure to air. A relatively soft element, scandium resembles yttrium and the rareearth metals more than it resembles aluminum or titanium. It is a very light metal and has a 184

much higher melting point than aluminum, making it of interest to designers of spacecraft. Scandium is not attacked by a 1:1 mixture of HNO3 and 48% HF. Scandium reacts repidly with many acids. Scandium is apparently a much more abundant element in the sun and certain stars than on eart. Source Scandium is apparently much more abundant (the 23rd most) in the sun and certain stars than on earth (the 50th most abundant). It is widely distributed on earth, occurring in very minute quantities in over 800 mineral species. The blue color of beryl (aquamarine variety) is said to be due to scandium. It occurs as a principal component in the rare mineral thortveitite, found in Scandinavia and Malagasy. It is also found in the residues remaining after the extraction of tungsten from Zinnwald wolframite, and in wiikite and bazzite. Most scandium is presently being recovered from thortveitite or is extracted as a byproduct from uranium mill tailings. Metallic scandium was first prepared in 1937 by Fischer, Brunger, and Grienelaus who electrolyzed a eutectic melt of potassium, lithium, and scandium chlorides at 700 to 800oC. Tungsten wire and a pool of molten zinc served as the electrodes in a graphite crucible. Pure scandium is now produced by reducing scandium fluoride with calcium metal. The production of the first pound of 99% pure scandium metal was announced in 1960. Uses The following uses for scandium are gathered from a number of sources as well as from anecdotal comments. I'd be delighted to receive corrections as well as additional referenced uses (please use the mechanism to add uses). •

isotope tracing in crude oil analysis

•

the iodide added to mercury vapour lamps and produces a highly efficient light source resembling sunlight, which is important for indoor or night-time colour TV transmission.

•

apparently scandium is a component of alloys used to make metallic baseball bats (mercifully metallic cricket bats were banned after Dennis Lillee tried to use one).

185

About 20 kg of scandium (as Sc2O3) are used yearly in the U.S. to produce highintensity lights. The radioactive isotope 46Sc is used as a tracing agent in refinery crackers for crude oil, etc. Scandium iodide added to mercury vapor lamps produces a highly efficient light source resembling sunlight, which is important for indoor or night-time color TV. Isolation Isolation: preparation of metallic samples of scandium is not normally necessary given that it is commercially avaialable. In practice littel scandium is produced. The mineral thortveitite contains 35-40% Sc2O3 is used to produce scandium metal but another important source is as a byproduct from uranium ore processing, even though these only contain 0.02% Sc2O3 . Handling Little is yet known about the toxicity of scandium, therefore it should be handled with care. Reactions Reaction of scandium with air Scandium metal tarnishes in air and burns readily to form scandium (III) oxide, Sc2O3. 4Sc + 3O2 → 2Sc2O3 Reaction of scandium with water When finely divided, or heated, scandium metal dissolves in water to form solutions containing the aquated Sc(III) ion together with hydrogen gas, H2. 2Sc(s) + 6H2O(aq) → 2Sc3+(aq) + 6OH-(aq) + 3H2(g) Reaction of scandium with the halogens

186

Scandium is very reactive towards the halogens fluorine, F2, chlorine, Cl2 bromine, Br2, and iodine, I2, and burns to form the trihalides scandium(III) fluoride, ScF3 , scandium(III) chloride, ScCl3, scandium(III) bromide, ScBr3, and scandium(III) iodide, ScI3 respectively. 2Sc(s) + 3F2(g) → 2ScF3(s) 2Sc(s) + 3Cl2(g) → 2ScCl3(s) 2Sc(s) + 3Br2(g) → 2ScBr3(s) 2Sc(s) + 3I2(g) → 2ScI3(s) Reaction of scandium with acids Scandium metal dissolves readily in dilute hydrochloric acid to form solutions containing the aquated Sc(III) ion together with hydrogen gas, H2. 2Sc(s) + 6HCl(aq) → 2Sc3+(aq) + 6Cl-(aq) + 3H2(g)

YTTRIUM (Y)

Physic, chemical, and atomic properties 1. Name

: Yttrium

2. Symbol

:Y

3. Atomic number

: 39

4. Atomic weight

: 88.9059

5. Atomic radius

: 181 pm

6. Standard state

: solid at 298 K 187

7. Melting point

: 15220C

8. Boiling Point

: 33450C

9. Group in periodic table

:3

10. Period in periodic table

:5

11. Block in periodic table

: d-block

12.Electron configuration

: [Kr]5s14d1

13. Oxidation States

:3

14.Relative Atomic Mass

: 88.91

15. Liquid range

: 1810 K

Electronegativities The most used definition of electronegativity is that an element's electronegativity is the power of an atom when in a molecule to attract electron density to itself. The electronegativity depends upon a number of factors and in particuler as the other atoms in the molecule. The first scale of electronegativity was developed by Linus Pauling and on his scale yttrium has a value of 1.22 on a scale running from from about 0.7 (an estimate for francium) to 2.20 (for hydrogen) to 3.98 (fluorine). Electronegativity has no units but "Pauling units" are often used when indicating values mapped on to the Pauling scale. On the interactive plot below you may find the "Ball chart" and "Shaded table" styles most useful. Electronic configuration The following represents the electronic configuration and its associated term symbol for the ground state neutral gaseous atom. The configuration associated with yttrium in its compounds is not necessarily the same. •

Ground state electron configuration: [Kr].4d1.5s2

•

Shell structure: 2.8.18.9.2

•

Term symbol:

D3/2

2

188

A schematic representation of the shell structure of yttrium - not what the atom of yttrium "looks like". One measure of size is the element-element distance within the element. The bond length in YY is: 355.1 pm. It is not always easy to make sensible comparisons between the elements however as some bonds are quite short because of multiple bonding (for instance the O=O distance in O2 is short because of the the double bond connecting the two atoms. There are several other ways ways to define radius for atoms and ions. Follow the appropriate hyperlinks for literature references and definitions of each type of radius. All values of radii are given in picometres (pm). Conversion factors are: •

1 pm = 1 x 10-12 metre (meter)

•

100 pm = 1 Ångstrom

•

1000 pm = 1 nanometre (nm, nanometer)

Discovery Yettrium was discovered by Johann Gadolin at 1794 in Finland. Origin of name : after the village of “Ytterby” near Vaxholm in Sweden. Yttria (yttrium oxide, Y2O3) was discovered by Gadolin in 1794 in a mineral called gadolinite from Yetterby. Ytterby is the site of a quarry which contains many unusual minerals containing erbium, terbium, and ytterbium as well as yttrium. Fredrich Wohler obtained the impure element in 1828 by reduction of the anhydrous chlorida (YCl3) with patassium. In 1843 Mosander showed that yttira could be resolved into the oxides (or earths) of three elements. The name yttria was reserved for the most basic one; the others were named erbia and terbia Appearance Yttrium has a silver-metallic luster and is relatively stable in air. Turnings of the metal, however, ignite in air if their temperature exceeds 400 oC. Finely divided yttrium is very unstable in air. Yttrium is found in most rare earth minerals. Moon rocks contain yttrium is uses as a “phosphor” to produce the red colour in television screens. Source 189

Yttrium occurs in nearly all of the rare-earth minerals. Analysis of lunar rock samples obtained during the Apollo missions show a relatively high yttrium content. It is recovered commercially from monazite sand, which contains about 3%, and from bastnasite, which contains about 0.2%. Wohler obtained the impure element in 1828 by reduction of the anhydrous chloride with potassium. The metal is now produced commercially by reduction of the fluoride with calcium metal. It can also be prepared by other techniques. Uses Yttrium oxide is one of the most important compounds of yttrium and s for the largest use. It is widely used in making YVO4 europium, and Y2O3 europium phosphors to give the red color in color television tubes. Hundreds of thousands of pounds are now used in this application. Yttrium oxide also is used to produce yttrium-iron-garnets, which are very effective microwave filters. Yttrium iron, aluminum, and gadolinium garnets, with formulas such as Y 3Fe5O12 and Y3Al5O12, have interesting magnetic properties. Yttrium iron garnet is also exceptionally efficient as both a transmitter and transducer of acoustic energy. Yttrium aluminum garnet, with a hardness of 8.5, is also finding use as a gemstone (simulated diamond). Small amounts of yttrium (0.1 to 0.2%) can be used to reduce the grain size in chromium, molybdenum, zirconium, and titanium, and to increase strength of aluminum and magnesium alloys. Alloys with other useful properties can be obtained by using yttrium as an additive. The metal can be used as a deoxidizer for vanadium and other nonferrous metals. The metal has a low cross section for nuclear capture. exists in equilibrium with its parent

Y, one of the isotopes of yttrium,

90

Sr, a product of nuclear explosions. Yttrium has been

90

considered for use as a nodulizer for producing nodular cast iron, in which the graphite forms compact nodules instead of the usual flakes. Such iron has increased ductility. Yttrium also can be used in laser systems and as a catalyst for ethylene polymerization reactions. It also has potential use in ceramic and glass formulas, as the oxide has a high melting point and imparts shock resistance and low expansion characteristics to glass. The following uses for yttrium are gathered from a number of sources as well as from anecdotal comments. I'd be delighted to receive corrections as well as additional referenced uses (please use the mechanism to add uses).

190

•

YVO4 europium, and Y203 europium phosphors give the red colour in colour television tubes

•

the oxide is used to produce yttrium-iron-garnets, which are very effective microwave filters

•

yttrium iron, aluminum, and gadolinium garnets have interesting magnetic properties. Yttrium iron garnet is also exceptionally efficient as both a transmitter and transducer of acoustic energy

•

yttrium aluminium garnet is a gemstone (simulated diamond)

•

used in laser systems

•

used as a catalyst for ethene polymerization

•

potential use in ceramic and glasses as the oxide has a high melting point and imparts shock resistance and low expansion characteristics to glass

•

increases increase the strengths of alloys of metals such as chromium, aluminium, and magnesium.

Isolation Isolation: yttrium metal is available commercially so it is not normally necesary to make it in the laboratory. Yttrium is found in lathanoid minerals and the extraction of the yttrium and the lanthanoid metals from the ores is highly complex. Initially, the metals are extractedas salts from the ores by extraction with sulphuric acid (H 2SO4), hydrochloric acid (HCl), and sodium hydroxide (NaOH). Modern purification techniques for these lanthanoid salt mixtures involve selective complexation techniques, solvent extractions, and ion exchange chromatography. Pure yttrium is available through the reduction of YF3 with calcium metal. 2YF3 + 3Ca → 2Y + 3CaF2 Reactions Reaction of yttrium with air Yttrium metal tarnishes slowly in air and burns readily to form yttrium (III) oxide, Y2O3. 4Y + 3O2 → 2Y2O3 Reaction of yttrium with water 191

When finely divided, or heated, yttrium metal dissolves in water to form solutions containing the aquated Y(III) ion together with hydrogen gas, H2. 2Y(s) + 6H2O(aq) → 2Y3+(aq) + 6OH-(aq) + 3H2(g) Reaction of yttrium with the halogens Yttrium is very reactive towards the halogens fluorine, F2, chlorine, Cl2 bromine, Br2, and iodine, I2, and burns to form the trihalides yttrium(III) fluoride, YF3 , yttrium(III) chloride, YCl3, yttrium(III) bromide, YBr3, and yttrium(III) iodide, YI3 respectively. 2Y(s) + 3F2(g) → 2YF3(s) 2Y(s) + 3Cl2(g) → 2YCl3(s) 2Y(s) + 3Br2(g) → 2YBr3(s) 2Y(s) + 3I2(g) → 2YI3(s) Reaction of yttrium with acids Yttrium metal dissolves readily in dilute hydrochloric acid to form solutions containing the aquated Y(III) ion together with hydrogen gas, H2. 2Y(s) + 6HCl(aq) → 2Y3+(aq) + 6Cl-(aq) + 3H2(g)

LANTHANUM (La)

192

Physic, chemical, and atomic properties 1. Name

: Lanthanum

2. Symbol

: La

3. Atomic number

: 57

4. Atomic weight

: 138.9055

5. Atomic radius

: 187.7 pm

6. Melting point

: 9180C

7. Boiling Point

: 34640C

8. Standard state

: solid at 298 K

9. Liquid range

: 2550 K

10. Period in periodic table

:6

11. Block in periodic table

: f-block

12.Electron configuration

: [Xe]6s25d1

13. Oxidation States

:2

14.Relative Atomic Mass

: 139

15. Relative Density

: 6.146

16. Colour

: silvery white

17. Classification

: Metallic

Electronegativitie The most used definition of electronegativity is that an element's electronegativity is the power of an atom when in a molecule to attract electron density to itself. The electronegativity depends upon a number of factors and in particuler as the other atoms in the molecule. The first scale of electronegativity was developed by Linus Pauling and on his scale lanthanum has a value of 1.10 on a scale running from from about 0.7 (an estimate for francium) to 2.20 (for hydrogen) to 3.98 (fluorine). Electronegativity has no units but "Pauling units" are often used when indicating values mapped on to the Pauling scale. On the interactive plot below you may find the "Ball chart" and "Shaded table" styles most useful. 193

Electronic configuration The following represents the electronic configuration and its associated term symbol for the ground state neutral gaseous atom. The configuration associated with lanthanum in its compounds is not necessarily the same. •

Ground state electron configuration: [Xe].5d1.6s2

•

Shell structure: 2.8.18.18.9.2

•

Term symbol:

D3/2

2

A schematic representation of the shell structure of lanthanum - not what the atom of lanthanum "looks like". One measure of size is the element-element distance within the element. The bond length in LaLa is: 373.9 pm. It is not always easy to make sensible comparisons between the elements however as some bonds are quite short because of multiple bonding (for instance the O=O distance in O2 is short because of the the double bond connecting the two atoms. There are several other ways ways to define radius for atoms and ions. Follow the appropriate hyperlinks for literature references and definitions of each type of radius. All values of radii are given in picometres (pm). Conversion factors are: •

1 pm = 1 x 10-12 metre (meter)

•

100 pm = 1 Ångstrom

•

1000 pm = 1 nanometre (nm, nanometer)

Discovery Lanthanum was discovered by Carl Gustaf Mosander at 1839 in Sweden. Origin of name: from the Greek word "lanthanein" meaning "to lie hidden". Carl Gustav Mosander recognized the element lanthanum in impure cerium nitrate in 1839. His extraction resulted 194

in the oxide lanthana (La2O3). A number of other lanthanides (rare-earths) were later discovered by identification of the impurities in yttrium and cerium compounds. (Greek lanthanein: to lie hidden) Mosander in 1839 extracted lanthana from impure cerium nitrate and recognized the new element. Lanthanum was isolated in relatively pure form in 1923. Iron exchange and solvent extraction techniques have led to much easier isolation of the socalled "rare-earth" elements.

Appearance Lanthanum is silvery white, malleable, ductile, and soft enough to be cut with a knife. It is one of the most reactive of the rare-earth metals. It oxidizes rapidly when exposed to air. Cold water attacks lanthanum slowly, while hot water attacks it much more rapidly. The metal reacts directly with elemental carbon, nitrogen, boron, selenium, silicon, phosphorus, sulfur, and with halogens. At 310 °C, lanthanum changes from a hexagonal to a face-centered cubic structure, and at 865 °C it again transforms into a body-centered cubic structure. Pure metal lutetium has been isolated only in recent years and is one of the more difficult to prepare. It can be prepared by the reduction of anhydrous LuCl 3 or LuF3 by an alkali or alkaline earth metal. The metal is silvery white and relatively stable in air. It is a rare earth metal and perhaps the most expensive of all rare elements. It is found in small amounts with all rare earth metals, and is very difficult to separate from other rare elements. Source Lanthanum is found in rare-earth minerals such as cerite, monazite, allanite, and bastnasite. Monazite and bastnasite are principal ores in which lanthanum occurs in percentages up to 25 percent and 38 percent respectively. Misch metal, used in making lighter flints, contains about 25 percent lanthanum. The availability of lanthanum and other rare earths has improved greatly in recent years. The metal can be produced by reducing the anhydrous fluoride with calcium . Uses Rare-earth compounds containing lanthanum are extensively used in carbon lighting applications, especially by the motion picture industry for studio lighting and projection. This application consumes about 25 percent of the rare-earth compounds produced. La 2O3

195

improves the alkali resistance of glass, and is used in making special optical glasses. Small amounts of lanthanum, as an additive, can be used to produce nodular cast iron. There is current interest in hydrogen sponge alloys containing lanthanum. These alloys take up to 400 times their own volume of hydrogen gas, and the process is reversible. Every time they take up the gas, heat energy is released; therefore these alloys have possibilities in an energy conservation system. The following uses for lanthanum are gathered from a number of sources as well as from anecdotal comments. I'd be delighted to receive corrections as well as additional referenced uses (please use the mechanism to add uses). •

rare-earth compounds containing lanthanum are extensively used in carbon lighting applications, especially by the motion picture industry for studio lighting and projection

•

La improves the alkali resistance of glass, and is used in making special optical

203

glasses •

small amounts as an additive are used to produce nodular cast iron

•

hydrogen sponge alloys containing lanthanum reversibly take up to 400 times their own volume of hydrogen gas. Heat is released, therefore these alloys have potential in energy conservation systems

•

lighter flints

•

alloys

Handing Lanthanum and its compounds have a low to moderate acute toxicity rating; therefore, care should be taken in handling them. Isolation Isolation: lanthanum metal is available commercially so it is not normally necessary to make it in the laboratory, which is just as well as it is difficult to separate it from as the pure metal. This is largely because of the way it is found in nature. The lanthanoids are found in nature in a number of minerals. The most important are xenotime, monazite, and bastnaesite. The first two are orthophosphate minerals LnPO4 (Ln deonotes a mixture of all the lanthanoids except promethium which is vanishingly rare) and the third is a fluoride carbonate LnCO3F. Lanthanoids with even atomic numbers are more common. The most 196

comon lanthanoids in these minerals are, in order, cerium, lanthanum, neodymium, and praseodymium. Monazite also contains thorium and ytrrium which makes handling difficult since thorium and its decomposition products are radioactive. For many purposes it is not particularly necessary to separate the metals, but if separation into individual metals is required, the process is complex. Initially, the metals are extracted as salts from the ores by extraction with sulphuric acid (H2SO4), hydrochloric acid (HCl), and sodium hydroxide (NaOH). Modern purification techniques for these lanthanoid salt mixtures are ingenious and involve selective complexation techniques, solvent extractions, and ion exchange chromatography. Pure lanthanum is available through the reduction of LaF3 with calcium metal. 2LaF3 + 3Ca → 2La + 3CaF2 This would work for the other calcium halides as well but the product CaF 2 is easier to handle under the reaction conditions (heat to 50°C above the melting point of the element in an argon atmosphere). Excess calcium is removed from the reaction mixture under vacuum. Reactions Reaction of lanthanum with air Lanthanum metal tarnishes slowly in air and burns readily to form lanthanum (III) oxide, La2O3. 4La + 3O2 → 2La2O3 Reaction of lanthanum with water The silvery white metal lanthanum is quite electropositive and reacts slowly with cold water and quite quickly with hot water to form lanthanum hydroxide, La(OH)3, and hydrogen gas (H2). 2La(s) + 6H2O(g) → 2La(OH)3(aq) + 3H2(g) Reaction of lanthanum with the halogens Lanthanum metal reacts with all the halogens to form lanthanum(III) halides. So, it reacts with fluorine, F2, chlorine, Cl2, bromine, I2, and iodine, I2, to form respectively lanthanum(III) 197

bromide,

LaF3,

lanthanum(III)

chloride,

LaCl3,

lanthanum(III)

bromide,

LaBr3,

and

lanthanum(III) iodide, LaI3. 2La(s) + 3F2(g) → 2LaF3(s) 2La(s) + 3Cl2(g) → 2LaCl3(s) 2La(s) + 3Br2(g) → 2LaBr3(s) 2La(s) + 3I2(g) → 2LaI3(s) Reaction of lanthanum with acids Lanthanum metal dissolves readily in dilute sulphuric acid to form solutions containing the aquated La(III) ion together with hydrogen gas, H 2. It is quite likely that La3+(aq) exists as largely the complex ion [La(OH2)9]3+ 2La(s) + 3H2SO4(aq) → 2La3+(aq) + 3SO42-(aq) + 3H2(g)

198

ACTINIUM (Ac)

Physical Information 1. Name

: Actinium

2. Symbol

: Ac

3. Atomic number

: 89

4. Atomic weight

: 227

5. Atomic radius

: 187,8 pm

6. Melting point

: 1051

7. Boiling Point

: 3159

8. Liquid range

: 2250 K

9. Standard state

: solid at 298 K

10. Group name

: Actinoid

11. Period in periodic table

: 7 (actinoid)

12. Block in periodic table

: f-block

13.Electron configuration

: [Rn]7s26d1

14. Oxidation States

:3

15.Relative Atomic Mass

: 227 (approximately)

16.Relative Density

: 10

17. Colour

: silvery 199

18. Classification

: Metallic

Electronegativities The most used definition of electronegativity is that an element's electronegativity is the power of an atom when in a molecule to attract electron density to itself. The electronegativity depends upon a number of factors and in particuler as the other atoms in the molecule. The first scale of electronegativity was developed by Linus Pauling and on his scale actinium has a value of 1.1 on a scale running from from about 0.7 (an estimate for francium) to 2.20 (for hydrogen) to 3.98 (fluorine). Electronegativity has no units but "Pauling units" are often used when indicating values mapped on to the Pauling scale. On the interactive plot below you may find the "Ball chart" and "Shaded table" styles most useful. Electronic Configuration The following represents the electronic configuration and its associated term symbol for the ground state neutral gaseous atom. The configuration associated with actinium in its compounds is not necessarily the same. •

Ground state electron configuration: [Rn].6d1.7s2

•

Shell structure: 2.8.18.32.18.9.2

•

Term symbol:

2

D3/2

One measure of size is the element-element distance within the element. The bond length in AcAc is: 375.6 pm. It is not always easy to make sensible comparisons between the elements however as some bonds are quite short because of multiple bonding (for instance the O=O distance in O2 is short because of the the double bond connecting the two atoms. There are several other ways ways to define radius for atoms and ions. Follow the appropriate 200

hyperlinks for literature references and definitions of each type of radius. All values of radii are given in picometres (pm). Conversion factors are: •

1 pm = 1 x 10-12 metre (meter)

•

100 pm = 1 Ångstrom

•

1000 pm = 1 nanometre (nm, nanometer)

Discovery Actinium, named from the Greek aktinos (ray) is a rare, extremely radioactive metal that glows in the dark (the photo shown above is of Ac 2O3). Discovered by Andre Debierne in 1899 in and independently by F. Giesel in 1902, both of whom obtained it while working on separation techniques for rare eart oxides. Occurs naturally in association with uranium minerals. Appearance Actinium-227, a decay product of uranium-235, is a beta emitter with a 21.6-year halflife. Its principal decay products are thorium-227 (18.5-day half-life), radium-223 (11.4-day half-life), and a number of short-lived products including radon, bismuth, polonium, and lead isotopes. In equilibrium with its decay products, it is a powerful source of alpha rays. Actinium metal has been prepared by the reduction of actinium fluoride with lithium vapor at about 1100 to 1300-degrees C. The chemical behavior of actinium is similar to that of the rare earths, particularly lanthanum. Purified actinium comes into equilibrium with its decay products at the end of 185 days, and then decays according to its 21.6-year half-life. It is about 150 times as active as radium, making it of value in the production of neutrons. Source Actinium occurs naturally in uranium minerals. It is made by the neutron bombardment of the radium isotope

Ra.

228

Uses Actinium is used as a source of Neutrons. The following uses for actinium are gathered from a number of sources as well as from anecdotal comments. I'd be delighted to receive corrections as well as additional referenced uses (please use the mechanism to add uses). 201

•

thermoelectric power

•

source of neutrons

General Information Actinium occurs in association with Uranium ores, as a decay product of Uranium-235. It has three Isotopes nuclide

:

atomic mass

:

natural abundance

:

half-life

Ac

225

Ac

227

Ac

228

227.03 0% :

10 days

trace 21.6 yrs

trace 6.13 h

Actinium 227, which has a half-life of 22 years, is formed by the irradiation of Radium.

202

Reference http://www.ucc.ie/academic/chem/dolchem/html/elem/group.html kamis 2010. pkl.18:16 http://www.periodic.lanl.gov/elements/89.html kamis 23 September

23 September

10.pkl.18:16

http://www.webelements.com/index.html kamis 23 September 2010 pkl 18.17

http://www.britannica.com/EBchecked/topic/74395/boron-group-element Contributor: Alan Gibbs Massey Rabu 22 September 2010 pkl. 13.16

Primary

http://www.lenntech.com/periodic/elements/al.htm#ixzz10KRFTEVJ Rabu 22 september 2010 pukul 14.13

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CHAPTER I INTRODUCTION

In periodicity, there are many elements that classification based on its properties. For example are metal and non metal. The elements are classification into groups and period. Periodicity has 144 elements with 16 groups (8 groups A and 8 groups B), Aktanida and Lantanida. This paper will discuss about elements in groups VA and VB. Groups VA consist of Nitrogen, Phospor, Arsenic, Antimony , and Bismut. Groups VB consist of Vanadium, Niobium, Tantalum, and Dubnium. The elements in group VA and VB has application in daily activity. Nitrogen has use for freeze. Phospor uses for toxic of rat and makes alloy. Bismut uses for makes tender iron, and catalyst for make akrilat fiber. Antanium(Sb) uses as semiconductor and batere. Arsen as additive of Ge and Si. Vanadium used in nuclear reactors. Niobium used in surgical implants because they do not react with human tissue. Tantalum resists corrosion and is almost impervious to chemical attack.

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CHAPTER II MATTER

2.1. Elements of group VA VA group consists of five elements, namely Nitrogen, Phospor, Bismut, Antimony, and Arsen. Some properties of the elements of this group we can see in the table bellow : Properties Atomic number

Atomic mass Electronegativity according to Pauling

Nitrogen

Phospor

Arsen

Antimony

Bismut

7

15

33

51

83

30.973762 74.9216 g.mol -1 g.mol -1

14.0067 g.mol -1

2.19

3.0 1.25*10 g.cm-3 at 20°C

1.823, 2.2–2.34, 2.36, 2.69 g·cm−3

-210 °C

44.2 °C, 610 °C

-3

Density

Melting point

2.0

5.7 g.cm-3 at 14°C

814 °C (36 atm)

121.75 g.mol -1

208.98040 g.mol -1

1.9

2.02

6.684 g.cm-3

9.78 g·cm−3

631 °C

544.7 K271.5 ° ,C520.7 ° , F

Boiling point

-195.8 °C

280.5 °C

615 °C (sublimation)

1380 °C

1837 K, 1564 °C, 2847 °F

Vanderwaals radius

0.092 nm

180 pm

0.139 nm

0.159 nm

207 pm

Ionic radius

0.171 nm (-3) ; 0.011 (+5) ; 0.016 (+3)

0.245 nm (3); 0.062 nm (+5); 0.076 nm (+3)

-

3s2 3p3

[ Ar ] 3d10 4s2 4p3

[ Kr ] 4d10 5s25p3

[Xe]4f145d106s2 6p3

1011.8

947 kJ.mol -1

834 kJ.mol -1

703 kJ·mol−1

-

[Ne] Electronic shell

[He]2s 2p

Energy of first ionisation

1402 kJ.mol -1

2

3

0.222 nm (2) 0,047 nm (+5) 0,058 (+3)

205

kJ·mol−1 Energy of second ionisation

2856 kJ.mol -1

1907 kJ·mol−1

1798 kJ.mol -1

1595 kJ.mol -1

1610 kJ·mol−1

Energy of third ionisation

4578 kJ.mol -1

2914.1 kJ·mol−1

2736 kJ.mol -1

2443 kJ.mol -1

2466 kJ·mol−1

2.1.1. Nitrogen Nitrogen (pronounced) is a chemical element that has the symbol N, atomic number of 7 and atomic mass 14.00674 u. Elemental nitrogen is a colorless, odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78.08% by volume of Earth's atmosphere. Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in breaking the bond to convert the N2 into useful compounds, but releasing large amounts of often useful energy, when these compounds burn, explode, or decay back into nitrogen gas. The element nitrogen was discovered by Scottish physician Daniel Rutherford in 1772. Nitrogen occurs in all living organisms. It is a constituent element of amino acids and thus of proteins, and of nucleic acids (DNA and RNA). It resides in the chemical structure of almost all neurotransmitters, and is a defining component of alkaloids, biological molecules produced by many organisms.

1. History Nitrogen is formally considered to have been discovered by Daniel Rutherford in 1772, who called it noxious air or fixed air. That there was a fraction of air that did not combustion was well-known to the late 18th century chemist. Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as "mephitic air" or azote, from the Greek word (azotos) meaning "lifeless".Animals died in it, and it was the principal component of air in which animals had suffocated and flames had burned to extinction. Lavoisier's name for nitrogen is used in many languages (French, Russian, etc.) and still remains in English in the common names of many compounds, such as

206

hydrazine and compounds of the azide ion. Compounds of nitrogen were known in the Middle Ages. The alchemists knew nitric acid as aqua fortis (strong water). The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold (the king of metals). The earliest military, industrial and agricultural applications of nitrogen compounds involved uses of saltpeter (sodium nitrate or potassium nitrate), notably in gunpowder, and later as fertilizer. In 1910, Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", an allotrope considered to be monatomic. The "whirling cloud of brilliant yellow light" produced by his apparatus reacted with quicksilver to produce explosive mercury nitride.

2. Properties Nitrogen is a nonmetal, with an electronegativity of 3.04. It has five electrons in its outer shell and is therefore trivalent in most compounds. The triple bond in molecular nitrogen (N2) is the strongest. The resulting difficulty of converting N2 into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N2, have dominated the role of nitrogen in both nature and human economic activities. At atmospheric pressure molecular nitrogen condenses (liquefies) at 77 K (−195.8 °C) and freezes at 63 K (−210.0 °C) into the beta hexagonal closepacked crystal allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the alpha cubic crystal allotropic form. Liquid nitrogen, a fluid resembling water in appearance, but with 80.8% of the density (the density of liquid nitrogen at its boiling point is 0.808 g/mL), is a common cryogen. Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like N3 and N4. Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced using a diamond anvil cell, nitrogen polymerizes into the single-bonded cubic gauche crystal structure. This structure is similar to that of diamond, and both have extremely strong covalent bonds. N4 is nicknamed "nitrogen diamond. a. Isotopes There are two stable isotopes of nitrogen: 14N and 15N. By far the most common is 14N (99.634%), which is produced in the CNO cycle in stars. Of the ten isotopes produced synthetically, 13N has a half-life of ten minutes and the remaining isotopes have half-lives on the order of seconds or less. Biologically mediated reactions (e.g., assimilation, nitrification, and denitrification)

207

strongly control nitrogen dynamics in the soil. These reactions typically result in 15N enrichment of the substrate and depletion of the product. A small part (0.73%) of the molecular nitrogen in Earth's atmosphere is the isotopologue 14N15N, and almost all the rest is 14N2. Radioisotope 16N is the dominant radionuclide in the coolant of pressurized water reactors during normal operation. It is produced from 16O (in water) via (n,p) reaction. It has a short half-life of about 7.1 s, but during its decay back to 16O produces highenergy gamma radiation (5 to 7 MeV). Because of this, the access to the primary coolant piping must be restricted during reactor power operation. 16N is one of the main means used to immediately detect even small leaks from the primary coolant to the secondary steam cycle. Electromagnetic spectrum:

A 1×5 cm vial of glowing ultrapure nitrogen tube.

Nitrogen discharge (spectrum)

Molecular nitrogen (14N2) is largely transparent to infrared and visible radiation because it is a homonuclear molecule and thus has no dipole moment to couple to electromagnetic radiation at these wavelengths. Significant absorption occurs at extreme ultraviolet wavelengths, beginning around 100 nanometers. This is associated with electronic transitions in the molecule to states in which charge is not distributed evenly between nitrogen atoms. Nitrogen absorption leads to significant absorption of ultraviolet radiation in the Earth's upper atmosphere and the atmospheres of other planetary bodies. For similar reasons, pure molecular nitrogen lasers typically emit light in the ultraviolet range. Nitrogen also makes a contribution to visible air glow from the Earth's upper atmosphere, through electron impact excitation followed by emission. This visible blue air glow (seen in the polar aurora and in the re-entry glow of returning spacecraft) typically results not from molecular nitrogen, but rather from free nitrogen atoms combining with oxygen to form nitric oxide (NO). b. Reactions

208

Structure of [Ru(NH3)5(N2)]2+.

Nitrogen is generally unreactive at standard temperature and pressure. N2 reacts spontaneously with few reagents, being resilient to acids and bases as well as oxidants and most reductants. When nitrogen reacts spontaneously with a reagent, the net transformation is often called nitrogen fixation. Nitrogen reacts with elemental lithium. Lithium burns in an atmosphere of N2 to give lithium nitride: 6 Li + N2 → 2 Li3N Magnesium also burns in nitrogen, forming magnesium nitride. 3 Mg + N2 → Mg3N2 N2 forms a variety of adducts with transition metals. The first example of a dinitrogen complex is [Ru(NH3)5(N2)]2+ (see figure at right). Such compounds are now numerous, other examples include IrCl(N2)(PPh3)2, W(N2)2(Ph2CH2CH2PPh2)2, and [(η5-C5Me4H)2Zr]2(μ2, η²,η²-N2). These complexes illustrate how N2 might bind to the metal(s) in nitrogenase and the catalyst for the Haber process. A catalytic process to reduce N2 to ammonia with the use of a molybdenum complex in the presence of a proton source was published in 2005. (See nitrogen fixation). The starting point for industrial production of nitrogen compounds is the Haber process, in which nitrogen is fixed by reacting N2 and H2 over an iron(III) oxide (Fe3O4) catalyst at about 500 °C and 200 atmospheres pressure. Biological nitrogen fixation in free-living cyanobacteria and in the root nodules of plants also produces ammonia from molecular nitrogen. The reaction, which is the source of the bulk of nitrogen in the biosphere, is catalyzed by the nitrogenase enzyme complex which contains Fe and Mo atoms, using energy derived from hydrolysis of adenosine triphosphate (ATP) into adenosine diphosphate and inorganic phosphate (−20.5 kJ/mol).

3. Occurrence Nitrogen is the largest single constituent of the Earth's atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air). It is created by

209

fusion processes in stars, and is estimated to be the 7th most abundant chemical element by mass in the universe. Molecular nitrogen and nitrogen compounds have been detected in interstellar space by astronomers using the Far Ultraviolet Spectroscopic Explorer. Molecular nitrogen is a major constituent of the Saturnian moon Titan's thick atmosphere, and occurs in trace amounts in other planetary atmospheres. Nitrogen is present in all living organisms, in proteins, nucleic acids and other molecules. It typically makes up around 4% of the dry weight of plant matter, and around 3% of the weight of the human body. It is a large component of animal waste (for example, guano), usually in the form of urea, uric acid, ammonium compounds and derivatives of these nitrogenous products, which are essential nutrients for all plants that cannot fix atmospheric nitrogen.

Applications •

Liquid : Nitrogen has use for freeze.

•

Nitrogen occurs naturally in many minerals, such as saltpetre

(potassium nitrate), Chile saltpetre (sodium nitrate) and sal ammoniac (ammonium chloride). Most of these are uncommon, partly because of the minerals' ready solubility in water. See also Nitrate minerals and Ammonium minerals.

2.1.2. Phospor History and discovery The discovery of phosphorus is credited to the German alchemist Hennig Brand in 1669, although other chemists might have discovered phosphorus around the same time. Brand experimented with urine, which contains considerable quantities of dissolved phosphates from normal metabolism. Working in Hamburg, Brand attempted to create the fabled philosopher's stone through the distillation of some salts by evaporating urine, and in the process produced a white material that glowed in the dark and burned brilliantly. It was named phosphorus mirabilis ("miraculous bearer of light"). His process originally involved letting urine stand for days until it gave off a terrible smell. Then he boiled it down to a paste, heated this paste to a high temperature, and led the vapours through water, where he hoped they would condense to gold. Instead, he obtained a white, waxy substance that glowed in the dark. Brand had discovered phosphorus, the first element discovered since antiquity. We now know that Brand produced ammonium sodium hydrogen phosphate, 210

(NH4)NaHPO4. While the quantities were essentially correct (it took about 1,100 L of urine to make about 60 g of phosphorus), it was unnecessary to allow the urine to rot. Later scientists would discover that fresh urine yielded the same amount of phosphorus. Since that time, phosphors and phosphorescence were used loosely to describe substances that shine in the dark without burning. However, as mentioned above, even though the term phosphorescence was originally coined as a term by analogy with the glow from oxidation of elemental phosphorus, is now reserved for another fundamentally different process—re-emission of light after illumination. Brand at first tried to keep the method secret,[24] but later sold the recipe for 200 thaler to D Krafft from Dresden,[4] who could now make it as well, and toured much of Europe with it, including England, where he met with Robert Boyle. The secret that it was made from urine leaked out and first Johann Kunckel (1630–1703) in Sweden (1678) and later Boyle in London (1680) also managed to make phosphorus. Boyle states that Krafft gave him no information as to the preparation of phosphorus other than that it was derived from "somewhat that belonged to the body of man". This gave Boyle a valuable clue, however, so that he, too, managed to make phosphorus, and published the method of its manufacture. Later he improved Brand's process by using sand in the reaction (still using urine as base material), 4 NaPO3 + 2 SiO2 + 10 C → 2 Na2SiO3 + 10 CO + P4 Robert Boyle was the first to use phosphorus to ignite sulfur-tipped wooden splints, forerunners of our modern matches, in 1680. In 1769 Johan Gottlieb Gahn and Carl Wilhelm Scheele showed that calcium phosphate (Ca3(PO4)2) is found in bones, and they obtained phosphorus from bone ash. Antoine Lavoisier recognized phosphorus as an element in 1777. Bone ash was the major source of phosphorus until the 1840s. Phosphate rock, a mineral containing calcium phosphate, was first used in 1850 and following the introduction of the electric arc furnace in 1890, this became the only source of phosphorus. Phosphorus, phosphates and phosphoric acid are still obtained from phosphate rock. Phosphate rock is a major feedstock in the fertilizer industry. Occurrence Due to its reactivity with air and many other oxygen-containing substances, phosphorus is not found free in nature but it is widely distributed in many different minerals. Phosphate rock, which is partially made of apatite (an impure tri-calcium phosphate mineral), is an important commercial source of this element. About 50 percent of the global phosphorus reserves are in the Arab nations. Large deposits of apatite are located in China, Russia, Morocco, Florida, Idaho, Tennessee, Utah, and elsewhere. Albright and Wilson in the United Kingdom and their Niagara Falls plant, for instance, were using phosphate rock in the 1890s and 1900s from Connetable, Tennessee and Florida; by 1950 they were using phosphate rock mainly from Tennessee and North Africa. In the early 1990s Albright and Wilson's purified wet phosphoric acid business was being adversely affected by phosphate rock sales by China and the entry of their long-standing Moroccan phosphate suppliers into the purified wet phosphoric acid business. 211

In 2007, at the current rate of consumption, the supply of phosphorus was estimated to run out in 345 years. However, scientists are now claiming that a "Peak Phosphorus" will occur in 30 years and that "At current rates, reserves will be depleted in the next 50 to 100 years. The stability of the +5 oxidation state is illlustrated by the wide range of phosphate materials available in the earth. Applications Match striking surface made of a mixture of red phosphorus, glue and ground glass. The glass powder is used to increase the friction. Widely used compounds

Use

Ca(H2PO4)2·H2O

Baking powder and fertilizers

CaHPO4·2H2O

Animal food additive, toothpowder

H3PO4

Manufacture of phosphate fertilizers

PCl3

Manufacture of POCl3 and pesticides

POCl3

Manufacturing plasticizer

P4S10

Manufacturing of additives and pesticides

Na5P3O10

Detergents

Phosphorus, being an essential plant nutrient, finds its major use as a constituent of fertilizers for agriculture and farm production in the form of concentrated phosphoric acids, which can consist of 70% to 75% P2O5. Global demand for fertilizers led to large increase in phosphate (PO43–) production in the second half of the 20th century. Due to the essential nature of phosphorus to living organisms, the low solubility of natural phosphorus-containing compounds, and the slow natural cycle of phosphorus, the agricultural industry is heavily reliant on fertilizers that contain phosphate, mostly in the form of superphosphate of lime. Superphosphate of lime is a mixture of two phosphate salts, calcium dihydrogen phosphate Ca(H2PO4)2 and calcium sulfate dihydrate CaSO4·2H2O produced by the reaction of sulfuric acid and water with calcium phosphate. • Phosphorus is widely used to make organophosphorus compounds, through the intermediates phosphorus chlorides and two phosphorus sulfides: phosphorus pentasulfide, and phosphorus sesquisulfide.[13] Organophosphorus compounds have many applications, including in plasticizers, flame retardants, pesticides, extraction agents, and water treatment.

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•

Phosphorus is also an important component in steel production, in the making of phosphor bronze, and in many other related products. Phosphorus is added to metallic copper during its smelting process to react with oxygen present as an impurity in copper and to produce oxygen-free copper or phosphorus-containing copper (CuOFP) alloys with a higher thermal and electrical conductivity than normal copper.

•

Phosphates are utilized in the making of special glasses that are used for sodium lamps.

•

Bone-ash, calcium phosphate, is used in the production of fine china.

•

Sodium tripolyphosphate made from phosphoric acid is used in laundry detergents in some countries, but banned for this use in others.

•

Phosphoric acid made from elemental phosphorus is used in food applications such as some soda beverages. The acid is also a starting point to make food grade phosphates. These include mono-calcium phosphate

that

is

employed

in

baking

powder

and

sodium

tripolyphosphate and other sodium phosphates. Among other uses these are used to improve the characteristics of processed meat and cheese. Others are used in toothpaste. Trisodium phosphate is used in cleaning agents to soften water and for preventing pipe/boiler tube corrosion. •

Phosphorus sesquisulfide is used in heads of strike-anywhere matches.

•

In trace amounts, phosphorus is used as a dopant for n-type semiconductors.

•

P and

32

P are used as radioactive tracers in biochemical laboratories

33

(see Isotopes). •

Phosphate is a strong complexing agent for the hexavalent uranyl (UO22+) species and this is the reason why apatite and other natural phosphates can often be very rich in uranium.

•

Tributylphosphate is an organophosphate soluble in kerosene and used to extract uranium in the Purex process applied in the reprocessing of spent nuclear fuel.

2.1.3. Arsen

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Arsenic is the chemical element that has the symbol As, atomic number 33 and atomic mass 74.92. Arsenic was first documented by Albertus Magnus in 1250. Arsenic is a notoriously poisonous metalloid with many allotropic forms, including a yellow (molecular non-metallic) and several black and grey forms (metalloids). Three metalloidal forms of arsenic, each with a different crystal structure, are found free in nature (the minerals arsenic sensu stricto and the much rarer arsenolamprite and pararsenolamprite). However, it is more commonly found as arsenide and in arsenate compounds, several hundred of which are known. Arsenic and its compounds are used as pesticides, herbicides, insecticides and in various alloys.

1. History Arsenic sulfides (orpiment, realgar) and oxides have been known and used since ancient times. Zosimos (circa 300 AD) describes roasting sandarach (realgar) to obtain cloud of arsenic (arsenious oxide) which he then reduces to metallic arsenic. As the symptoms of arsenic poisoning were somewhat illdefined, it was frequently used for murder until the advent of the Marsh test, a sensitive chemical test for its presence. (Another less sensitive but more general test is the Reinsch test.) Owing to its use by the ruling class to murder one another and its potency and discreetness, arsenic has been called the Poison of Kings and the King of Poisons. During the Bronze Age, arsenic was often included in bronze, which made the alloy harder (so-called "arsenical bronze"). Albertus Magnus (Albert the Great, 1193–1280) is believed to have been the first to isolate the element in 1250 by heating soap together with arsenic trisulfide. In 1649, Johann Schröder published two ways of preparing arsenic.

Alchemical symbol for arsenic

Cadet's fuming liquid (impure cacodyl), the first organometallic compound, was synthesized in 1760 by Louis Claude Cadet de Gassicourt by the reaction of potassium acetate with arsenic trioxide.

2. Properties a. Isotopes : Naturally occurring arsenic is composed of one stable isotope, 75As. As of 2003, at least 33 radioisotopes have also been synthesized, ranging in atomic mass from 60 to 92. The most stable of these is 73As with a half-life of

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80.3 days. Isotopes that are lighter than the stable 75As tend to decay by β+ decay, and those that are heavier tend to decay by β- decay, with some exceptions. At least 10 nuclear isomers have been described, ranging in atomic mass from 66 to 84. The most stable of arsenic's isomers is 68mAs with a half-life of 111 seconds. Allotropes

Structure of yellow arsenic As4 and white phosphorus P4

Like phosphorus, arsenic is an excellent example of an element that exhibits allotropy, as its various allotropes have strikingly different properties. The three most common allotropes are metallic grey, yellow and black arsenic. The most common allotrope of arsenic is grey arsenic. It has a similar structure to black phosphorus (β-metallic phosphorus) and has a layered crystal structure somewhat resembling that of graphite. It consists of many six-membered rings which are interlinked. Each atom is bound to three other atoms in the layer and is coordinated by each 3 arsenic atoms in the upper and lower layer. This relatively close packing leads to a high density of 5.73 g/cm3. Yellow arsenic (As4) is soft and waxy, somewhat similar to P4. Both have four atoms arranged in a tetrahedral structure in which each atom is bound to each of the other three atoms by a single bond, resulting in very high ring strain and instability. This form of arsenic is the least stable, most reactive, most volatile, least dense, and most toxic of all the allotropes. Yellow arsenic is produced by rapid cooling of arsenic vapour with liquid nitrogen. It is rapidly transformed into the grey arsenic by light. The yellow form has a density of 1.97 g/cm3. b. Chemical properties The most common oxidation states for arsenic are −3 (arsenides: usually alloy-like intermetallic compounds), +3 (arsenates(III) or arsenites, and most organoarsenic compounds), and +5 (arsenates: the most stable inorganic arsenic oxycompounds). Arsenic also bonds readily to itself, forming square As3−4 ions in the arsenide skutterudite. In the +3 oxidation state, the stereochemistry of arsenic is affected by the presence of a lone pair of electrons. Arsenic is very similar chemically to its predecessor in the Periodic Table, phosphorus. Like phosphorus, it forms colourless, odourless, crystalline oxides As2O3 and As2O5 which are hygroscopic and readily soluble in water to form acidic solutions. Arsenic(V) acid is a weak acid. Like phosphorus, arsenic forms an unstable, gaseous hydride: arsine (AsH3). The similarity is so great that arsenic will partly substitute for phosphorus in biochemical reactions and is thus poisonous. However, in subtoxic doses, soluble arsenic compounds act as stimulants, and were once popular in small doses as medicine by people in the mid 18th century. When heated in air, arsenic oxidizes to arsenic trioxide; the fumes from this reaction have an odour resembling garlic. This odour can be detected on

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striking arsenide minerals such as arsenopyrite with a hammer. Arsenic (and some arsenic compounds) sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state. The liquid state appears at 20 atmospheres and above, which explains why the melting point is higher than the boiling point. 3. Occurrence

A large sample of native arsenic.

Arsenopyrite, also unofficially called mispickel, (FeAsS) is the most common arsenic-bearing mineral. In the lithosphere, the minerals of the formula M(II)AsS, with M(II) being mostly Fe, Ni and Co, are the dominant arsenic minerals. Orpiment and realgar were formerly used as painting pigments, though they have fallen out of use owing to their toxicity and reactivity. Although arsenic is sometimes found native in nature, its main economic source is the mineral arsenopyrite mentioned above; it is also found in arsenides of metals such as silver, cobalt (cobaltite: CoAsS and skutterudite: CoAs3) and nickel, as sulfides, and when oxidised as arsenate minerals such as mimetite, Pb5(AsO4)3Cl and erythrite, Co3(AsO4)2·8H2O, and more rarely arsenites ('arsenite' = arsenate(III), AsO33− as opposed to arsenate (V), AsO43−). In addition to the inorganic forms mentioned above, arsenic also occurs in various organic forms in the environment.Other naturally occurring pathways of exposure include volcanic ash, weathering of the arsenic-containing mineral and ores as well as groundwater. It is also found in food, water, soil and air.

4. Applications a. Wood preservation : The toxicity of arsenic to insects, bacteria, and fungi led to its use as a wood preservative. In the 1950s a process of treating wood with chromated copper arsenate (also known as CCA or Tanalith) was invented, and for decades this treatment was the most extensive industrial use of arsenic. Due to an increased understanding of arsenic's high level of toxicity, most countries banned the use of CCA in consumer products. The European Union and United States led this ban, beginning in 2004. As of 2002, US-based industries consumed 19,600 metric tons of arsenic. 90% of this was used for treatment of wood with CCA. In 2007, 50% of the 5,280 metric tons of consumption was still used for this purpose. In the United States, the use of arsenic in consumer products was discontinued for residential and general consumer construction on December 31, 2003 and alternative chemicals are now used, such as Alkaline Copper Quaternary, borates, copper azole, cyproconazole, and propiconazole. b. Other uses of arsen :

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 Various agricultural insecticides, termination and poisons. For example

Lead hydrogen arsenate was used well into the 20th century as an insecticide on fruit trees. Its use sometimes resulted in brain damage to those working the sprayers. In the last half century, monosodium methyl arsenate (MSMA) and disodium methyl arsenate (DSMA), a less toxic organic form of arsenic, has replaced lead arsenate's role in agriculture.  Used in animal feed, particularly in the US as a method of disease

prevention and growth stimulation. One example is roxarsone which was used by 69.8 and 73.9% of the broiler starter and growers between 1995 to 2000.  Gallium arsenide is an important semiconductor material, used in

integrated circuits. Circuits made using the compound are much faster (but also much more expensive) than those made in silicon. Unlike silicon it is direct bandgap, and so can be used in laser diodes and LEDs to directly convert electricity into light.  Also used in bronzing and pyrotechnics.  Up to 2% of arsenic is used in lead alloys for lead shots and bullets.  Arsenic is added in small quantities to brass to make it dezincification

resistant. This grade of brass is used to make plumbing fittings.  Arsenic is also used for taxonomic sample preservation.  Until recently Arsenic was used in optical glass. Modern glass

manufacturers, under pressure from environmentalists, have removed it, along with Lead.

Trimethylarsine

Arsenobetaine

Inorganic arsenic and its compounds, upon entering the food chain, are progressively metabolised to less toxic forms of arsenic through a process of methylation. For example, the mold Scopulariopsis brevicaulis produce significant amounts of trimethylarsine if inorganic arsenic is present. The organic compound arsenobetaine is found in some marine foods such as fish and algae, and also in mushrooms in larger concentrations. The average person's intake is about 10–50 µg/day. Values about 1000 µg are not unusual following consumption of fish or mushrooms. But there is little danger in eating fish

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because this arsenic compound is nearly non-toxic. Some species of bacteria obtain their energy by oxidizing various fuels while reducing arsenate to arsenite. The enzymes involved are known as arsenate reductases (Arr).

5. Occupational exposures Industries that use inorganic arsenic and its compounds include wood preservation, glass production, nonferrous metal alloys, and electronic semiconductor manufacturing. Inorganic arsenic is also found in coke oven emissions associated with the smelter industry. Occupational exposure and poisoning may occur in persons working in these industries

2.1.4. Antimony Antimony (pronounced Latin: stibium) is a chemical element with the symbol Sb and an atomic number of 51. It has two stable isotopes, one with seventy neutrons, and the other with seventy-two. A silvery lustrous grey metalloid, it is found mainly as antimony sulfide, commonly known as stibnite. Elemental antimony has applications in electronics and as an alloy with other metals it is used for small arms ammunition.

1. History

One of the alchemical symbols for antimony

Antimony's sulfide compound, antimony(III) sulfide, Sb2S3 was recognized in antiquity, at least as early as 3000 BC. An artifact made of antimony dating to about 3000 BC was found at Tello, Chaldea (part of present-day Iraq), and a copper object plated with antimony dating between 2500 BC and 2200 BC has been found in Egypt. There is some uncertainty as to the description of the artifact from Tello. Although it is sometimes reported to be a vase, a recent detailed discussion reports it to be rather a fragment of indeterminate purpose. The first European description of a procedure for isolating antimony is in the book De la pirotechnia of 1540 by Vannoccio Biringuccio, written in Italian. This book precedes the more famous 1556 book in Latin by Agricola, De re metallica, even though Agricola has been often incorrectly credited with the discovery of metallic antimony. A text describing the preparation of metallic antimony that was published in in 1604 purported to date from the early fifteenth century, and if authentic it would predate Biringuccio. The book, written in Latin, was called "Currus Triumphalis Antimonyi" (The Triumphal Chariot of Antimony), and its putative author was a certain Benedictine monk, writing under the name Basilius Valentinus. Already in 1710 218

Wilhelm Gottlob Freiherr von Leibniz, after careful inquiry, concluded that the work was spurious, that there was no monk named Basilius Valentinus, and the book's author was its ostensible editor, Johann Thölde (ca. 1565-ca. 1624). There is now agreement among professional historians that the Currus Triumphalis.. was written after the middle of the sixteenth century and that Thölde was likely its author. An English translation of the "Currus Triumphalis" appeared in English in 1660, under the title The Triumphant Chariot of Antimony. The work remains of great interest, chiefly because it documents how followers of the renegade German physician, Philippus Theophrastus Paracelsus von Hohenheim (of whom Thölde was one), came to associate the practice of alchemy with the preparation of chemical medicines. According to the traditional history of Middle Eastern alchemy, pure antimony was well known to Jābir ibn Hayyān, sometimes called "the Father of Chemistry", in the 8th century. Here there is still an open controversy: Marcellin Berthelot, who translated a number of Jābir's books, stated that antimony is never mentioned in them, but other authors claim that Berthelot translated only some of the less important books, while the more interesting ones (some of which might describe antimony) are not yet translated, and their content is completely unknown. The first natural occurrence of pure antimony ('native antimony') in the Earth's crust was described by the Swedish scientist and local mine district engineer Anton von Swab in 1783. The type-sample was collected from the Sala Silvermine in the Bergslagen mining district of Sala, Västmanland, Sweden.

2. Properties a. Physical properties :

A vial containing a black allotrope of antimony

Native massive antimony with oxidation products

There are four known allotropes of antimony: a stable metallic form, and three meta-stable forms which are explosive, black and yellow. Each has its own distinct physical properties, the most common of which is metallic antimony, a brittle, silver-white shiny metal. When molten antimony is slowly cooled to metallic antimony, it forms with an hexagonal crystal structure, isomorphic with that of the grey form of arsenic. The explosive form of antimony is formed from the electrolysis of antimony(III) trichloride, under specific temperatures and concentration. In a bath of hydrochloric with an antimony anode and platinum foil cathode, explosive antimony is deposited on the latter. When scratched with a sharp implement, an exothermic reaction occurs and white fumes given off as metallic antimony is formed; alternatively, when rubbed with a pestle in a mortar, an strong detonation occurs. Black antimony is formed when gaseous metallic antimony is rapidly cooled. It oxidies in air and is sometimes

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spontaneously combustible. At 100 °C, it gradually transforms into the stable form. Finally, the yellow allotrope of antimony is the most unstable. While it cannot be produced as the black allotrope by rapid cooling, it can only be formed by introducing oxygen into antimony hydride at -90 °C. Above this temperature and in ordinary light, it transforms into the stabler black allotrope. b. Chemistry Antimony trioxide (Sb4O6) is formed when antimony is burnt in an excess of air. In the gas phase, this compound exists as Sb4O6, a species that is retained when cooled to its solid, cubic form. However, in the rhombic form, the molecules polymerise to form chains of [Sb2O3]x. Antimony pentoxide, (Sb4O10) can only be formed by oxidation by concentrated nitric acid. Antimony also forms a mixed-valence oxide, antimony tetroxide (Sb2O4), where it is found in both the +3 and +5 oxidation states. Unlike phosphorus and arsenic, these various oxides are amphoteric and do not form well-defined oxoacids and react with acids to form antimony salts. Antimony trioxide dissolves in concentrated acid to form antimony oxo- (antimonyl) compounds such as SbOCl and (SbO)2SO4. The hypothetical antimonous acid Sb(OH)3 only exists as its salts, [19]:763 such as sodium antimonyte ([Na3SbO3]4), formed by fusing sodium oxide and Sb4O6. Transition metal antimonytes are best described as mixed metal oxides. Antimonyc acid exists only as the hydrate HSb(OH)6, forming salts containing the antimonate anion Sb(OH)−6. Dehydrating metal salts containing this anion yields mixed oxides. Many antimony ores are sulfides, including stibnite (Sb2S3), pyrargyrite (Ag3SbS3), zinkenite, jamesonite, and boulangerite. Antimony pentasulfide is known, but is non-stoichiometric and contains only antimony in the +3 oxidation state. Several complex anions of antimony and sulfur are known, such as [Sb6S10]2− and [Sb8S13]2−. Antimony forms two series of halides: SbX3 and SbX5, where X is one of the halogens. The trihalides SbF3, SbCl3, SbBr3, and SbI3 are all molecular compounds having trigonal pyramidal molecular geometry. The trifluoride SbF3 is prepared by the reaction of Sb2O3 with HF. Sb2O3 + HF → 2 SbF3 + 3 H2O It is a strong Lewis acid that readily accepts fluoride ions to form the complex anions SbF−4 and SbF2−5. Molten SbF3 is a weak electrical conductor. The trichloride SbCl3 is prepared by dissolving Sb2S3 in hydrochloric acid: Sb2S3 + HCl → 2 SbCl3 + 3 H2S The pentahalides SbF5 and SbCl5 have trigonal bipyramidal molecular geometry in the gas phase, but in the liquid phase, SbF5 is polymeric, whereas SbCl5 is monomeric. SbF5 is a powerful Lewis acid used to make the superacid fluoroantimonyc acid (HSbF6), and is an important solvent used in the study of noble gas compounds. Antimony forms antimonydes with metals, such as indium antimonyde (InSb), and silver antimonyde (Ag3Sb). Treating antimonydes with acid produces the unstable toxic gas stibine, SbH3. Sb3− + 3 H+ → SbH3

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Stibine may also be produced by reacting Sb3+ salts with sources of the hydride ion H−. Antimony does not react with hydrogen directly to form stibine.

3.

Occurrence and Production

The abundance of antimony in the Earth's crust is estimated at 0.2 to 0.5 parts per million, comparable to thallium at 0.5 parts per million and silver at 0.07 ppm. Even though this element is not abundant, it is found in over 100 mineral species. Antimony is sometimes found native, but more frequently it is found in the sulfide stibnite (Sb2S3) which is the predominant ore mineral. Commercial forms of antimony are generally ingots, broken pieces, granules, and cast cake. Other forms are powder, shot, and single crystals. In 2005, China was the top producer of antimony with about 84% world share followed at a distance by South Africa, Bolivia and Tajikistan, reports the British Geological Survey. The mine with the largest deposits in China is Xikuangshan mine in Hunan Province with a estimated deposit of 2.1 million metric tons. Antimony is isolated from its ore by a reduction with scrap iron: Sb2S3 + 3Fe → 2Sb + 3FeS Isolating antimony from its oxide, is performed by a charcoal reduction: 2Sb2O3 + 3C → 4Sb + 3CO2 Applications Elemental antimony and alloys Elemental antimony is increasingly being used in the semiconductor industry as a dopant for ultra-high conductivity n-type silicon wafers in the production of diodes, infrared detectors, and Hall-effect devices. It is also used as an alloy, to increase lead's hardness and mechanical strength, as in lead-acid batteries, which is the most common use of antimony. It is used in antifriction alloys, such as Babbit metal. It is used as an alloy in small arms ammunition, buckshot, tracer ammunition, cable sheathing, type metal (e.g. for linotype printing machines), solder – some "lead-free" solders contain 5% Sb in pewter, and in hardening alloys with low tin content in the manufacturing of organ pipes. In the 1950s, tiny beads of a lead-antimony alloy were used to dope the emitters and collectors of NPN alloy junction transistors with antimony. A coin made of antimony was issued in the Keichow Province of China in 1931. The coins were not popular, being too soft and they wore quickly when in circulation. After the first issue no others were produced. Elemental antimony as an antimony pill was once used as a medicine. It could be reused by others after ingestion.

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Treatments principally containing are known as antimonyals and are used as emetics.Antimony compounds are used as antiprotozoan drugs. Antimony potassium tartrate, or tartar emetic, has been used in the past as an anti-schistosomal drug, later replaced by praziquantel. Antimony and its compounds are used in several veterinary preparations like anthiomaline or lithium antimony thiomalate, which is used as a skin conditioner in ruminants. Antimony has a nourishing or conditioning effect on keratinized tissues, at least in animals. Antimony-based drugs, such as meglumine antimonyate, are also considered the drugs of choice for treatment of leishmaniasis in domestic animals. Unfortunately, as well as having low therapeutic indices, the drugs are poor at penetrating the bone marrow, where some of the Leishmania amastigotes reside, and so cure of the disease – especially the visceral form – is very difficult. a. Other uses of antimony : In the heads of some safety matches in nuclear reactors together with beryllium in startup neutron sources; in the form of antimony oxides, antimony sulfides, and antimony trichloride are used in the making of flame-proofing compounds, ceramic enamels, glass, paints, and pottery. Antimony trioxide is the most important of the antimony compounds and is primarily used in flameretardant formulations. These flame-retardant applications include such markets as children's clothing, toys, aircraft and automobile seat covers. It is also used in the fiberglass composites industry as an additive to polyester resins for such items as light aircraft engine covers. The resin will burn while a flame is held to it but will extinguish itself as soon as the flame is removed. 2.1.5. Bismut Bismuth is a chemical element that has the symbol Bi and atomic number 83. This trivalent poor metal chemically resembles arsenic and antimony. Bismuth is heavy and brittle; it has a silvery white color with a pink tinge owing to the surface oxide. Bismuth is the most naturally diamagnetic of all metals, and only mercury has a lower thermal conductivity. It is generally considered to be the last naturally occurring stable, non-radioactive element on the periodic table, although it is actually slightly radioactive. Its only non-synthetic isotope bismuth-209 decays via alpha decay into thallium-205, with an extremely long half-life of 1.9 × 1019 years. Bismuth compounds are used in cosmetics, medicines, and in medical procedures. As the toxicity of lead has become more apparent in recent years,

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alloy uses for bismuth metal as a replacement for lead have become an increasing part of bismuth's commercial importance.

Bismuth crystal with an iridescent oxide surface

Bismuth crystals

Bismuth is a brittle metal with a white, silver-pink hue, often occurring in its native form with an iridescent oxide tarnish showing many colors from yellow to blue. The spiral stair stepped structure of a bismuth crystal is the result of a higher growth rate around the outside edges than on the inside edges. The variations in the thickness of the oxide layer that forms on the surface of the crystal causes different wavelengths of light to interfere upon reflection, thus displaying a rainbow of colors. When combusted with oxygen, bismuth burns with a blue flame and its oxide forms yellow fumes. Its toxicity is much lower than that of its neighbors in the periodic table such as lead, tin, tellurium, antimony, and polonium. Although ununpentium is theoretically more diamagnetic, no other metal is verified to be more naturally diamagnetic than bismuth. (Superdiamagnetism is a different physical phenomenon). Of any metal, it has the second lowest thermal conductivity (after mercury) and the highest Hall coefficient. It has a high electrical resistance. When deposited in sufficiently thin layers on a substrate, bismuth is a semiconductor, rather than a poor metal. Elemental bismuth is one of very few substances of which the liquid phase is denser than its solid phase (water being the best-known example). Bismuth expands 3.32% on solidification; therefore, it was long an important component of low-melting typesetting alloys, where it compensated for the contraction of the other alloying components.

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Though virtually unseen in nature, high-purity bismuth can form distinctive hopper crystals. These colorful laboratory creations are typically sold to collectors. Bismuth is relatively nontoxic and has a low melting point just above 271°C, so crystals may be grown using a household stove, although the resulting crystals will tend to be lower quality than lab-grown crystals. History Bismuth (New Latin bisemutum from German Wismuth, perhaps from weiße Masse, "white mass") was confused in early times with tin and lead because of its resemblance to those elements. Bismuth has been known since ancient times, and so no one person is credited with its discovery. Agricola, in De Natura Fossilium states that bismuth is a distinct metal in a family of metals including tin and lead in 1546 based on observation of the metals and their physical properties. Claude François Geoffroy demonstrated in 1753 that this metal is distinct from lead and tin. "Artificial bismuth" was commonly used in place of the actual metal. It was made by hammering tin into thin plates, and cementing them by a mixture of white tartar, saltpeter, and arsenic, stratified in a crucible over an open fire. Bismuth was also known to the Incas and used (along with the usual copper and tin) in a special bronze alloy for knives. Occurrence and production

Bismite mineral

In the Earth's crust, bismuth is about twice as abundant as gold. It is not usually economical to mine it as a primary product. Rather, it is usually produced as a byproduct of the processing of other metal ores, especially lead, tungsten (China), tin, copper, and also silver (indirectly) or other metallic elements. The most important ores of bismuth are bismuthinite and bismite. In 2005, China was the top producer of bismuth with at least 40% of the world share followed by Mexico and Peru, reports the British Geological Survey. Native bismuth is known from Australia, Bolivia, and China. Recycling While bismuth is most available today as a byproduct, its sustainability is more dependent on recycling. Bismuth is mostly a byproduct of lead smelting, along with silver, zinc, antimony, and other metals, and also of tungsten production, along with molybdenum and tin, and also of copper production. Recycling bismuth is difficult in many of its end uses, primarily because of scattering. Probably the easiest to recycle would be bismuthcontaining fusible alloys in the form of larger objects, then larger soldered objects. Half of the world solder consumption is in electronics (i.e., circuit

224

boards). As the soldered objects get smaller or contain little solder or little bismuth, the recovery gets progressively more difficult and less economic, although solder with a sizable silver content will be more worth recovering. Next in recycling feasibility would be sizeable catalysts with a fair bismuth content, perhaps as bismuth phosphomolybdate, and then bismuth used in galvanizing and as a free-machining metallurgical additive. Finally, the bismuth in the uses where it gets scattered the most, in stomach medicines (bismuth subsalicylate), paints bismuth vanadate on a dry surface, pearlescent cosmetics (bismuth oxychloride), and bismuth-containing bullets. The most important sustainability fact about bismuth is its byproduct status, which can either improve sustainability (i.e., vanadium or manganese nodules) or, for bismuth from lead ore, constrain it; bismuth is constrained. The extent that the constraint on bismuth can be ameliorated or not is going to be tested by the future of the lead storage battery, since 90% of the world market for lead is in storage batteries for gasoline or diesel-powered motor vehicles. The life-cycle assessment of bismuth will focus on solders, one of the major uses of bismuth, and the one with the most complete information. The average primary energy use for solders is around 200 MJ per kg, with the highbismuth solder (58% Bi) only 20% of that value, and three low-bismuth solders (2% to 5% Bi) running very close to the average. The global warming potential averaged 10 to 14 kg carbon dioxide, with the high-bismuth solder about twothirds of that and the low-bismuth solders about average. The acidification potential for the solders is around 0.9 to 1.1 kg sulfur dioxide equivalent, with the high-bismuth solder and one low-bismuth solder only one-tenth of the average and the other low-bismuth solders about average. There is very little life-cycle information on other bismuth alloys or compounds. Chemistry properties Bismuth forms trivalent and pentavalent compounds. The trivalent compounds are more common. Many of its chemical properties are similar to other elements in its group; namely, arsenic and antimony, although it is less toxic than those elements. Bismuth is stable to both dry and moist air at ordinary temperatures. At elevated temperatures, the vapours of the metal combine rapidly with oxygen, forming the yellow trioxide, Bi2O3 On reaction with base, this oxide forms two series of oxyanions: BiO−2, which is polymeric and forms linear chains, and BiO3−3. The anion in Li3BiO3 is actually a cubic octameric anion, Bi8O24−24, whereas the anion in Na3BiO3 is tetrameric. Bismuth sulfide, Bi2S3, occurs naturally in bismuth ores. It is also produced by the combination of molten bismuth and sulfur. Unlike earlier of group 15 elements such as nitrogen, phosphorus, and arsenic, and similar to the previous group 15 element antimony, bismuth does not form a stable hydride analogous to ammonia and phosphine. Bismuth hydride, bismuthine (BiH3), is an endothermic compound that spontaneously decomposes at room temperature. It is stable only below −60°C. The halides of bismuth in low oxidation states have been shown to have unusual structures. What was originally thought to be bismuth(I) chloride, BiCl, turns out to be a complex compound consisting of Bi5+9 cations and BiCl2−5

225

and Bi2Cl2−8 anions. The Bi5+9 cation has a distorted tricapped trigonal prismic molecular geometry, and is also found in Bi10Hf3Cl18, which is prepared by reducing a mixture of hafnium(IV) chloride and bismuth chloride with elemental bismuth, having the structure [Bi+][Bi5+9][HfCl2−6]3. Other polyatomic bismuth cations are also known, such as Bi2+8, found in Bi8(AlCl4)2. Bismuth also forms a low-valence bromide with the same structure as "BiCl". There is a true monoiodide, BiI, which contains chains of Bi4I4 units. BiI decomposes upon heating to the triiodide, BiI3, and elemental bismuth. A monobromide of the same structure also exists. In oxidation state +3, bismuth forms trihalides with all of the halogens: BiF3, BiCl3, BiBr3, and BiI3. All of these, except BiF3, are hydrolysed by water to form the bismuthyl cation, BiO+, a commonly encountered bismuth oxycation.[14] Bismuth(III) chloride reacts with hydrogen chloride in ether solution to produce the acid HBiCl4. Bismuth dissolves in nitric acid to form bismuth(III) nitrate, Bi(NO3)3. In the presence of excess water or the addition of a base, the Bi3+ ion reacts with the water to form BiO+, which precipitates as (BiO)NO3. The oxidation state +5 is less frequently encountered. One such compound is BiF5, a powerful oxidising and fluorinating agent. It is also a strong fluoride acceptor, reacting with xenon tetrafluoride to form the XeF+3 cation: BiF5 + XeF4 → XeF + 3BiF−6 The dark red bismuth(V) oxide, Bi2O5, is unstable, liberating O2 gas upon heating. In aqueous solution, the Bi3+ ion exists in various states of hydration, depending on the pH: pH range Species <3

Bi(H2O)3+6

0-4

Bi(H2O)5OH2+

1-5

Bi(H2O)4(OH)2+

5-14

Bi(H2O)3(OH)3

>11

Bi(H2O)2(OH)4−

These mononuclear species are in equilibrium. Polynuclear species also exist, the most important of which is BiO+, which exists in hexameric form as the octahedral complex [Bi6O4(OH)4]6+ (or 6 [BiO+]·2 H2O). Applications Bismuth oxychloride is sometimes used in cosmetics. Bismuth subnitrate and bismuth subcarbonate are used in medicine. Health •

Bismuth subsalicylate (the active ingredient in Pepto-Bismol and (modern) Kaopectate) is used as an antidiarrheal and to treat some other gastro-intestinal diseases (oligodynamic effect). The means by

226

which this appears to work is still not well-documented. It is thought to be some combination of: •

Killing some bacteria that cause diarrhea

•

Reducing

inflammation/irritation

of

stomach

and

intestinal lining •

Retarding the expulsion of fluids into the digestive system

by irritated tissues, by "coating" them. •

The product Bibrocathol is an organic molecule containing Bismuth and is used to treat eye infections.

•

Bismuth subgallate (the active ingredient in Devrom) is used as an internal deodorant to treat malodor from flatulence (or gas) and faeces.

•

Historically Bismuth compounds were used to treat Syphilis and today Bismuth subsalicylate and Bismuth subcitrate are used to treat the Peptic ulcer.

Other uses •

Strong permanent magnets can be made from the alloy Bismanol (BiMn).

•

Bismuth has a potential role in electronic circuits and in manufacturing next-generation solar cells which would have a greater efficiency. Bismuth allows for the creation of new diodes that can reverse their direction of current flow.

•

Many bismuth alloys have low melting points and are widely used for fire detection and suppression system safety devices.

•

Bismuth is used as an alloying agent in production of malleable irons.

•

It is also used as a thermocouple material.

•

A carrier for U-235 or U-233 fuel in nuclear reactors.

•

Bismuth is also used in solders. The fact that bismuth and many of its alloys expand slightly when they solidify makes them ideal for this purpose.

•

Bismuth subnitrate is a component of glazes that produces an iridescent luster finish.

•

Bismuth telluride (a semiconductor) is an excellent thermoelectric material. Bi2Te3 diodes are used in mobile refrigerators and U coolers. Also used as detectors in Infra red spectrophotometers.

•

A replacement propellant for xenon in Hall effect thrusters 227

•

Bismuth is included in BSCCO (Bismuth Strontium Calcium Copper Oxide) which is a group of similar superconducting compounds discovered in 1988 among which is the highest temperature superconductor yet known with a transition temperature of 110K.[23]

•

Bi-213 can be produced by bombarding radium with bremsstrahlung photons from a linear particle accelerator. In 1997 an antibody conjugate with Bi-213, which has a 45 minute half-life, and decays with the emission of an alpha-particle, was used to treat patients with leukemia. This isotope has also been tried in cancer treatment, e.g. in the Targeted Alpha Therapy (TAT) program.[24]

•

The delta form of bismuth oxide when it exists at room temperature is a solid electrolyte for oxygen. This form normally only exists above and breaks down below a high temperature threshold, but can be electrodeposited well below this temperature in a highly alkaline solution.

•

RoHS initiative for reduction of lead has broadened bismuth's use in electronics as a component of low-melting point solders, as a replacement for traditional tin-lead solders.

•

As noted above, bismuth has been used in lead-free solders; its low toxicity will be especially important for solders to be used in food processing equipment and copper water pipes, although it can also be used in other applications including those in the automobile industry, in the EU for example.

•

A catalyst for making acrylic fibers

•

Ingredient in lubricating greases

•

Dense material for fishing sinkers

2.2. Elements of group VB VB group consists of four elements, namely Vanadium (V), niobium (Nb), Tantalum (Ta). Some properties of the elements of this group we can see in the following : Properties

Vanadium

Niobium

Tantalium

Dubnium

228

Configuration

(V23)

(Nb4i)

(Ta73)

(Dbios)

2, 8, 11, 2

2, 8, 18, 12, 1

2, 8, 18, 32, 11, 2

2, 8, 18, 32, 32, 11, 2

5d3 6s2

5f14 6d3 7s2

Outermost electrons

3d3 4s2

Mass atom (g/mol)

50, 94

92, 91

130, 95

268

6,1

8,4

16,6

-

Melting point (0C)

1900

2468

2996

-

Boiling point (0C)

3450

3300

5425

-

Radius atom (0A)

1,34

1,46

1,49

-

Ionic radius (0A)

0,59

0,70

0,73

-

Inization potensial first

6,74

6,80

7,00

-

Electronegativity

1,6

1,6

1,5

-

-1, 0, +1, +2, +3, +4, +5

-1,+2,+3,+4, +5

-1, +2, +3, +4, +5

+3, +4, +5

Steam Heat

106

-

180

-

Heat of melting

4,2

6,4

6,8

-

Capasity calor (250C)J(mol 0K)

24.89

24.60

25.36

Magnetic properties

Paramagne tic

-

-

-

Electrical properties (N0m)

(20 °C) 197

(0 °C) 152

(20 °C) 131

-

Conductivity thermal (300K) (W.m-1K-1)

30.7

53.7

57.5

-

8.4

7.3

6.3

-

3480 m/s

3400 m/s

-

Period type

Oxidation

4d4 5s1

Thermal Expansion

Speed of sound (On 4560 m/s from wire) (200C)

229

Young’s Modulus

128 GPa

105 GPa

186 GPa

-

Shear Modulus

47 GPa

38 GPa

69 GPa

-

6.7

6.0

6.5

-

Hardness scale Mohs

Outermost elektron V ( 3d3 4S2 ), Nb ( 4d4 4S1 ), Ta ( rd3 4S2 ). Nomer oxsidation varied, stability of +5 oxsidation state increases from V-Nb-Ta. Thus V+5 easily reduced to V+2 is Nb+5 and Ta+5 remained stable, V+5 is a good oxidant. The unique nature of each elements diminished view of the reduced size of cations. Thus, the unique nature of V+4 > V+3 > V+2. Consequently VC14 covalen character. Oxide properties , V205 amphoter but more acidic, while Nb205 and Ta205 fever bases. Unreactive at room temperature but on heating reacts to from VC15, VC14, VC13, dan VI3. Nb and Ta are formed only halide type MX5. All halides are covalen and volatile. With H2 from non-stoichiometric compuonds, VH0,7 ; NbH0,86 and TaH0,76. Tendency from a complex : V > Nb > Ta. Radius (ionic and atomic) Nb with Ta is almost the same as a result lanthanida contraction, so that bolt are almost the same properties. Compounds these metals with an oxidation number lower winnowing colored because of the d orbitals that contain partially.

2.2.1. Vanadium 1. History and discovery : In 1831, Swedish chemist, Niel Grabiol Sefstrom discover new elements in the iron ore in Sweden. The elements was named Vanadium as a means goddess Vanadis gorgeous. Year 1865 Roscor and Thorpe found this elements to be with layers of copper and lawer standstone of Cheshire. Vanadium compounds are abundant in scattered earth’s crust. Some prominent vanadium minerals is : ♣ vanadinite : 3 Pb3(VO4)2 . PbCl2 ♣ carnotite

: K2O . 2UO3 . V2O53H2O

♣ patronite

: V2S5 . 3CuS2

Vanadium also occurs in the clay, rocks, coal and crude oil with a small degree.

230

How to get the vanadium them is by extraction from several compounds, namey : a) From vanadinite, extraction from ore involves several stages: ♣ Separation of PbCl2 Ore is reacted with concentrated HCl, PbCl2 will settle, dioxovandium chlotida (VO2Cl) remained in solution. ♣ Making V2O5 PbCl2 after inseparable, solution plus NH4C1 and saturated with NH3, forming NH4VO3 which when formed V205. ♣ Reduction of V205. V205 reduced by Ca in the 900 - 950 ° C to obtain pure vanadium. b) From carnotite. ♣ The making of sodium orthovanadate. Carnotite melted with Na2C03, the liquid obtain extracted with water to precipitate Fe(OH)3, evaporated and cooled the importance of the Na3V04. ♣ Making V205. Solution containing Na3V04 given NH4C1 and saturated NH3, forming NH4V03 (amonium metavanadate), heated to obtain V205. ♣ Reduction of V205. Rich Mardenand way of vanadium metal was obtained pure.

2. Manufacture of metals : This metal is very hard to obtain in a pure state because the melting and high reactivity toward O2, N2 and C at high temperature. ♣ Vanadium ± 99 % can be obtained by V205 with Al (thermit process). ♣ Pure Vanadium was obtained by reducing VC13 with Na or with H2 at temperature 900 ° C. VC13 obtained from V205 with S2C12 at 300 °C.

231

♣ VC14 reduction with Mg can be obtained 99,3 % vanadium.

3. Aliase vanadium : Commercial products are mainly as aliase vanadium, ♣ Ferro vanadium and Cupro vanadium Both are made by reducing vanadium mixed oxide with Fe or Cu with carbon preformance in the electric furnace. ♣ Nikelo vanadium, made by healting a mixture of V205 + NiO. ♣ Obalto vanadium, created by mixing sediment (from the reactive Na-vanadate solutions with cobalto sulphate) with Na 2C03 electric furnace.

4. Usage : ♣ Addition of 0,1 to 0,3 % V steel will increase the power range. ♣ Vanadium is important for the tools of high speed steel. ♣ V205 used as a catalyst in oxidation naphtalen and also in making H2S04 process.

5. Properties : ♣ Healted H2 (no other gases) at 1100°C to vanadium hydride is stable. ♣ These metal reactive in cold conditions, when heated formed V20 (brown), hested hold formed V203 (black), V204 (blue), end the V205 (orange). These metals burn with a bright flame with oxygen. ♣ When heated with Cl2 fromed dry VC14. ♣ This metal does not reactive with bromine water, HCl/winter, releasing H2 with HF and forming a green solution.

6. Substances : Vanadium form compounds with oxidation numbr +5, +4, +3 and +2. Compounds with lower oxidation number is reducing agent,is unique and colorful.

232

1. Compounds V+5 (colorless) is reducted with a reducing agent according to the following changes :

a. Vanadium pentoksida, V2O5 Made from: ♣ Oxidation/ metal or oxide , toasting with low oxidation states. V2O5 as the final result. ♣ Hyidrolysis VOCl3. ♣ Heating amonium vanadate. Usage : ♣ As a catalyst in the oxidation of SO2 → SO3, in sulfuric acid. V2O5 2SO2 + O2 ↔ 2SO3 ♣ Catalyst in a alcohol oxidation and hydrogenation of olefin. b. Vanadium pentaflourida, VF5. This compounds is expressed as sublimat pure white. Be made with heating VF4 in a nitrogen environment, at a temperature 350°C – 650°C. This compounds is very solube in water or organic solvents. c. Vanadium oxitrikhlorida, VOCl3. This compounds is made by ing Cl2 dried at VO3to be heated. This compounds is yellow with a translucent boiling point 127 ° C. d. Vanadium pentasulfida, V2S5. This compounds is made by heating mixture of vanadium trisulfida, with sulfur without air at 400 ° C. This compounds form black powder.

2. Compound V+4. Compounds with a +4 oxidation state is stable, easy to manufacture.

233

a. Vanadium titroksida, V2O4 atau VO2. Created by heating a mixture of vanadium trioksida and vanadium pentoksida without air with a molar amount same. These compounds form a dark blue cystal, easily soluble acid or alkaline. b. Vanadium titraflourida, VF4. Created from the reaction of HF anhydride with VCl4. Start temperature –28°C and rose slowly to 0°C. Flouride is a brownish yellow powder, soluble in water forming a blue solution.

3. Compounds Vanadil This compound contains a cation vanadil (VO+2) where the number oxidation +4, is unique, colored blue. Vanadil chloride made from the hydrolysis VCl4 VCl4 + H2O → VOCl2 + 2HCl Or from the hydrolysis V2O5 with HCl V2O5 + HCl → 2VOCl2 +3H2O + Cl2 VOCl2 compounds are strong reducing agents that are used commercial in coloration. Only the E° of VO+2/ VO3 is –1 volt.

4. Compound V+3. a. Vanadium trioxside, V2O3. Created by reducing V2O5 with hydrogen. V2O3 is alkaline, soluble in acid gives hezaquo ions, V(H2O)63+. b. oxihalida halide and Vanadium. Vanadium triflourida, VF3.3H2O created when the V2O3 dissolved HF. Trihalida the other is the VCl3 and VBr3, moderate VI3 not known. Vanadium oxihalida is known VOCl and VOBr. Both are insoluble in water but soluble in acid.

5. Compound V+2. Compound V+2 colored and paramagnetic ion V+2 is a strong reducing agent. V+2 dilute solution (violet) to reduce water liberate H2. 234

V+2 + H+ → V+3 + ½ H2 (violet)

(green)

6. Compound V+1, V-1 and V°. Oxidation is not common, stabilized by the ligand acid Π. Oxidation numer +1 was found in compound V(CO)6-1.

2.2.2. Niobium and Tantalum Although these two elements are metals in some respects, chemical properties such as non logam Although these two elements has no cation real, but formed several anions. Halides and very oxihalida hydrolysed volatile and easily.

Niobium

Tantalum

1. History of Niobium and Tantalum Niobium was discovered by British chemist from the Charles Hatchett in the year 1801. Hatchett found in samples of niobium minerals columbite and called the new element columbium. Columbium discovered by Hatchett in a mixture with Tantalum. Tantalum itself was discovered by a chemist Anders Ekeberg Sweden. Then confusion arises between the two element are both are the same element or not. Then in 1809, English chemist William Hyde Wollaston comparing shapes oxidation that occurred columbium – columbite, with a density 5918 g/ml, and tantalum – tantalite with a density 7935 g/ml, and concluded that both elements are different because the course has different densities, so the name became niobium and columbium be Tantalum fixed Tantalum.

235

2.

Properties :

Both metals are very difficult to separate. Niobium metal is thin, soft, grayish, shiny, can be bent, high melting point (Nb= 2468 ° C). Tantalum metal is dark, dense, can be bent, the haerder than Niobium, electrical conductivity and high heat, high melting point (Ta = 2996° C), very resistant to acid. Both can be dissolved with HNO3, HF and dissolve very slowly in the alkaline liquid.

3. Compounds : a) Compound Nb+5 and Ta+5 •

Nb2O5 and Ta2O5

Created by dihydroksioksida hydrate (often called acid niobat or tantalat), or by roasting a particular compound with excess oxygen. Both these compound are in powder from dense, relatively inert chemically, almost did not react with acid except for concentrated HF. These compound can be dissolved with melted with alkaline hydrogen sulphate, alkali carbonate or alkali hydroxide. •

NbX5 and TaX5 (X = halide).

Compounds NbF5 and TaF5 prepared by reaction direct flourinasi metal or pentachlorida. Both are solid white. It’s easy yawn. Melting point of Nb = 80 ° C, Ta = 95 ° C. boiling point Nb = 235 ° C, Ta = 229 ° C, forming a colorless liquid and vapor. Compound other halide yellow to brown, made with the reaction direct metal with excess halogen. Halide is dotted halides liquid and boiling point between 200 - 300 ° C, soluble in organic solvents such us ether, CC14, and so forth.

b) Compound Nb and Ta with low oxidation state. •

Nb02Oxide and TaOx (x = 2 s.d 2,5)

•

Tetrahalida.

All halides known except TaF4. Compound NbF4 not very volatile, paramagnetic. Tetrakhlorida and colored tetrakronida dark brown ao black. Nbl4 can be obtained easily by heating NbI5 up to 300° C. these compounds diamagnetism.

4. Applications  Niobium

236

•

As a material for nuclear powder plant construction.

•

As a mixture of rust resistant metal (example Niobium foil), caused by the presence of Niobium carbide and compound Niobium Nitrit, with the concentration of Niobium in the compound approximately 0.1%.

•

As a superconducting magnet (3 tesla clinical Magnetic resonance imaging scanner), and superconducting radio frequency.

•

The manufacture of coin currency (example Austria 2003, Latvia 2004).

•

In medical equipment, Pace maker.

•

In the manufacture of jewelry.

 Tantalum •

Used in the manufacture of children in a loboratory scale.

•

Used in making electronic devices.

•

In making the camera lens.

•

To produce variations of the alloy that has a point high boiling and the forces of good.

•

Preparation of equipment made of metal carbide.

•

Used in the manufacture of jet engine components.

2.2.3. Dubnium Dubnium is an element of group Vb transition metals are made though nuclesr fusion reactions. This eleemnt was discovered by Albert Ghiorso in 1970. Because of very large nuclei dubnium then dubnium constituents stable and not readily disintegrate. Dubnium elements can br created by firing element amerisium with atoms of neon, adn produces isotopes dubnium isotopes, and quickly decays by emitting in the from of electromagnetic radition. Reaction as follows :

Compound s that can be formed for example Db205 (Dubnium pentoxide), DbX5 (Dubnium Halide), halide complexes Db043" , DbF6", DbF83. Other information about the elements Dubnium not clesrly known.

237

CHAPTER III SUMMARY

1. Groups VA consist of Nitrogen, Phospor, Arsenic, Antimony , and Bismut. Groups VB consist of Vanadium, Niobium, Tantalum, and Dubnium. 2. Nitrogen has use for freeze. 3. Match striking surface made of a mixture of red phosphorus, glue and ground glass. The glass powder is used to increase the friction. 4. The toxicity of arsenic to insects, bacteria, and fungi led to its use as a wood preservative. 5. Elemental antimony is increasingly being used in the semiconductor industry.

6. Bismuth oxychloride is sometimes used in cosmetics. Bismuth subnitrate and bismuth subcarbonate are used in medicine. 7. Outermost elektron V ( 3d3 4S2 ), Nb ( 4d4 4S1 ), Ta ( rd3 4S2 ). Nomer oxsidation varied, stability of +5 oxsidation state increases from V-Nb-Ta. 8. Tendency from a complex : V > Nb > Ta. 9.

Compounds these metals with an oxidation number lower winnowing colored because of the d orbitals that contain partially.

10.Vanadium is important for the tools of high speed steel, V205 used as a catalyst in oxidation naphtalen and also in making H2S04 process. 11. Niobium as a material for nuclear powder plant construction and as a mixture of rust resistant metal. 12.Tantalum used in the manufacture of children in a loboratory scale and used in making electronic devices.

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REFERENCES

Chang, Raymond. 2005. Kimia Dasar Jilid I. Jakarta: Erlangga

Cotton dan Wilkinson. 1989. Dasar Anorganik Kimia. Jakarta : UI-Press

Gianto. 2009. “Sistem Periodik Unsur dan Struktur Atom”. http://arsipegianto.tripod.com/perkspu.pdf, tanggal 20 september 2010, pukul 12:09:10

http://alvina.blog.uns.ac.id/2008/12/30/unsur-golongan-vb/, date September 18th 2010, time 10:17:26 AM

http://indice.blog.uns.ac.id/files/2010/05/unsur-gol-vb.pdf, date September 7th 2010, time 03:48:11 PM

http://www.lenntech.com/periodic/elements/n.htm#ixzz1bXce3Hgh, date September 13th 2010, pukul 07:35:57 PM

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