Chemical Bonding 2t19

A Fourth Grading Project in Advanced Chemistry

JOANNA EVE ALEXANDRA O. RAMOS REGINE J. REMOROZA GLENN OLIVER L. FERRER ALLAN L. ESCANILLA RAMDOLF GENER IV - Diamond

Lewis dot symbol consists of the symbol of an element and one dot for each valence electron in

an atom of the element.

Covalent chemical bonds involved the sharing of a pair of valence electrons by two atoms, in contrast to the transfer of electrons in ionic

bonds. Hydrogen gas forms the simplest covalent bond in the diatomic gases by forming covalent

bonds. The nitrogen and oxygen which makes up the bulk of the atmosphere also exhibits covalent bonding in forming diatomic molecules.

Electrons are shared between two atoms with the same electronegativity values. Difference = 0 Examples: N2

Br2

Electrons are shared nonmetal atoms.

between different

Examples: O-Cl

O-S

N-Cl

Covalent bonding can be visualize with the aid of Lewis Diagrams

LEWIS DIAGRAMS – the inner closed shells of electrons can be considered as included in chemical symbol for the element, and the outer shell or valence electrons are represented by dots.

THE CONCEPT OF RESONANCE 

Our drawing of the Lewis Structure for ozone satisfied the octet rule for the central atom because we placed a double bond between it and one of the two end 0 atoms. In fact, we can put the double bond at either end of the molecule, as shown by these two equivalent Lewis structures: O == O – O

O – O ==O

The octet rule refers to the tendency of atoms to prefer to have eight electrons in

their valence shell. When atoms have less than eight electrons, they tend to react and form more

stable compounds. An octet corresponds to an electron configuration ending with s2p6.

INCOMPLETE OCTET The number of electrons surrounding the central atom in a stable is fewer than eight.

Example: B2O3

The oxygens follow the octet rule, but the borons have only six electrons each. Each bond consists of two electrons being shared.

Example 2: 1.BeH2

Be

¨ ¨

¨

F

¨

¨

F

B

¨

F

¨

3.BF3

¨ ¨

Cl

¨

Mg

¨

Cl

¨



¨ ¨

2.MgCl2

H

¨

H

Some elements, especially nitrogen, have an odd number of electrons and will form somewhat stable elements. Nitric oxide has the formula NO. No matter how electrons are shared between the nitrogen and oxygen atoms, there is no way for nitrogen to have an octet. It will have seven electrons instead. An atom with an unpaired electron is called a free radical and is highly reactive.

Nitrogen dioxide has an unpaired electron. (Note the positive charge above the N).

In Period 3, the elements on the right side of the periodic table have empty d orbitals. The d orbitals may accept electrons, allowing elements like sulfur and phosphorus to have more than an octet. Compounds such as PCl5 and SF6 can form. These compounds have 10 and 12 electrons around their central atoms, respectively.

Xenon hexafluoride uses delectrons to form more than an octet. This compound shows another exception: a noble gas compound.

EXAMPLES: 1.SF6 F F F S F F F

¨ ¨¨ ¨¨ ¨¨ ¨¨

¨ ¨ ¨ ¨

¨ ¨ ¨ ¨ ¨¨ ¨

2. PCl5

¨ ¨

¨ ¨ ¨¨ ¨

¨

¨ ¨

Cl

Cl

P

Cl

¨ ¨ Cl

Cl

Illustrate the Chemical bonds of the following compounds:

1.) Water 2.) Ethylene 3.) Acetylene 4.) Sulfate 5.) Phosphate 6.) Carbonate 7.) Phosphite

1.)

2.)

3.)

4.)

5.)

6.)

7.)

VSEPR

s

for

the

geometric

arrangements of electron pairs around a central

atom in of the electrostatic repulsion between electron pairs.

N o.

Basic Geometry 0 lone pair

1 lone pair

2 lone pairs

3 lone pairs

2

Linear 3 Trigonal Planar

Bent

Tetrahedral

Trigonal Pyramid

Bent

Trigonal bipyramid

Seesaw

T-shaped

4

5

Linear

N o.

Basic Geometry 0 lone pair

1 lone pair

Octahedral

Square pyramid

Pentagonal byramid

Pentagonal pyramid

2 lone pairs

6

7

Square planar

3 lone pairs

Geometry

Examples

Molecular Type

Shape

AX1En

Diatomic

AX2E0

Linear

AX2E1

Bent

NO2−, SO2, O3

AX2E2

Bent

H2O, OF2

AX2E3

Linear

XeF2, I3−

HF, O2 BeCl2, HgCl2, C O2

Molecular Type

Shape

Examples

Geometry

BF3, CO32−, NO3 −, SO 3

AX3E0

Trigonal Planar

AX3E1

Trigonal Pyramidal

NH3, PCl3

AX3E2

T-shaped

ClF3, BrF3

AX4E0

Tetrahedral

AX4E1

Seesaw

CH4, PO43−, SO42 −, ClO − 4

SF4



Determine what shape of molecule shown on each illustration:

1.

2.

3.

4.

1.

Square planar

2.

Square pyramidal

3.

Tetrahedral

4.

Trigonal planar

HYBRIDIZATION OF ATOMIC ORBITALS The solution to the Schrodinger Equation provides the wavefunctions for the following atomic orbitals: 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f, etc. For atoms containing two or more electrons, the energy levels are shifted with respect to those of the H atom. An atomic orbital is really the energy state of an electron bound to an atomic nucleus. The energy state changes when one atom is bonded to another atom.

At this level, we consider the following hybrid orbitals:  sp sp2 sp3 sp3d sp3d2 



The sp hybrid atomic orbitals are possible states of electron in an atom, especially when it is bonded to others. These electron states have half 2s and half 2p characters. From a mathematical view point, there are two ways to combine the 2s and 2p atomic orbitals: sp1 = 2s + 2p sp2 = 2s - 2p



he energy states of the valence electrons in atoms of the second period are in the 2s and 2p orbitals. If we mix two of the 2p orbitals with a 2s orbital, we end up with three sp2 hybridized orbitals. These three orbitals lie on a plane, and they point to the vertices of a equilateral triangle as shown here.When the central atom makes use of sp2 hybridized orbitals, the compound so formed has a trigonal shape. BF3 is such a molecule:



Mixing one s and all three p atomic orbitals produces a set of four equivalent sp3 hybrid atomic orbitals. The four sp3 hybrid orbitals points towards the vertices of a tetrahedron, as shown here in this photograph.When sp3 hybrid orbitals are used for the central atom in the formation of molecule, the molecule is said to have the shape of a tetrahedron.

 Metallic  Ionic

Bond

Bond

 Covalent

Bond

 Hydrogen

Bond

The properties of metals suggest that their atoms possess strong bonds, yet the ease of conduction of heat and electricity suggest that

electrons can move freely in all directions in a metal. The general observations give rise to a picture of “positive ions in a sea of electrons” to

describe metallic bonding.



Formed between atoms of metallic elements.



Electron cloud around atoms.



Good conductor at all states, lustrous, very high melting points.

In chemical bonds, atoms can either transfer or share their valence electrons. In the extreme case where one or more atoms lose electrons and

other atoms gain them in order to produce a noble gas electron configuration, the bond is called an ionic bond. Typical of ionic bonds are those in the alkali halides such as sodium chloride (NaCl).



Crystalline solids (made of ions)



High melting and boiling points



Conduct electricity when melted



Soluble in water but not in nonpolar liquid

Hydrogen bonding differs from other uses of the word “bond" since it is a force of attraction between a hydrogen atom in one molecule and a small atom of high electronegativity in another molecule. That is, it is an intermolecular force, not an intermolecular force as in the common use of the word bond. Hydrogen bonding has a very important effect on the properties of water and ice. Hydrogen bonding is also very important in proteins and nucleic acids and therefore in life processes. The "unzipping" of DNA is a breaking of hydrogen bonds which help hold the two strands of the double helix together.

A.

B.

C.

Water molecules are assymetrical. The positively-charged portions of one are attracted to the negatively-charged parts of another. It takes a lot of energy to pull them apart. The assymetrical charge distribution on water molecule makes it very effective in dissolving ionically-bonded materials. However, it is not an effective solvent of covalently bonded materials. When water freezes, it assumes a very open structure and actully expands. Most materials shrink when they freeze and sink in their liquid phases.



Ionic bonding holds rocks and minerals together.



Covalent bonding organisms together.



Metallic bonding holds civilization together.



Hydrogen bonding gives water its heat-retaining and solvent properties.

holds

people

and

other