Water And Wastewater Analysis: Ph, Acidity, Alkalinity And Hardness 5z4957

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| Intensity factor of acidity and indicates ± hydrogen ion activity ± intensity of acidic or basic character of a solution at a given temperature ± N/10 solution of H2SO4 and of acitic acid do not show same pH (depends on dissociation and H+ ion release)

Most important and most frequently measured parameter Neutralization, softening, coagulation, precipitation, disinfection, corrosion control, etc., aspects of water supply and wastewater treatment are pH dependent Buffer capacity: Amount of strong acid or base needed to change pH of 1 liter of sample by one pH unit

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At neutral pH pH is defined as







Ion product of water



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pKw is constant for a given temperature Neutral pH varies with temp. (7.5 at 0uè & 6.5 at 60uè) If pH increases pOH decreases and vice-versa Natural water pH is in the range of 4-9 Natural waters are slightly basic due carbonates and bicarbonates Relationships exist between pH, acidity and alkalinity

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| | pH meter is used ± Involves potentiometric measurement of hydrogen ion activity ± capable of reading both pH and millivolts ± A pH meter with good electrodes measures pH with 0.1 pH units accuracy under normal conditions

pH meter has ± ± ± ±

A potentiometer A glass electrode A reference electrode A temperature compensation device

pH meters usually have two controls ± Intercept control ± parallelly shifts the response curve, between emf and pH, for giving 0 emf with pH 7 buffer ± Slope control ± rotates response curve about isopotential point

| | lass electrode ± A sensor electrode ± Electro motive force (emf) produced in the glass electrode system linearly varies with the pH of the sample ± Using buffers of known pH values emf is measured by glass electrode system and plotted against pH for calibrating the meter ± With the calibrated meter emf produced by the sample is measured and pH is estimated by extrapolation and interpolation

Reference electrode ± ± ± ±

A half cell providing constant electrode potential èalomel electrode or silver: silver chloride electrode is used Has a liquid junction The electrode is filled by an electrolyte to proper level to ensure proper wetting of the liquid junction

èombination electrode: both glass electrode and reference electrode are incorporated into a single probe

* Keep the electrodes wet when the pH meter is not in use ± Follow manufacturer¶s instructions ± Use tap water with conductivity >4000 µmhos/cm rather than distilled water for short-term electrode storage ± pH 4 buffer is best for glass electrode storage ± Saturated Kèl solution is good for glass electrodes and combination electrodes

Before use remove electrodes from storage solution, rinse with distilled water, and blot dry with soft tissue ± Rinsing and blotting dry are also needed for electrode transfer from one solution to the next

Prior to use, conditioning of the electrode in a small portion of the sample for a minute is recommended ± in case of poorly buffered samples the conditioning can be in 3 or 4 successive portions of the sample ± The conditioned electrode is not rinsed, it is only blot dried

è | Transfer electrode(s) into a standard buffer of neutral pH and set isopotential point on the meter (point of 0¶ emf) Transfer electrode(s) into 2nd standard buffer of pH within 2 units from the sample ± Ensure same temperature for both sample and 2nd buffer ± Record temp. and adjust temp. on the meter ± Adjust meter pH to that of the buffer

Transfer electrode(s) to 3rd standard buffer of pH <10 and within 3 pH units from the sample ± èheck if the meter pH is within 0.1 units of the actual - If not then the pH meter is faulty

èalibration frequency ± needed prior to each set of pH measurements ± If pH values vary widely within a set, check with a 3rd buffer of pH within 1 or 2 units from the sample is needed

| Standard buffer solutions of known pH are needed for pH meter calibration ± èommercially available buffer tablets, powders or solutions can be used ± can be prepared in the laboratory in distilled water ± 10.12 g of potassium hydrogen phthalate in 1000 ml solution made in distilled water gives 4.004 pH at 25uè ± 2.092 g of NaHèO3 and 2.64 g of NaèO3 in 1000 ml solution made in distilled water gives 10.014 pH at 25uè ± Distilled water with <2µmhos/cm conductivity is used after boiling and cooling (pH should be 6-7 after addition of a drop of saturated Kèl solution per 50 ml)

pH of buffer solutions change with temperature ± One can refer standard tables for pH of various buffer solutions at different temperatures

| To know whether the problem is with the meter ± Disconnect electrodes using short-circuit strap ± èonnect reference electrode terminal to glass electrode terminal ± Observe pH change with calibration knob adjustment - rapid and even response over wide range indicates no problem with the meter ± Switch to milliVolt scale ± if the meter reads zero then there is no problem with the meter

To know whether the problem is with the electrode pair ± Substitute one electrode at a time and cross check with two buffers (4 pH units apart) ± deviation <0.1 pH units indicates no problem with the electrode

| Failure of glass electrode ‡ Scratches, deterioration or accumulation of debris can be responsible ‡ Rejuvenate the electrode by alternatively immersing 3 times in 0.1 N NaOH and 0.1 N Hèl ‡ If not rejuvenated then do ± immerse the electrode in KF solution for 30 seconds ± soak in PH 7 buffer overnight ± rinse and store in 7 pH buffer ± rinse in distilled water prior to use

‡ KF solution: 2 g of KF in 2 ml conc. H2SO4, dilute resultant solution to 100 ml with distilled water If protein coat is suspected on glass electrode remove it by soaking in 10% pepsin solution adjusted to 1-2 pH

| èhecking of the reference electrode for failure ± Plug another same type of reference electrode in good condition into glass electrode jack and shift meter to read millivolts ± Dip both electrodes first in same electrolyte and read meter ± Dip both electrodes then in the same buffer solution and read the meter ± meter reading of 0±5 mV indicates no problem with electrode

‡ èhecking can also be done with a different type of electrode (Ag: Agèl electrode with calomel electrode & vice-versa) Meter reading of 44 5 mV indicates no problem with electrode

èlogged junction can be the cause for the problem Problem may be visible as increased response time and/or as drifts in the reading Applying suction or boiling the electrode tip in distilled water can clear the c the junction èlogged junction can also be replaced

ù | Use of special low sodium error electrodes is needed for measuring pH >10, at high temperature, accurately Liquid membrane electrodes are good for pH below 1 pH measurement can not be made accurately in nonaqueous media, suspensions, colloids or high ionic strength solutions Temperature can affect pH measurement by ± èhanging the properties of the electrodes ± Brining in chemical equilibrium changes

Buffer solutions deteriorate from mold growth or contamination hence replace them every 4 weeks once Store buffer solutions and samples in polyethylene bottles

Definition: Quantitative capacity of a water or wastewater sample to react with a strong base till neutralized to a designated pH An aggregate property contributed by strong mineral acids, weak acids (carbonic acid, acetic acid, etc.) and hydrolyzing salts (iron and aluminum sulfates) ± Strong acids are essentially neutralized completely at 4 pH ± èarbon dioxide will not depress pH below 4 and its neutralization is completed at 8.5 pH ± Presence of the acidity contributed by hydrolyzing salts is indicated by formation of precipitate during neutralization

Acids contribute to corrosiveness, influence chemical reaction rates, chemical separation and biological processes, and reflect change in source quality of water

Ä ‡ Recording of pH after successive small measured additions of titrant and plotting against cumulative addition of the titrant ‡ Inflection points on the curve accurately indicate end points ± Appropriate for samples with single acidic species ± but not for buffered and complex mixtures ± some arbitrary pH is used as an end-point (3.7 pH for mineral acidity and 8.3 for total acidity) ± èolour change of an indicator can indicate end points bromocresol blue and methyl orange for 3.7 pH phenolphthalein and metacresol purple for 8.3 pH

‡ Using titration curve acidity of the sample with respect to any pH can be known and buffering capacity at different pH values can be known

Ä | 8.3 is accepted as a standard end point for titration of total acidity for unpolluted surface water samples ± èorresponds to stoichiometric neutralization of carbonic acid into bicarbonate ± Most of the weak acids are neutralized ± èolour change of phenolphthalein or metacresol purple indicator can be used to indicate the end point

3.7 pH and 8.3 pH are used in standard acidity determinations for more complex or buffered solutions ± bromocresol blue or methyl orange for 3.7 pH is used as indicators ± Acidity to pH 3.7 is methyl orange acidity (mineral acidity) and that to pH 8.3 is phenolphthalein acidity or total acidity

‡ Dissolved gases contributing acidity/alkalinity (èO2, H2S, NH3) can be lost/gained during sampling, storage or titration ± Titrate promptly after opening the sample container ± Protect the sample from atmosphere during titration - avoid vigorous shaking or mixing ± Do not allow the sample to become warmer than it was at the time of collection

‡ Oily matter, suspended solids, precipitates and other waste matter can coat the glass electrode and cause sluggish response ± Pause between successive titrant additions ± èlean the electrodes occasionally

‡ Oxidizable or hydrolysable ions (of iron, aluminum, manganese) can cause drifting end points ± Boil the sample at <4 pH with hydrogen peroxide for a few minutes

‡ èoloured or turbid samples can obscures colour change ± Avoid indicator titrations

‡ Free residual chlorine can bleach the indicator ± Eliminate the interference by adding a drop sodium thiosulfate

‡ Sodium carbonate present in sodium hydroxide can introduce errors associated with the neutralization of carbon dioxide = ï

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‡ Sample collection and storage ± èollect samples in polyethylene or borosilicate glass bottles ± Store samples at low temperatures to minimize microbial action and loss/gain of èO2 and other gases ± Analyse sample within a day (if biological activity suspected analyse within 6 hours)

‡ èalibrated pH meter for titration (indicators can be alternative) ± Use pH meter that can read 0.05 pH unit ± If not having temperature compensation provision then titrate at 25±5uè

‡ Use sufficiently large titrant volume & sufficiently small sample volume for volumetric precision & for sharp end points If sample¶s acidity is <1000 mg/l then use sample volume with acidity <50 mg and titrate with 0.02N NaOH If sample¶s acidity is >1000 mg/l then use sample volume with acidity <250 mg and titrate with 0.1N NaOH

Hot peroxide treatment ± Required by samples containing hydrolysable metal ions (iron, aluminum or manganese), or reduced forms of polyvalent cation (iron pickle liquors, acidmine drainage, & other industrial wastes) ± pH of a measured volume of sample is reduced to below 4 by adding 5 ml increments of 0.02N H2SO4 ± Add 5 drops of 30% H2O2 and boil for 2 to 5 minutes ± èool sample to room temperature and titrate with standard alkali to 8.3 pH ± Make correction for the acid added for pH adjustment

Titration to end point using indicator ± If sample is suspected to have residual chlorine first add a drop of 0.1M Na2S2O3 and then add 5 drops of the indicator solution ± Titrate over a white surface to a persistent colour change

Potentiometric titration curve ± Measure sample pH ± Add standard alkali in increments of 0.5 ml or less (change of pH with each incremental addition should be <0.2) ± After each addition mix the sample gently with magnetic stirrer and record pH when constant reading is obtained ± èontinue incremental addition till pH 9 is reached ± èonstruct titration curve (pH versus cumulative titrant added) ± Determine acidity relative to a particular pH from the titration curve

Potentiometric titration to pH 3.7 or 8.3 ± Titrate the sample to preselected end point pH and record the amount of titrant added A is ml of NaOH



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è is ml of H2SO4

B is normality of NaOH D is normality of H2SO4

± Report acidity as ³The acidity to pH ------- = ------- mg èaèO3/L´

è Nomographic chart can be used ± Requires pH, alkalinity, dissolved solids and temperature of the sample for the measurement

èan also be measured from the following expression





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Here [H2èO3 is sum of carbonic acid and free èarbon dioxide (99% is free èO2)

± If pH is not accurately measured èO2 measurement will be erroneous (25% error in èO2 measurement for 0.1 unit error in pH measurement)

‡ Acid neutralizing capacity of sample ‡ Aggregate property of water (contributed by strong and weak bases and by salts of weak acids) ± Most of the alkalinity is contributed by hydroxides, carbonates and bicarbonates - hydroxide and bicarbonate may not coexist ± Borates, phosphates, silicates, etc. of weak acids; salts of humic acid; and H2S and ammonia can also contribute

‡ Boiler water and soft water from lime-soda process may contain hydroxide and carbonate alkalinity ‡ èhemically treated waters are usually alkaline (have high pH) ‡ Supernatant of properly operating anaerobic digesters has high alkalinity (2000-4000 mg/l) ‡ Water bodies with high algal activity can be alkaline (removal of carbon dioxide turns water alkaline and raises pH to 9-10)

‡ Is of little public health concern (highly alkaline water could be unpalatable) ‡ Alkalinity contributed by alkaline earth metals makes water not fit for irrigation ‡ Alkalinity (specially by salts of weak acids and strong bases) serves as a buffer system in water bodies ‡ Alkalinity measurements are used in the interpretation and control of water and wastewater treatment processes

‡ Measured titrimetrically and reported in mg/l as èaèO3 ± Titrated in 2 stages (first to 8.3 pH and then to 4.5 pH) and reported as phenolphthalein and total alkalinities respectively ± Inflection points on titration curve are used as end points

‡ Phenolphthalein alkalinity: ± End point of 8.3 pH corresponds to total conversion of carbonate into bicarbonate ± Phenolphthalein (pink to colourless) or metacresol purple (sharp colour change) as indicator ± Total of hydroxide, half of carbonate and none of bicarbonate

‡ Total alkalinity ± End point of 4.5 pH corresponds to total conversion of bicarbonate into carbonic acid ± Methyl orange/bromocresol green (green end point) as indicator ± Represents total of hydroxide, carbonate and bicarbonate

* Methods of estimation 1. 2. 3.

Estimation from phenolphthalein and total alkalinities Estimation from alkalinity and pH measurements Estimation from equilibrium equations

Estimation from phenolphthalein alkalinity (P) and total alkalinity (T) ‡ Sample can have any of the 5 combinations of alkalinities ± ± ± ± ±

only hydroxide only bicarbonate only carbonate hydroxide + carbonate carbonate + bicarbonate



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* * ù ‡ Sample¶s pH is indicative of the type(s) of alkalinity it has ± Samples with hydroxide alkalinity have pH >10 ± Samples with carbonate alkalinity have pH >8.3 ± Samples with only bicarbonate alkalinity have pH <8.3

‡ Types of alkalinities are related to P and T alkalinities as ± P = 0 means Bicarb Alk. = T (and no other alkalinities) ± P < 0.5T means èarb. Alk. = 2P and Bicarb. Alk. = T-2P ± P=0.5T means èarb Alk. = 2P (and no other alkalinity) ± P>0.5T means Hydrox. Alk. = 2P-T and èarb. Alk. = 2(T-P) ± P=T means Hydrox. Alk. = T (and no other alkalinities)

* Estimation from P and T alkalinities ‡ Presupposes incompatibility of hydroxide and bicarbonate alkalinities ± èarbonate alkalinity is present when P-alkalinity is present and is less than the T-alkalinity ± Hydroxide alkalinity is present if P is >0.5 of T ± Bicarbonate alkalinity is present if P is less than half of T

‡ For samples with pH >9 the estimations are only approximations ‡ Assumes absence of other inorganic or organic acids, such as, silicic, phosphoric and boric acids

* Estimation from alkalinity and pH measurements ‡

From measured pH hydroxide alkalinity is measured by ï å å !





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One mole of OH- is equivalent to 50,000 mg of èaèO3

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At 24uè pKw is 14, at 0uè it is 14.94 and at 40uè it is 13.5

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Temperature measurement is needed to select pKw

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Nomographs can be used to read Hydrox. Alk. from pH and temperature measurements

èarb. Alk. = 2(P ± Hydrox. Alk.) Bicarb. Alk. = T ± (èarb. Alk. ± Hydrox. Alk.)

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Accurate measurement of pH is needed for hydroxide alkalinity estimation

* * Measurement of pH, total alkalinity, TDS and temperature are needed Sum of equivalent concentrations of cations must be equal to the sum of equivalent concentrations of anions From pH and temperature measurements [H+] and [OH-] can be estimated



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è V At 20uè KA2 is 4.7X10-11 = V= Temp. and ionic concentrations influence

èarb. Alk. and Bicarb. Alk. èan be measured by

the value of KA2 Nomographs can be used to know KA2 from temp and TDS measurements



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Sample should not be filtered, diluted, concentrated or altered prior to alkalinity measurement For selecting method determine acid consumption for changing pH by 0.3 units on a sample portion ± Reduction of pH by 0.3 units corresponds to doubling of hydrogen ion concentration

Sample size and titrant (sulfuric acid or hydrochloric acid) strength ± Preliminary titration can help in the selection ± For low alkalinity, sample size is 200 ml and titrant strength is 0.02N H2SO4 otherwise the strength is 0.1N

Report alkalinity <20 mg/l only if it is measured by low alkalinity method

Hardwater produce scales in hot water pipes, heaters, boilers and other units ± ù Sodium when present in higher concentration through common ion effect interferes with the normal behavior of soap

Surface waters are softer than ground waters ‡ Hardness in water is derived largely from with soil and rock formations ‡ roundwater hardness reflects the nature of geological formations with which water came in <75 mg/L of hardness ± soft water 75 - <150 mg/L hardness ± moderately hard water 150 - <300 mg/L hardness ± hard water >300 mg/L hardness ± very hard water

Sum of the calcium and magnesium concentration expressed as èaèO3 in mg/l ± Knowing calcium and magnesium hardness is important to calculate chemical requirement for water softening by limesoda process

A measure of the capacity to precipitate soap ± chiefly by calcium and magnesium ions ± other polyvalent cations also precipitate soap ± Principal hardness causing cations ± calcium, magnesium, strontium, ferrous iron & manganous ions ± Associated anions ± bicarbonates, sulfates, chlorides, nitrates and silicates

èalcium hardness and magnesium hardness and total hardness èarbonate hardness (temporary hardness): hardness equivalent to carbonate and bicarbonate alkalinity ± Tends to precipitate at elevated temperatures

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Non carbonate (permanent) hardness: amount in excess of carbonate and bicarbonate alkalinity ± Associated with sulfate, chloride and nitrate anions ± Samples with hardness less than carbonate and bicarbonate alkalinity have no non-carbonate hardness

Hardness by calculation ‡ Separate determination of calcium and magnesium concentration by mineral analysis ‡ Hardness estimation by





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Ä Eriochrome Black T or èalmagite when added to water sample containing èa+2 and Mg+2 ions gives wine red colour at a pH 10 0.1 The colour is due to weak complexing of the indicator with èa+2 and Mg+2 ions

When titrated with EDTA, the èa+2 and Mg+2 are complexed with EDTA (chelated soluble complex) ± When all èa+2 and Mg+2 are complexed the sample turns to blue colour from wine red ± Indicator indicates presence of excess EDTA

Satisfactory end point needs presence of magnesium ions ± Small amount of complexometrically neutral magnesium salt of EDTA is added to the buffer for this purpose

Dissolve 16.9 grams of ammonium chloride in 143 ml of concentrated ammonium hydroxide, add 1.25 grams of magnesium salt of EDTA and dilute to 250 ml with distilled water ± In place of Mg-EDTA 1.179 grams disodium salt of ethylenediaminetetraacetic acid dihydrate and 780 mg of magnesium sulfate or 644 mg of of magnesium chloride dissolved in 50 ml distilled water can also be used

‡ Store in a plastic or borosilicate glass container for upto one month while stopper tightly to avoid loss of ammonia or pickup of èO2 ‡ Adding 1 or 2 ml to the sample should produce 10 0.1 pH at the titration end point

Odourless buffers as alternative ± To 55 ml conc. Hèl in 400 ml distilled water, while stirring, slowly add 300 ml of 2-aminoethanol, then add 5 g MgEDTA and adjust volume to 1 liter with distilled water

èommercial preparations incorporating buffer and complexing agent (with heavy metal interferences) are available ± supposed to maintain pH 10±0.1 during titration and give clear & sharp end point

* * Standard 0.01M solution of disodium ethylenediaminetetraacetate dihydrate is used ± Store the solution in polyethylene or borosilicate glass bottles (can extract hardness causing cations from soft-glass containers)

Standardized against standard calcium solution Standard calcium solution (1 ml = 1 mg èaèO3) ± Take 1 g anhydrous èaèO3 powder in erlenmeyer flask, add 1:1 Hèl until dissolved and add 200 ml distilled water ± Boil for a few minutes to expel èO2 ± Add a few drops of methyl red indicator and adjust colour to intermediate orange by adding 3N NH4OH or 1:1 Hèl ± Transfer contents quantitatively and dilute with distilled water to one liter

@ Eriochrome Black T ± 0.5 g sodium salt of 1-(1-hydroxy-2-naphthylazo)-5-nitro-2naphthol-4-sulfonic acid in 100 g triethanolamine or ethylene glycol monomethyl ether ± Use 2 drops per 50 ml solution

èalmagite ± 0.1 g 1-(1-hydroxy-4-methyl-2-phenylazo)-2-naphthol-4sulfonic acid in 100 ml distilled water ± Use 1 ml per 50 ml solution

Deteriorate with aging and give indistinct end points ± Alkaline solutions of Eriochrome Black T are sensitive to oxidants ± Aqueous or alcoholic solutions are unstable

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Some metal ions interfere by causing fading, by indistinct end points or by stoichiometric consumption of EDTA ± èertain inhibitors (inhibitor-1 or inhibitor-2) can be added before titration to reduce interferences ± but these are toxic or malodorous

@ adjust sample pH to >6 and add 250 mg sodium cyanide powder and raise pH to 10.0±0.1 @ 5 g sodium sulfide nonahydrate (Na2S.9H2O) or 3.7 g Na2S.5H2O in 100 ml distilled water ± Produces sulfide precipitates of heavy metals ± Exposure to air oxidizes the inhibitor ± Use is not desirable when polyphosphate level in the sample is high (>10 mg/l)

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Mg-èDTA (Magnesium salt of 1, 2-èyclohexane diamine tetra acetic acid) can be added to take care of the interferences ± Add 250 mg to 100 ml sample and dissolve completely prior to adjusting pH to 10.0±0.1 ± Selectively complexes with heavy metals but releases magnesium (+ve error) ± èan be used only when the magnesium release is insignificant

When heavy metal concentrations are high determine hardness by non-EDTA method (by calculation)

@ Suspended and colloidal organic matter can interfere with the end point ± Evaporate sample to dryness, heat in muffle furnace at 550°è, dissolve residue in 20 ml 1N Hèl, neutralize to 7 pH with 1N NaOH and adjust final volume to initial volume with distilled water ± Polluted waters and wastewaters may need pre-treatment through nitric acid ± sulfuric acid, or nitric acid ± perchloric acid digestion

@ Increasing pH can increase sharpness of the end point but higher pH precipitates èaèO3 and Mg(OH)2 ± To minimize the tendency of precipitation of èaèO3 titrate the sample within 5 minutes ± Dilute the sample with distilled water to reduce èaèO3 concentration ± If rough idea about hardness is there add 90% of the titrant to the sample before adjusting its pH to 10±0.1 ± Acidify sample and stir for 2 minutes to expel èO2 prior to pH adjustment to 10 0.1

Titration should be at normal room temperature temperatures make colour change slow

freezing

Select sample size that consumes <15 ml of EDTA solution, makeup volume to 50 ml and add 1 to 2 ml of buffer solution and 1 to 2 drops of indicator solution Titrate with standard EDTA while continuously stirring until the last reddish tinge disappears and end point of blue colour appears Low hardness samples ‡ Take 100 to 1000 ml sample and add proportionately larger amounts of buffer, inhibitor and indicator ‡ Titrate with standard EDTA solution ‡ Run a blank of redistilled, distilled or deionized water of same volume ‡ Subtract volume of the EDTA used for blank from the volume of EDTA used for the sample

V 22 V 2 è è 2 A is mL of EDTA used B is mg èaèO3 equivalent of 1 mL of EDTA