Water analysis methods_EB

E.Bergmann 1 Sec/Tert Interface 2009

The Royal Australian Chemical Institute Inc Chemistry Serving Australia

ABN 69 030 287 244

Qld Branch Chemical Education Group

Secondary-Tertiary Interface

22 August 2009

Techniques for water analysis

Prepared by Elaine Bergmann

With thanks to

Rob Gray and Kathy Gray

for their editorial input.


Contents

Dissolved oxygen 3

Hardness 7

Alkalinity 11

Chloride ion 14

Hypochlorite in bleach solutions 16

Phosphate 18


1.0 Dissolved Oxygen (Winkler/iodometric method)

(adapted from Eaton, A, Clesceri, L, Rice, E and Greenberg, A 2005, Standard Methods for the

Examination of Water and Wastewater 21st edition, American Public Health Association, American Water Works Association and Water Environmental Federation, Washington)

1.1 Introduction

Dissolved oxygen (DO) levels in natural and wastewaters depend on the physical, chemical, and biochemical activities in the water body. The analysis of DO is a key test in water pollution and waste treatment process control.

The iodometric method is the most precise and reliable titrimetric procedure for DO analysis. It is based on the sequence of reactions given below.

Interferences

The presence of oxidising or reducing materials in the sample may cause interference with the standard method given here. Certain oxidising agents liberate I2 from iodides (positive interference) and certain reducing agents reduce I2 to iodides (negative interference). Most organic matter is oxidised partially when the oxidised manganese precipitate is acidified, thus causing an error.

A number of modifications are available to minimise interference. The azide modification removes interference due to nitrite (common in effluents and incubated BOD samples). The permanganate modification is used in the presence of Fe2+.

Reaction sequence of iodometry

In the presence of excess OH-, Mn

2+ ions (from MnSO4) form a precipitate of Mn(OH)2. Mn

2+

ions (in equilibrium with Mn(OH)2 precipitate), are oxidised by O2 to Mn3+ forming Mn(OH)3.

O2 + 4 Mn(OH)2 + 2 H2O 4 Mn(OH)3 Addition of acid neutralises OH- and liberates free Mn3+ ions.

Mn(OH)3 + 3H+ Mn

3+ + 3 H2O

I- ions reduce Mn3+ to Mn2+, liberating I2 equivalent to the original DO content. 2 Mn

3+ + 2I

- 2 Mn

2+ + I2

I2 molecules, which are not soluble in water, combine with iodide ions in a reversible reaction to form the brown, soluble tri-iodide ion. I2 + I- I3

-

In the titration reaction, iodine is then reduced by thiosulfate I2 + 2 S2O3

2- 2 I- + S4O62-

As the titration proceeds, the brown colour due to the triiodide ion fades as the iodine is reduced by the thiosulfate. The disappearance of the triiodide is a difficult endpoint to detect, so starch is added as the brown colour fades, forming a blue complex with I2. The endpoint is signalled by the disappearance of the blue colour of the starch-iodine complex.


1.2 Sampling

The reliability of the results depends heavily upon the sampling method. Since dissolved oxygen will be measured back in the laboratory, care must be taken to prevent the loss or addition of oxygen to the sample once it is collected. Because the amount of dissolved O2 is also temperature dependent, the temperature of the sample should also be maintained (or stored at a lower temperature). Chemical and biochemical activity may consume O2. Ideally, Steps 1 - 5 should be performed immediately after sample collection. Once O2 has been consumed, (after formation of I3

-) treated samples can be transported back to the laboratory for titration.

Collect the sample by gently lowering a bottle (about 1 L and preferably wide-necked) into the water and filling the bottle completely beneath the surface. Cap the bottle underneath the water surface so that no bubbles are present in the sample bottle. You need to ensure a smooth flow of water into the bottle here, to avoid aeration. You should endeavour to collect the water sample on the same day that you intend to do the DO assay.

1.3 On-site measurements

Record all information regarding the sample (i.e. place, time, temperature, pH, general description of site). It is most important that you measure the temperature of the water

where you have collected your sample.

1.4 Analysis

1.4.1 Apparatus

1 x 250 mL conical flask with a stopper

1 x 500 mL conical flask

1 x 250 mL measuring cylinder

1 x 10 mL pipette and filler

1 x 50 mL burette

2 x 2mL pipettes

1.4.2 Reagents (The following should be enough to do 4 titrations of a single water sample.)

10 mL 2.2M MnSO4

10 mL alkaline KI reagent (0.9 M KI and 12.5M KOH)

10 mL conc H2SO4

60 mL 0.025M Na2S2O3 (can be standardised with KIO3, see 1.4.5)

10 mL 1% starch suspension

1 L water sample

1.4.3 Safety Latex gloves, safety glasses, remember that the KI reagent and conc. H2SO4 are both extremely corrosive.


1.4.4 Method

1. Carefully fill a 250 mL conical flask (with a stopper) to nearly full with water from your sampling site. Pour the sample water down the side of the flask in order to minimise bubbles. You should leave a space at the top that will accommodate about 5 mL of reagents. (Check with tap water before you start. Ideally you should not have any air bubbles after adding the first 2 reagents and capping the solution.)

2. Using a 2 mL pipette, quickly add 2.0 mL MnSO4 solution, mix and then add 2.0 mL of the alkaline KI reagent.

3. A precipitate of manganese(II) hydroxide will form. Stopper the flask and invert several times to mix and bring about the reaction with the O2(aq). When the precipitate settles, leaving about 100 mL of clear supernatant, repeat the inversions.

4. Allow the precipitate to settle so that about 100 mL of clear solution is produced.

5. Add 2.0 mL of conc H2SO4 (wearing latex gloves) by allowing the acid to run down the neck of the flask. Stopper the flask and invert/mix several times until all the precipitate dissolves.

6. Using a 250 mL measuring cylinder, take 203 mL from the flask and transfer to a 500 mL conical flask. (The 203 mL allows for the addition of the other reagents and is equivalent to 200 mL of the original water sample.) It does not matter that the solution is now open to the air (why not?)

7. Record the initial volume of sodium thiosulfate in the burette. Titrate with this thiosulfate to a pale straw yellow colour. THEN add starch indicator (~0.5 mL). (The starch reacts with iodine to give a deep blue colour.) Continue the titration to the first disappearance of blue colour. Record the final volume of sodium thiosulfate in the burette. Calculate volume of thiosulfate used.

8. Repeat steps 1 – 7 for further aliquots.

1.4.5 Standardisation of thiosulfate solutions

Thiosulfate solutions can be standardised by reaction with iodate. Molarity of iodate solution should be about one-sixth that of the thiosulfate being used.

The overall reaction is as follows:

2 IO3- + 12 H

+ + 12 S2O3

2- 2 I

- + 6 S4O6

2- + 6 H2O

Method

1. Accurately weigh approximately 0.1 g KIO3 (dried at 120˚C for at least 2 hours prior to weighing) and make up to 100 mL in a volumetric flask.

2. Pipette out 25 mL of this solution and add 10 mL of 10% sulfuric acid solution and 2 g of KI.

3. Titrate with thiosulfate solution, adding starch as colour fades to straw yellow.


1.4.6 Calculations and Discussion

1. Calculate the molarity of the thiosulfate from the solution standardisation procedure.

2. Calculate the Dissolved Oxygen (check your notes for some help) for your sample.

Shortcut (if thiosulfate concentration is accurate): When a 200 mL water sample is used, 1mL of sodium thiosulfate solution (0.025 M) is equivalent to 1 mg/L (1ppm) dissolved oxygen in the sample. (If 8.7 mL of sodium thiosulfate was used, then the DO of the sample is 8.7mg/L i.e. 8.7 ppm. 3. Compare your results with the Oxygen Solubility Data (Table 1).

(a) Calculate the percentage saturation of your sample.

(b) Make some suggestions about any differences between your results and the saturation values given in the table.

Table 1. Oxygen Solubility Reference Table

The following values refer to the solubility of oxygen in chloride-free water at various concentrations. The water is in equilibrium with water-saturated air at a total pressure of 1.013 x 103 kPa and a partial pressure of oxygen of 0.213 x 103kPa.

Temp (˚C)

DO (ppm)

Temp (˚C)

DO (ppm)

Temp (˚C)

DO (ppm)

Temp (˚C)

DO (ppm)

0 14.62 11 11.08 22 8.83 33 7.3

1 14.23 12 10.83 23 8.68 34 7.2

2 13.84 13 10.60 24 8.53 35 7.1

3 13.48 14 10.37 25 8.38 36 7.0

4 13.13 15 10.15 26 8.22 37 6.9

5 12.80 16 9.95 27 8.07 38 6.8

6 12.48 17 9.74 28 7.92 39 6.7

7 12.17 18 9.54 29 7.77 40 6.6

8 11.87 19 9.35 30 7.7 41 6.5

9 11.59 20 9.17 31 7.5 42 6.4

10 11.33 21 8.99 32 7.4 43 6.3


2. 0 Water Hardness by Complexation Titration

(Adapted from Greenberg, A.E. et al, 1975 Standard methods for the examination of water and

wastewater 14th edition, American Public Health Association, Washington)

2.1 Introduction

In the past, water hardness was defined as a measure of the capacity of water to precipitate soap. Current laboratory practices, however, define total hardness as the sum of divalent ion concentrations, especially those of calcium and magnesium, expressed in terms of mgCaCO3/L. There are no known adverse health effects of hard or soft water, but the presence of hard waters results in two economic considerations:

(1) hard waters require considerably larger amounts of soap to foam and clean materials and (2) hard waters readily precipitate carbonates (known as boiler scale) in piping systems at high temperatures.

Calcium and magnesium carbonates are two of the few common salts whose solubility decreases with increasing temperature. This is due to the removal of dissolved CO2 as temperature increases. The advent of synthetic detergents has significantly reduced the problems associated with hard water and the “lack of foaming”. However, scale formation continues to be a problem. The method described below relies on a “competition” between EDTA and an indicator (Eriochrome Black T). Both are organic compounds which can form a complex with divalent metal ions. The chemical structure of EDTA is shown below. Note the lone pairs of electrons on the two nitrogens. These, combined with the dissociated carboxyl groups, create a 1:1 hexadentate complex with each divalent ion in solution. However, the complexation constant is a function of pH. Virtually all common divalent ions will be complexed at pH values greater than 10, the pH used in this titration experiment and in most hardness tests. Thus, the value reported for hardness includes all divalent ions in a water sample.

A simplified representation of the neutral form of EDTA. In fact, N atoms are protonated and two of the carbonyl groups are deprotonated, as in the zwitter-ion of amino acids.

The tetra-anion form of EDTA exists at high pH and is able to form 6 bonds to a central metal ion, forming an octahedral complex.


At high pH (>10), each of the 4 carboxyl groups will be deprotonated as will the N atoms (being in a basic solution), so the EDTA will exist in this solution as a tetra-anion.

EDTA is supplied as the disodium salt, (Na+)2H2EDTA

2-, in which 2 protons have been

removed, forming a di-anion.

In solution, the ions dissociate, so H2EDTA2- ions are free in solution. In the presence of excess OH

- ions, the remaining H

+ ions are removed from the carboxyl groups, to form the

tetra-anion. This has lone pairs on the 2 N atoms of the amino groups, as well as on O of each carboxylate anion. These are able to form co-ordinate covalent bonds with empty valence orbitals of metal ions. The EDTA is said to act as a ligand or chelate, and one EDTA forms 6 bonds to each divalent metal ion – hence it’s said to be a hexadentate ligand.

In this analysis you will be titrating Ca2+ in a standard solution and water samples with EDTA. Both Ca2+ solutions and EDTA are colourless so an indicator is needed to signal the reaction completion. The indicator of choice is Eriochrome Black T (EBT) which forms a wine-red complex with Mg2+. A very small amount of Mg2+ will be bound to the indicator through most of the titration. When all of the Ca2+ (and any other divalent metal ions present) has been removed from solution by reaction with EDTA, the Mg

2+ originally bound to the

indicator will combine with the EDTA. The indicator then returns to a sky-blue form signaling the end-point. (Calcium ions are “better” at bonding to EDTA than magnesium, and when there are no free calcium ions left, the EDTA is “better” at bonding with magnesium ions than the indicator is. So at the end point the indicator is not bound to any metal ion. This can be related to the K values (stability constants) for the formation of each complex with EDTA and with EBT given below. Stability constants for the formation of each complex ion are as follows:

[CaEDTA]2- log K = 10.65 K = 4.5 x 1010 [CaEBT]

2+ log K = 5.4 K = 2.5 x 10

5

[MgEDTA]2- log K = 8.85 K = 7.1 x 108 [MgEBT]2+ log K = 7 K = 1.0 x 107

(Aylward, G. and Findlay, T. SI Chemical Data 5th edition (pp. 132 - 133), John Wiley and Sons, Australia Limited)


2.2 Analysis

2.2.1 Apparatus

20 mL pipette and filler 50 mL burette 250 mL conical flask pH meter or indicator strips to give pH 10 – 10.5 50 mL pipette (or measuring cylinder)

2.2.2 Reagents

250 mL 0.01M CaCl2 solution (see Section 2.2.3.1)

250 mL 0.01M EDTA solution (Weigh approximately 0.9 g of AR grade disodium ethylenediamine tetraacetate dihydrate (Na2H2C10H12O8N2.2H2O), dissolve in distilled water, and dilute to 250 mL. Store in a labelled plastic bottle.)

pH 10 buffer containing Mg2+ (Dissolve 16.9 g NH4Cl in 143 mL conc NH4OH, to which has been added 1.25 g of the Mg salt of EDTA. (If Mg salt of EDTA is not available, dissolve 1.18g of MgSO4.7H2O or 645mg of MgCl2.6H2O in 50 mL H2O and add to buffer solution. Make up to 250 mL with distilled water. Store in a plastic bottle, tightly stoppered.)

Eriochrome Black T indicator (For a powder: mix 0.5 g of the dye and 100 g NaCl For a solution: mix 0.5 g of the dye with 4.5 g of hydroxylamine hydrochloride, then dissolve in 100 mL of 95% ethyl or isopropyl alcohol.)

2.2.3 Standardisation of EDTA Solution

2.2.3.1 Preparation of a standard solution of 0.01M CaCl2

1. Accurately weigh approximately 0.2500 g of dried anhydrous CaCO3, record its mass and quantitatively transfer to a 250 mL Erlenmeyer flask.

2. Place a funnel in the neck of the flask and add approximately 25 mL of distilled water, then 2 mL of 6M HCl. Mix, and if the mixture remains cloudy, add another drop of HCl. Continue adding HCl dropwise until the solution is clear.

3. Add 50 mL distilled water and boil for a few minutes to expel CO2.

4. Cool, add a few drops of methyl red indicator, and adjust to the intermediate orange colour (approx pH 5) by adding 3M NH4OH or 6M HCl, as required.

5. Transfer quantitatively to a 250 mL volumetric flask and fill to the mark with distilled water. Stopper, invert 10 – 15 times to mix thoroughly.

If 0.2500 g of CaCO3 is used, this standard is equivalent to 1.00 mg CaCO3/1.00 mL. (1000 ppm) If a slightly different amount of CaCO3 is used, then a factor must be applied e.g. 0.2432/0.2500)


2.2.3.2 Standardisation of 0.01M EDTA solution

1. Rinse a 50 mL burette with distilled water then with the EDTA solution. Fill with EDTA.

2. Rinse a pipette with distilled water then with CaCl2. Pipette 20 mL of the CaCl2 solution into a rinsed 250 mL Erlenmeyer flask.

3. Add pH 10 buffer solution to give a pH of 10 – 10.1, and 2 drops of Eriochrome Black T indicator. Mix well. What colour should the solution be at this point?

4. Titrate the sample in the flask with the EDTA, slowing additions (as red becomes more purple) as you near the sky-blue endpoint (wait 3–5 seconds between additions, as the blue colour does not develop immediately). Record the volume required to titrate the CaCl2

sample. Keep the first sample as a colour reference. Your titration should not take longer than 5 minutes, from the time the buffer is added to endpoint.

5. Prepare a second portion of CaCl2 in the same way as above (steps 2 - 3) and repeat the titration, recording all data. If the volumes of EDTA agree to within 0.4 mL, proceed to the next section. If not, repeat the titration until 2 or more titres agree within 0.4 mL.

2.2.4 Determination of water hardness

1. Measure100 mL of tap water into a 250mL flask. Add pH 10 buffer (about 5 times the volume you added to the standard, to bring pH to 10 – 10.1 as above) and 5 drops of Eriochrome Black T. Mix well. Titrate with the standardised EDTA solution. (100 mL of SEQ water should give a titre of around 10 mL; the volume of water can be varied for water with more or less calcium. If your water is very hard, then reduce the volume of water in the flask.)

2. Repeat until you have 2 titres which are in agreement as Step 5 above.

2.2.5 Calculations

1. Determine the moles of CaCO3 (moles Ca2+) in the 250 mL of standard solution, and in each 20.00mL aliquot. 2. EDTA and Ca2+ react in a 1:1 ratio. Calculate the average volume of EDTA which reacted with the standard solution, and hence calculate EDTA concentration. 3. Calculate the average volume of EDTA which reacted completely with the tap water sample, and hence calculate the concentration of calcium (and other divalent ions) in the water. Express this result both as a molarity and as mg CaCO3 per litre (ppm). 4. The Brisbane City Council website states that the average total hardness of Brisbane tap water is 100mg as sCaCO3/L. How closely does your result agree with this value? If it does not, suggest possible reasons for the difference. Consider both your laboratory procedures and other factors.


3.0 Alkalinity



3.1.1 Introduction

This procedure could be taught in conjunction with work on equilibrium and buffers with specific reference to the carbonate/bicarbonate equilibrium.

Alkalinity is a measure of the acid neutralising capacity of a water or wastewater and is a useful component in the interpretation of water quality for various uses.

As neutralising capacity depends mainly on carbonate, bicarbonate and hydroxide, it can be used as an indicator of the concentrations of these components.

For groundwaters, the main source of carbonate and bicarbonate is rocks through which the water has passed.

3.1.2 Summary of Test Method

The total alkalinity is determined by titrating a portion of sample with standardised sulfuric acid to pH 4.5 end point using bromocresol green-methyl red indicator solution, or for samples with pH greater than 8.3, “Phenolphthalein Alkalinity” to pH 8.3 end point using phenolphthalein indicator.

3.1.3 Interferences

Strongly coloured or turbid samples may hinder end point detection. To overcome this, a smaller aliquot of sample may be used or the procedure for low alkalinity may be followed. Sample alkalinity less than 20 mg CaCO3/L can only be reported if the low alkalinity procedure is used.

3.1.4 Sample preservation

Samples should be collected in glass or plastic bottles with no entrapped air if possible. If stored at 0-4

˚C immediately upon receipt, samples can be kept for up to 24 hours prior to

analysis. Do not filter samples. Samples may be subject to microbial action and to loss or gain of gases when exposed to air, so they should be analysed without delay. Avoid sample agitation and prolonged exposure to air.

3.2 Analysis

3.2.1 Apparatus

150 or 250 mL Erlenmeyer Flasks 50 mL burette 10 mL pipette. 50 mL pipette For low alkalinity procedure. 20 mL microburette pH meter 250 mL beaker magnetic stirrer


3.2.2 Reagents

• 0.01M sulfuric acid (see 3.2.3.1 for standardisation procedure)

• mixed bromocresol green - methyl red indicator solution (Dissolve 100 mg bromocresol green and 20 mg methyl red into 100 mL 95% ethyl or isopropyl alcohol; or

100mg bromocresol green sodium salt and 20mg methyl red sodium salt in 100 mL water.)

• phenolphthalein solution (Dissolve 1 g phenolphthalein in 100 mL 95% ethyl or isopropyl alcohol and add 100 mL distilled water.)

Note: The colour response of the mixed bromocresol green-methyl red is approximately as follows: above pH 5.2, greenish blue; pH 5.0, light blue with lavender grey; pH 4.8, light pink grey with bluish cast; pH 4.5, light pink; and below pH 4.0, strong pink.

3.2.3 Method

3.2.3.1 Standardisation of 0.01M H2SO4

Make up a 0.025M Na2CO3 solution.

1. Dry 3 to 5 g of primary standard Na2CO3 at 250C for 4 hours and cool in a desiccator.

2. Weigh 2.5 +/- 0.2g (to the nearest mg), transfer to a 1L volumetric flask, fill flask to the mark with distilled water, and dissolve and mix reagent. Do not keep longer than a week.

3. Titrate 10 mL of this solution with the 0.01M H2SO4 using 5 drops of the bromocresol green–methyl red indicator.

4. Repeat the titration as necessary.

5. An indicator blank should be titrated as well, to ensure that this result is not being affected by residual alkalinity of the distilled water used to make up the standard solution. Add 5 drops of indicator to 10mL of the water, and titrate as above. If this value is significant, then the volume should be subtracted from the volume used in the standardization titration before determining concentration of sulfuric acid.

Calculations

1. Write the balanced equation for the reaction between Na2CO3 and H2SO4.

2. Find the number of moles of Na2CO3 (nNa2CO3) weighed out to make up the standard solution.

3. Calculate the molarity of the Na2CO3 solution (M Na2CO3).

4. Subtract the value for the indicator blank from the average volume of acid that was added to the Na2CO3 solution. (VH2SO4)

5. Use this value to calculate the molarity of the H2SO4 solution.


3.2.3.2 Total alkalinity

•Test the pH of the sample before beginning titration. If it is above 8.3, then you must use the phenolphthalein procedure. If below 8.3, then proceed as follows:

1. Pipette 50 mL of sample (or less if alkalinity is high) into an Erlenmeyer flask.

2. Add approximately 5 drops of mixed indicator to the sample.

3. Over a white surface, titrate with 0.01M H2SO4 to a persistent colour change characteristic of the equivalence point (this can be found by making a pH buffer solution to pH 4.5 and adding 5 drops of mixed indicator).

4. Repeat steps 1 – 3 as required.

3.2.3.3 Phenolphthalein alkalinity

If solution pH is greater than 8.3 proceed as in Total Alkalinity above, substituting phenolphthalein for the mixed indicator.

3.2.3.4 Low alkalinity samples

Use a 200 mL sample volume, and titrate using a 20 mL microburette and the pH 4.5 endpoint procedure.

3.2.4 Calculations

As alkalinity is reported as the equivalent mass of CaCO3, calculations are based on the reaction between CaCO3 and H2SO4.

As the stoichiometric ratio is 1:1, at the end point, n(CaCO3) = n(H2SO4).

M(CaCO3)= M(H2SO4) x V(H2SO4)/V(water sample)

Alkalinity (as mg/L CaCO3) = M(CaCO3) x Molar Mass (CaCO3) x 103mg/g

(n = no. of moles; M = molarity; V = volume of solution or sample)


4. 0 Chloride ion titration by Mohr (argentometric) method



4.1 Introduction

The method is suitable for the analysis of chloride ion concentrations in relatively clean water when 0.15 to 10 mg of Cl- is present in the portion titrated. The titrant is a standardised solution of AgNO3. Chloride ions are quantitatively precipitated as AgCl before the red silver chromate precipitate is formed. Ag

+(aq) + Cl

- (aq) AgCl(s) (Ksp = 1.8 x 10

-10; s = 1.34 x 10

-5 M)

2Ag+(aq) + CrO42- (aq) Ag2CrO4(s) (Ksp = 2.6 x 10-12; s = 8.66 x 10-5 M)

4.2 Sample collection

Collect samples in clean chemically resistant glass or plastic bottles. Maximum sample portion required for a single titration is 100 mL.

4.3 Analysis

4.3.1 Apparatus

250 mL flask 50 mL burette pH probe or indicator strips for pH range 7 - 10 1 mL pipette 100 mL pipette or measuring cylinder

4.3.2 Reagents

K2CrO4 indicator solution (Dissolve 50 g K2CrO4 in a little distilled water. Add AgNO3 solution until a definite red precipitate is formed. Let stand for 12 hours, filter (Whatman No. 42) and dilute to 1 L with distilled water.)

standard silver nitrate titrant, 0.0141 M (Dissolve 2.395 g AgNO3 in distilled water and dilute to 1000 mL. Standardise against NaCl by the procedure described below. Store in a brown bottle.)

standard NaCl solution, 0.0141 M. (Dissolve 824.0 mg NaCl (dried at 140˚C) in distilled water and dilute to 1000 mL

0.5 M H2SO4

1.0M NaOH


4.3.3 Method

4.3.3.1 Standardisation of AgNO3 titrant

1. Establish a reagent blank value by titrating 100mL of distilled water plus 1.0 mL of K2CrO4 indicator with the AgNO3 titrant.

2. Titrate 20 mL of the standard NaCl solution with the AgNO3 to a pinkish yellow end point. Be consistent in end point recognition.

3. Subtract the reagent blank value from the titre and use the difference to calculate the concentration of the AgNO3.

4.3.3.2 Sample preparation

1. Use a 100 mL sample or a suitable portion diluted to 100mL if chloride ion concentration is high.

2. Test pH of sample. Adjust sample pH to 7 - 10 with 0.5M H2SO4 or 1M NaOH as necessary.

4.3.3.3 Titration

1. Add1.0 mL K2CrO4 indicator solution.

2. Titrate with standard AgNO3 solution to a pinkish yellow end point as for the standardisation above.

Chloride ion concentration of sample is reported as mg Cl- per L (ppm).


5.0 Sodium hypochlorite (free chlorine) levels in Bleach

5.1 Introduction

Note: Standard titrimetric methods for determination of free chlorine (hypochlorite ion, hypochlorous acid and chlorine) are not sensitive enough for the analysis of municipal water supplies. This method is suitable for the analysis of pool water, or to investigate the effect of factors such as UV light or addition of organic materials to pool water.

A redox titration is used to determine the concentration of free chlorine in a sample of bleach. A diluted sample of commercial bleach is firstly reacted in a flask with a solution containing an excess of acidified iodide ion. This reaction produces I2. (This in turn combines with iodide ion to form the very visible soluble brown triiodide ion I3

-.)

ClO- + 2I- + 2 H+ I2 + Cl- + H2O (1) (faint yellow) (colourless) (brown)

We can’t detect an end-point for the reaction above, so the I2 product is titrated with thiosulfate. As the iodine is used, the brown colour due to the triiodide ion fades to yellow, and then disappears. This is a hard endpoint to detect, so when the brown starts to fade to yellow, starch can be added. Starch forms a blue complex with I2, and this colour disappears when the stoichiometric amount of thiosulfate has been added to the flask. 2 S2O3

2- + I2 2 I- + S4O62- (2)

From equations (1) and (2), it can be seen that 1 mole of ClO- reacts to form 1 mole I2, which consumes 2 moles of S2O3

2-. Hence, 1 mole of hypochlorite is equivalent to 2 moles of

thiosulfate. Typical bleach concentrations are quoted as 35 g/L available Cl as sodium hypochlorite (approx 1M) or approximately 5.25% NaClO by mass (approx 0.7M), so it is advisable to dilute the bleach by a factor of 1 in 10.

5.2 Analysis

5.2.1 Apparatus

pipettes 10 mL, 20 mL and 25 mL and filler burette (50 mL) 100 mL volumetric flask 250 mL conical flasks

5.2.2 Reagents

bleach (freshly purchased)

0.2 M Na2S2O3 solution

acidified NaI or KI solution (12.6 g NaI or 14.0 g KI + 40 mL glacial acetic acid in 1 L)

1% starch suspension


5.2.3 Method

5.2.3.1 Part A standardisation of thiosulfate

Thiosulfate solutions can be standardised by reaction with iodate. Molarity of iodate solution should be about one-sixth that of the thiosulfate being used.

The overall reaction is as follows:

2 IO3- + 12 H

+ + 12 S2O3

2- 2 I

- + 6 S4O6

2- + 6 H2O

1. Accurately weigh approximately 0.7 g KIO3 (dried at 120˚C for at least 2 hours prior to weighing) and make up to 100 mL in a volumetric flask.

2. Pipette out 25 mL of this solution and add 10 mL of 10% sulfuric acid solution and 2 g of KI.

3. Titrate with thiosulfate solution, adding starch as colour fades to straw yellow.

5.2.3.2 Part B titration of bleach solution

1. Prepare the diluted bleach solution by pipetting 10 mL into a 100 mL volumetric flask and filling to the mark with distilled water. Stopper and invert several times to mix.

2. Rinse a 20 mL (or 25 mL) pipette with the bleach solution, and transfer an aliquot to a conical flask.

3. Add 50 mL of the acidified iodide solution to the bleach in the flask. It should go brown. If it does not, check the bleach – it decomposes over time.

4. Rinse and fill the burette with Na2S2O3 solution, and titrate until the brown colour fades to a straw yellow. Add a few drops of starch indicator. A blue starch-iodine complex should form.

5. Continue adding Na2S2O3 to the flask until the blue colour disappears.

6. Repeat titrations.

5.2.3.3 Calculations

1. Use the stoichiometric information given above to find the molarity of your diluted bleach solution. 2. Calculate the molarity of the neat bleach.


6.0 Colorimetric Phosphate Determination (stannous chloride method)



6.1 Introduction

Phosphorus occurs in natural waters and in wastewaters almost solely as phosphates. These

are classified as orthophosphates (phosphate PO43-

), condensed phosphates (polyphosphates)

and organically bound phosphates. They occur in solution, in particles or detritus, or in the

bodies of aquatic organisms.

These forms of phosphate arise from a variety of sources. Small amounts of orthophosphate

or certain condensed phosphates are added to some water supplies during treatment. Larger

quantities of the same compounds may be added during laundering or other cleaning, because

these materials are major constituents of many commercial cleaning preparations.

Orthophosphates applied to agricultural or residential cultivated land as fertilisers are carried

into surface waters with storm runoff. Organic phosphates are formed primarily by biological

processes. They are contributed to sewage by body wastes and food residues.

Phosphorus is essential to the growth of organisms and can be the nutrient that limits the

primary productivity of a body of water. In instances where phosphate is a growth-limiting

nutrient, the discharge of raw or treated wastewater, agricultural drainage, or certain industrial

wastes to that water may stimulate the growth of photosynthetic aquatic micro- and macro-

organisms in nuisance quantities. Phosphates also occur in bottom sediments and in biological

sludges, both as precipitated inorganic forms and incorporated into organic compounds.

Analysis of phosphorus involves 2 steps: (i) digestion or conversion of the phosphorus to

dissolved orthophosphate; (ii) colorimetric determination of dissolved orthophosphate. The

digestion can be accomplished in 3 ways: (i) perchloric acid recommended only for difficult

samples such as sediments; (ii) nitric acid-sulfuric acid can be used for most samples, but the

simplest is (iii) persulfate oxidation. This method oxidises all nitrogen compounds to nitrate

and all phosphorus compounds to orthophosphate.

In this colorimetric determination of phosphate concentration, ammonium molybdate reacts

under acid conditions to form a heteropoly acid, molybdophosphoric acid. This is then

reduced by stannous chloride to intensely coloured moleybdenum blue. The concentration of

the orthophosphate is then determined colorimetrically.


6.2 Analysis

6.2.1 Apparatus

hotplate

10 mL pipette

25 mL pipette

50 mL conical flask x 4

100 mL volumetric flask

25 mL volumetric flask x 6

80 mL conical flask x 10

spectrophotometer

6.2.2 Reagents

6.2.2.1 Part A persulfate digestion

• phenolphthalein indicator solution

• strong acid solution

(slowly add 300 mL conc H2SO4 to 600 mL distilled water. When cool, add 4.0 mL conc HNO3 and dilute to 1L.)

• sulfuric acid solution

(Carefully add 300mL conc H2SO4 solution to approx 600 mL distilled water and dilute to 1L with distilled water.)

• potassium persulfate, K2S2O8, solid or ammonium persulfate, (NH4)2S2O8 , solid

• 1.0M sodium hydroxide

6.2.2.2 Part B Colour development

• phenolphthalein indicator solution

• strong acid solution (as above)

• ammonium molybdate reagent

(Dissolve 25 g of (NH4)6Mo7O24.4H2O in 175 mL of distilled water. Cautiously add 280 mL conc. H2SO4 to 400 mL distilled water. Cool, add molybdate solution, and dilute to 1 L.)

• stannous chloride reagent

(Dissolve 2.5 g fresh SnCl2.2H2O in 100 mL of glycerol. Heat in a water bath and stir with a glass rod to hasten dissolution. This reagent is stable and requires neither preservative nor special storage.)

6.2.2.3 Part C Standard curve

• standard phosphate solution

(Dissolve in distilled water 219.5 mg anhydrous KH2PO4 and dilute to 1 L (1.00 mL = 50.0 µg phosphorus; or P concentration is 50.0 ppm.)


6.2.3Method

Acid washing of all glassware

All glassware should be rinsed with hot dilute (5%) HCl solution to remove traces of

phosphates left by detergents. Rinse with distilled water.

6.2.3.1 Part A - Acid persulfate digestion (for each water sample and for the calibration

curve standards)

1. Use 50 mL or a suitable portion of a thoroughly mixed sample. Add 0.05 mL (1 drop) of

phenolphthalein indicator. If sample turns pink, add strong acid solution dropwise to just

discharge the colour. Then add 1 mL H2SO4 solution and either 0.4 g solid (NH4)2S2O8 or

0.5g of solid K2S2O8.

2. Boil gently on a preheated hotplate for 30 to 40 minutes or until a final volume of 10 mL is

reached. Cool, dilute to 30 mL with distilled water, add 0.05 mL (1 drop) of phenolphthalein

indicator solution, and neutralise to a faint pink colour with NaOH.

3. Make up to 100 mL with distilled water. In some samples, a precipitate may form at this

stage, but do not filter. For any subsequent subdividing of the sample, shake well. The

precipitate (which is possibly a calcium phosphate) redissolves under the acidic conditions of

the colorimetric reactive phosphorus test.

6.2.3.2 Part B Colorimetric development

• Preliminary sample treatment To 100 mL of digested sample containing not more than

200µg of phosphate and free from colour and turbidity, add 0.05 mL (1 drop) of

phenolphthalein indicator. If sample turns pink, add strong acid solution dropwise to

discharge the colour. If more than 0.25 mL (5 drops) is required, take a smaller sample and

dilute to 100 mL with distilled water after first discharging the pink colour with acid.

• Colour development Add, with thorough mixing after each addition, 4.0 mL molybdate

reagent and 0.5 mL (10 drops) of stannous chloride reagent. Rate of colour development and

intensity of colour depend on temperature of the final solution, each 1˚C increase producing

about 1% increase in colour. Hence, hold samples, standards and reagents within 2˚C of one

another and the temperature range between 20 and 30˚C.

• Colour measurement After 10 minutes, but before 12 minutes, using the same specific

interval for all determinations, measure colour photometrically at 690 nm, using a distilled

water blank. Use a calibration curve prepared as described in Part C below to find the

concentration of P.

Always run a blank on reagents and distilled water. Because the colour at first develops

progressively and later fades, maintain equal timing conditions for samples and standards.

Prepare at least one standard with each set of samples or once each day that tests are made.

The calibration curve may deviate from a straight line at the upper concentrations of the 0.3 to

2.0 mg/L range.


6.2.3.3 Part C Preparation of calibration curve

• Prepare a 5 ppm P standard solution by pipetting 10.0 mL of the stock (50ppm P) solution

into a 100 mL volumetric flask and making up to 100 mL.

• From this, prepare a range of standards from 0 - 5 ppm according to the volumes given

below.

Concentration of P

(ppm, mg/L or µµµµg/mL)

Volume of 5ppm

solution

(mL)

Volume of water

(mL)

5 25 0

4 20 5

3 15 10

2 10 15

1 5 20

0 0 25

• To prepare each standard solution for colorimetric measurement, apply the persulfate

digestion and colorimetric development procedures as given above.

• Prepare the calibration curve.

Water analysis methods_EB

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