Field Methods of Monitoring Aquatic Systems

Unit 5 – pH, Acidity and Alkalinity

Copyright © 2008 by DBS

Title

Pure water is neither acidic or basic because it contains equal concentrations of hydroxide and hydronium ions

Role of pH in Water Quality

Brønsted-Lowry definition• Acid is a proton donor

HCl + H2O → H3O+ + Cl-

• Base is a proton acceptor

NH3 + H2O → NH4+ + OH-

Acidic: H+ > OH- Basic: OH- > H+

pH Scale

pH = -log10 [H+]

[H+] = 10-pH

Typically 0 – 14 (can go beyond this)

[H+] = [OH-] = 1.0 x 10-7 moles L-1 (pH = 7, neutral)

For each change of one pH unit [H+] changes x10

Or pH = -log10 [H3O+]

pH of Common Substances

Substance pH

Battery acid 0.3

Lemon juice 2.4

Urine 4.8 - 7.5

Rainwater 5.5 - 6.0

Blood 7.35 - 7.45

Bleach 10.5

Ammonia 11.5

Typical pH Values

Reeve, 2002

Rainwater

• Unpolluted rain water is slightly acidic due to dissolved CO2

(NO2 and SO2), pH ~ 5.6

H2O(l) + CO2(g) ⇌ H2CO3(aq)

⇌ H+(aq) + HCO3

-(aq) ⇌ 2H+

(aq) + CO32-

(aq)

Gas Natural Anth.

CO2

NO2

SO2

PA Acid Deposition

Aerochem Metrics wet/dry precipitation collector

http://www.dep.state.pa.us/dep/deputate/airwaste/aq/acidrain/acidrain.htm

Question

We must always hold an objective view. If you look for it there is a positive side of the existence of acid rain. What could this be?

Acid rain cleans the atmosphere of pollutants

Alkalinity

• Measure of the ability of a water body to neutralize acidity

• Dissolution of limestone and other minerals produces alkalinity

e.g.

CaCO3 ⇌ Ca2+ + CO32-

CO32- + H2O ⇌ HCO3

- + OH-

• Water supply with high total alkalinity is resistant to pH change

• Two samples with identical pH but different alkalinity behave differently on addition of acid

– Different capacity to neutralize acid

Mineral Composition

Calcite CaCO3

Magnesite MgCO3

Dolomite CaCO3.MgCO3

Brucite Mg(OH)2

Alkalinity

• Measurement of the buffer capacity (resistance to pH change)

e.g. Carbonate neutralization reactionCO3

2- + H+ ⇌ HCO3-

Bicarbonate neutralization reactionHCO3

- + H+ ⇌ H2O.CO2 ⇌ H2O + CO2

Hydroxide neutralization reactionH+ + OH- ⇌ H2O

Alkalinity = [OH-] + [HCO3-] + 2[CO3

2-] – [H+]

• Units are mg L-1 CaCO3 or mEq L-1 (regardless of species)

• Acid titration to change the pH to 4.5 (methyl orange end-point)

• If pH < 4.5 there is no acid neutralizing capacity i.e. no need to measure alkalinity

Biological and Chemical Effects

• Sensitivity of fish populations– Salmon populations decrease below 6.5– Perch below 6.0– Eels below 5.5

• Increases solubility of metals– Toxic Al3+ and Pb 2+ release– Particuarly from soils (aluminosilicates)

• Increases weathering of minerals and crustaceans

of acidification of waters

Water Quality

• Public Health Service Act accepted level 6.5-8.5• Public health concern is corrosion and leaching of toxic metals

(Pb, Cu, Zn, Fe) from metal pipes

Measuring pH

• Electrochemical• Colorimetric

Remove sample from refrigerator ~30 mins prior to analysis Measure on unfiltered samplesSamples may be stored for 24 hrs at 4 °C prior to analysis

Electrochemical

• Electrodes generate a voltage directly proportional to the pH of the solution

– pH 7 potential is 0 V– < 7 +ve V, > 7 –ve V

Analogy:

Battery where +ve is measuring electrode, -ve is reference electrode

• Electrochemical potential - known pH liquid inside the glass H+ sensitive membrane electrode vs. unknown outside

• Circuit is closed through the solutions - internal and external - and the pH meter

Flowing• Internal KCl slowly flows to

the outside through the junction (salt bridge)

• Must be refilled!

Gelled• Slows leak but gets

contaminated (shorter life-span)

Source: http://www.ph-meter.info

Thin Glass Membrane

• Aluminosilicate (Al2SiO5)

• Kegley description is incorrect, not controlled by H+ but Na+

Electrochemical Potential

Nernst equation

• Ecell = constant – 0.059 pH (at 25 °C)

• Calibrated with buffer solutions of known pH

• Straight line plot of Ecell vs. pH

Colorimetric

• Acid-base indicator solution or indicator paper

• Indicators are large organic molecules that change color depending on pH

e.g, cresol red is yellow < 7.0 and red > 8.8 and various shades in between

Indicator Color(acidic → basic)

pH Range

Malachite green yellow → green 0.2 -1.8

Thymol blue red → yellowyellow → blue

1.2 - 2.88.0 - 9.6

Methyl orange red → yellow 3.2 – 4.4

Bromocresol green

Yellow → blue 3.8 -5.4

Methyl red Red → yellow 4.8- 6.0

Bromothymol blue Yellow → blue 6.0 - 7.6

Cresol red Yellow → red 7.0 - 8.8

Phenolphthalein Colorless → pink 8.2 - 10.0

Thymolphthalein Colorless → blue 9.4 - 10.6

Alizarin yellow Yellow → red 10.1 -12.0

Measuring Total Alkalinity

• To unfiltered sample add strong acid of known concentration, (0.0100 M H2SO4) titrate to pH 4.5

CaCO3 + H2SO4 → H2CO3 + CaSO4

Net ionic: CO32- + 2H+ → H2CO3

• Range 30 - 500 mg CaCO3 L-1

– Rainwater < 10– Surface water < 200– Groundwater > 1000 (due to MO decomposition)

Remove sample from refrigerator ~30 mins prior to analysis Measure on unfiltered samples

Indicator

• Methyl Orange end-point ~4.5• Difficult to see

• More precise indicator is a bromocresol green/methyl red mixture5.2 – green-blue5.0 – light blue with lavender grey4.8 – light pink with blue cast4.6 light pink

Question

What is the total alkalinity for a sample requiring 21.25 mL of 0.0100 M H2SO4?

0.02125 L x 0.0100 mol L-1 = 2.125 x 10-4 mol H2SO4

Mole ratio is 1:1

2.125 x 10-4 moles H2SO4 = 2.125 x 10-4 moles CaCO3

2.125 x 10-4 mol CaCO3 x 100.09 g / mol = 2.13 x 10-2 g = 21.3 mg

21.3 mg CaCO3 = 213 mg CaCO3 L-1

0.100 L

Units

• Units are mg L-1 CaCO3 or mEq L-1 (regardless of species)mEq L-1 = mg L-1 CaCO3 divided by 50

• CaCO3 + 2H+ ⇌ H2CO3

mg x 1 mmol x 2mEq = mEq L 100 mg mmol L

mg x 1/50 = mEq L L

Field Method / High-Throughput Labs

• Hach Titrator

– Cartridge based system

– 100 mL cylinder

– 250 mL beaker

Source: http://www.hach.com

Text Books

• Rump, H.H. (2000) Laboratory Manual for the Examination of Water, Waste Water and Soil. Wiley-VCH.

• Nollet, L.M. and Nollet, M.L. (2000) Handbook of Water Analysis. Marcel Dekker.

• Keith, L.H. and Keith, K.H. (1996) Compilation of Epa's Sampling and Analysis Methods. CRC Press.

• Van der Leeden, F., Troise, F.L., and Todd, D.K. (1991) The Water Encyclopedia. Lewis Publishers.

• Kegley, S.E. and Andrews, J. (1998) The Chemistry of Water. University Science Books.

• Narayanan, P. (2003) Analysis of environmental pollutants : principles and quantitative methods. Taylor & Francis.

• Reeve, R.N. (2002) Introduction to environmental analysis. Wiley.

• Clesceri, L.S., Greenberg, A.E., and Eaton, A.D., eds. (1998) Standard Methods for the Examination of Water and Wastewater, 20th Edition. Published by American Public Health Association, American Water Works Association and Water Environment Federation.