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 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 reaction CO32- + H+ ⇌ HCO3- Bicarbonate neutralization reaction HCO3- + H+ ⇌ H2O.CO2 ⇌ H2O + CO2 Hydroxide neutralization reaction H+ + OH- ⇌ H2O Alkalinity = [OH-] + [HCO3-] + 2[CO32-] – [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 of acidification of waters 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

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 samples Samples may be stored for 24 hrs at 4 °C prior to analysis

Electrochemical Analogy: 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 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

Slows leak but gets contaminated (shorter life-span) 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

red → yellow yellow → blue Colorimetric Indicator Color (acidic → basic) pH Range Malachite green yellow → green 0.2 -1.8 Thymol blue red → yellow yellow → blue 1.2 - 2.8 8.0 - 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 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 10.1 -12.0 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

Measuring Total Alkalinity Remove sample from refrigerator ~30 mins prior to analysis Measure on unfiltered samples 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)

Indicator Methyl Orange end-point ~4.5 Difficult to see More precise indicator is a bromocresol green/methyl red mixture 5.2 – green-blue 5.0 – light blue with lavender grey 4.8 – light pink with blue cast 4.6 light pink From Boehnke, D.N.

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.