Field Methods of Monitoring Aquatic Systems Unit 5 – pH, Acidity and Alkalinity Copyright © 2008...
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Transcript of Field Methods of Monitoring Aquatic Systems Unit 5 – pH, Acidity and Alkalinity Copyright © 2008...
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
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.