Transcript Field Methods of Monitoring Atmospheric Systems

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
Or
pH = -log10 [H3O+]
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
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
CO2
NO2
SO2
Natural
Anth.
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
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Measure of the ability of a water body
to neutralize acidity
Dissolution of limestone and other
minerals produces alkalinity
e.g.
CaCO3 ⇌ Ca2+ + CO32CO32- + H2O ⇌ HCO3- + OHWater 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
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Measurement of the buffer capacity (resistance to pH change)
e.g.
Carbonate neutralization reaction
CO32- + H+ ⇌ HCO3Bicarbonate neutralization reaction
HCO3- + H+ ⇌ H2O.CO2 ⇌ H2O + CO2
Hydroxide neutralization reaction
H+ + OH- ⇌ H2O
Alkalinity = [OH-] + [HCO3-] + 2[CO32-] – [H+]
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Units are mg L-1 CaCO3 or mEq L-1 (regardless of species)
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Acid titration to change the pH to 4.5 (methyl orange end-point)
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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
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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
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
Indicator
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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
pH Range
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
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
Yellow → red
10.1 -12.0
Malachite green
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Color
(acidic → basic)
Alizarin yellow
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
Net ionic:
CaCO3 + H2SO4 → H2CO3 + CaSO4
CO32- + 2H+ → H2CO3
• Range 30 - 500 mg CaCO3 L-1
– Rainwater < 10
– Surface water < 200
– Groundwater > 1000 (due to MO decomposition)
Indicator
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Methyl Orange end-point ~4.5
Difficult to see
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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
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
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Units are mg L-1 CaCO3 or mEq L-1 (regardless of species)
mEq L-1 = mg L-1 CaCO3 divided by 50
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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
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Hach Titrator
– Cartridge based system
– 100 mL cylinder
– 250 mL beaker
Source: http://www.hach.com
Text Books
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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.