Chem/Envr 251
Spring 2005
Chapter 9: The Chemistry of Natural Waters
Planetary Water Resources
Major Stocks of Water on Earth
(From P. H. Glieck, The World's Water 2001-02, TD345.W674 2000)
Volume
Percentage of
Percentage of
(1000 km3)
Total Water
Total Fresh Water
Salt water stocks
Oceans
1,338,000
96.54
Saline/brackish groundwater
12,870
0.93
Saltwater lakes
85
0.006
Freshwater stocks
Glaciers, permanent snowcover
20,064
1.74
68.70
Fresh groundwater
10,530
0.76
30.06
Ground ice, permafrost
300
0.022
0.86
Freshwater lakes
91
0.007
0.26
Soil moisture
16.5
0.001
0.05
Atmospheric water vapor
12.9
0.001
0.04
Marshes, wetlands
11.5
0.001
0.03
Rivers
2.12
0.0002
0.006
Incorporated in biota
1.12
0.0001
0.003
1,386,000
100
Total Water
35,029
2.53
100
Total Freshwater
Oceans
Continent
Europe
Asia
Africa
North America
South America
Australia/Oceania
Antarctica
Total Land
Water Flows
(Glieck, 2000)
Precipitation
Evaporation
(km3/yr)
(km3/yr)
458,000
505,000
8,290
32,200
22,300
18,300
28,400
7,080
2,310
118,880
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5,320
18,100
17,700
10,100
16,200
4,570
0
71,990
Runoff
(km3/yr)
2,970
14,100
4,600
8,180
12,200
2,510
2,310
46,870
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Hydrologic Cycle
Composition of Natural Waters
Seawater
-Seawater contains about 3.5% dissolved solids and ions.
3.5% (by mass) = 35,000 ppm = 35 g dissolved solids/L water
-Result of ions carried in from runoff which remain when seawater evaporates.
Major Constituents of Seawater (Girard, Prin. Env. Chem., 2005)
Ion
Conc. (ppm)
Ion
Conc. (ppm)
+
Chloride, Cl
19,000
Potassium, K
380
Sodium, Na+
10,600
Bicarbonate, HCO3140
2Sulfate, SO4
2,600
Bromide, Br
65
Magnesium, Mg+2
1,300
Others
34
Calcium, Ca+2
400
Total
34,519
Freshwater
-Freshwater generally contains less than 0.1% (1000 ppm) dissolved solids and ions.
-Composition is variable, dependant on climate and composition of rocks and soils.
-Common cations include: Ca+2, Mg+2, Na+, K+, and possibly NH4+
-All but Na+ derive from the weathering of rock, soils, and minerals in the ground.
-Na+ is present in the atmosphere from salt water spray and is easily rained out.
-NH4+ is from atmospheric NH3.
-Common anions include: Cl-, SO4-2, HCO3-, and possibly NO3-Derive from minerals, but sulfate, nitrate, and especially bicarbonate come from
atmospheric gases and chloride from sea spray.
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Some Unique Properties of Water
-hydrogen bonding in water:
-raises the expect b.p. to from ~-80°C to +100°C, therefore a liquid at "normal" temps
-in ice, forces the lower density packing that allows ice to float on water.
-polar--dissolves ionic and polar species
-high heat capacity--can absorb a large quantity of heat with minimal temperature changes
-temperature moderation due to oceans and large lakes.
-high heat of vaporization--requires a large amount of energy to evaporate
-evaporative cooling
-I'd call this evidence of design in the universe.
Redox Chemistry in Natural Waters
Dissolved Oxygen
-All surface water is in contact with atmospheric oxygen, O2.
-There is an equilibrium between gaseous oxygen and dissolved oxygen:
O2(g)  O2(aq)
-Based on Henry's law:
[O ]
K H  2  0.0013
PO2
-Solubility of oxygen decreases as the temperature increases: thermal pollution.
Example
Determine the concentration of dissolved oxygen in water when PO2 = 0.21 atm.
Oxygen Demand
-The presence of dissolved oxygen means that, like the air, natural waters are oxidizing:
-Primary half-reaction:
O2 + 4H+ + 4e-  2H2O(l)
-This is capable of oxidizing organic matter.
-As an approximation, assume organic matter is carbohydrate: CH2O
CH2O(aq) + O2(aq)  CO2(g) + H2O(l)
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BOD-Biochemical (or Biological) Oxygen Demand
-Bacteria in water can consume organic matter.
-BOD is the measure of the quantity of oxygen consumed during the oxidation of the organic
matter.
-Measure the dissolved oxygen (DO) content in a sample of water, add bacteria, seal the
container for 5 days, measure again. The change in DO content is the BOD.
-Typical surface water has a BOD of 0.7 mg/L, well below the maximum solubility of oxygen.
-Does not account for oxygen needs of fish or the like.
-Polluted water and sewage can have BODs of 10's to 100's mg/L.
DO: Problem 9-3
COD-Chemical Oxygen Demand
-If you are not patient, COD gives a fast result.
-dichromate, Cr2O7-2 is a strong oxidizing agent that will oxidize most organic matter:
Cr2O7-2 + 14H+ + 6e-  2Cr3+ + 7H2O
-Reaction is quick, but:
-oxidizes all organic matter, including that not readily available to organisms.
-oxidizes some inorganics, such as Cl-Therefore COD gives a high result.
Example
Find the balanced reaction between dichromate and a typical carbohydrate.
DO: Problem 9-5
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TOC and DOC
-Total Organic Carbon: all suspended and dissolved organic matter.
-Dissolved Organic Carbon: only that carbon which is actually dissolved.
Decomposition of Organic Matter in Water
-Oxidation of organic matter is an aerobic process.
-In anaerobic conditions, oxygen is absent, and reduction processes tend to occur:
2CH2O  CH4 + CO2
-Fermentation.
-CH4 formed in bogs is called swamp or marsh gas.
-Also occurs in landfills and, intentionally, in digestors for solid waste/sewage.
-Moving water is rarely low in oxygen.
-In lakes, the water can stratify in summer.
-In winter, lakes will be at about 0°C at the surface and 4°C near the bottom.
-No mixing.
-In spring, surface water starts to warm, density increases, and winds can turnover lake.
-Mixes lake and brings DO to lower levels.
-In summer, surface water warms further, mixing stops.
-Epilimnion-surface layer, warm, high DO
-Thermocline-layer where temp drops
-Hypolimniom-lower layer, can have low or high DO depending upon degree of
biological activity in the lake.
-Eutrophic: low DO, high biological activity.
-Oligotrophic: high DO, low biological activity.
-In fall, another turnover, winter stratifies again.
-Figure 9-2
-In high DO conditions, get oxidized species: HCO3-, SO4-2, NO3-, Fe(OH)3
-In low DO conditions, get reduced species: CH4, H2S, NH3/NH4+, Fe+2
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The pE Scale
-pE is a redox equivalent to pH. Also called Eh
-Low pE, highly reducing (e- available) species available.
-High pE, highly oxidizing (e- deficient) species available.
-Usually one species is important, O2 for oxidizing species.
-The pE-pH diagram shows regions of stability of various redox species
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Sulfur Compounds
-Table 9-1 shows common oxidation states of sulfur compounds.
-All but the sulfite species are accessible thermodynamically under the proper conditions.
-Sulfates are present under high pE (oxidizing) conditions.
-Sulfides are present under low pE conditions (anaerobic).
-Sulfur can be formed in acidic, moderate pE conditions.
-H2S is formed in acidic, anaerobic conditions, gives rise to smell of rotten eggs.
-gets oxidized back to sulfate in the atmosphere.
Acid Mine Drainage
skip
Nitrogen Compounds
-The two common aqueous forms are the -3 state, NH3 and NH4+, and +5, NO3-.
-Table 9-2
-These are the extremes
-N2 is a mostly insoluble gas.
-Some N2 can be fixed into amines by certain bacteria.
-amines turn into ammonia in microbial decay.
-nitrification: oxidation of ammonia to nitrate
-catalyzed by various microorganisms in aerobic environments.
-denitrification: reduction of nitrate to nitrite or further to ammonia.
-anaerobic environments
-nitrate can be used as an oxidizing agent in anaerobic conditions:
4H+ + 4NO3- + 5CH2O  2N2 + 5CO2 + 7H2O
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Acid-Base Chemistry in Natural Waters: The Carbonate System
-Natural water is never "pure" in the laboratory sense.
-It is in contact with air, which has soluble gases.
-The most important are oxygen (as above) and carbon dioxide (CO2).
-It is, or has been, in contact with minerals, of which calcium carbonate (CaCO3) is a model
system.
The Equilibrium of Water and CO2(g)
-Carbon dioxide is slightly soluble in water, forming carbonic acid:
CO2(g) + H2O(l)  H2CO3(aq)
(1)
-Actually, dissolved CO2(aq) is important, but your text lumps together the two species:
[H2CO3(aq)]text = [H2CO3(aq)] + [CO2(aq)]
-This is not terribly important, other than the fact that most (99.8%) of the dissolved species
is CO2(aq).
-Reaction (1) has the following equilibrium constant (Henry's Law):
[H CO (aq)]
K1  2 3
 3.44 102
PCO 2
Example
What is the concentration of carbonic acid in water in contact with air? Normal air has a CO2
concentration of 373 ppm.
-Carbonic acid is a weak diprotic acid. The first acid dissociation is given by:
H2CO3(aq) + H2O(l)  H3O+(aq) + HCO3-(aq)
-The equilibrium constant for this reaction should be familiar as an acid dissociation:
[H O  (aq)][ HCO 3 (aq)]
K 2  K a1  3
 4.5 10 7
[H 2CO3 (aq)]
(2)
Example
What is the pH of a solution saturated with CO2?
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The CO2-Carbonate System
-Above we explored the acid properties of CO2 in contact with water.
-Natural water is also in contact with minerals.
-Limestone is an important type of mineral and serves as a good model system. Limestone is
made of calcite, one form of the compound calcium carbonate, CaCO3.
-Silicate minerals are much less soluble than calcite.
-Magnesium carbonate is important, too, and we will expand our discussion later.
-Other carbonates are of low abundance.
-Calcium carbonate is slightly soluble in water:
CaCO3(s)  Ca+2(aq) + CO3-2(aq)
(3)
-The equilibrium constant for this reaction is a solubility product:
K3  K sp  [Ca2 ][CO32 ]  4.6 109
Example
What is the solubility, or concentrations of Ca+2 and CO3-2, in a saturated solution of CaCO3,
assuming no other reactions occur?
-However, carbonate is a weak base:
CO32-(aq) + H2O(l)  HCO3-(aq) + OH-(aq)
(4)
-This is the base equivalent of the acid dissociation of the bicarbonate ion, HCO3 :
HCO3-(aq) + H2O(l)  CO3-2(aq) + H3O+(aq)
Ka2 = 4.7 x 10-11
-The base reaction has the equilibrium constant:
[HCO 3 ][OH ] K w 1.0 10 14
K 4  Kb 


 2.110 4
2
11
[CO3 ]
K a 2 4.7 10
-This complicates the solubility because of the common ion effect. As the carbonate reacts to
form bicarbonate, more calcium dissolves.
-A simplistic approach assumes that the bicarbonate ion is more important, and that little CO3-2 is
present. This allows us to combine reactions 3 and 4 to get:
CaCO3(s) + H2O(l)  Ca+2(aq) + HCO3-(aq) + OH-(aq)
(5)
-When we add reactions, we multiply equilibrium constants:
K 5  K 3 K 4  [Ca 2 ][ HCO 3 ][OH ]  9.7 10 13
_In this case there is a 1:1:1 ratio of the three products.
Example
Assuming reaction (5) is important, what are the concentrations of calcium, bicarbonate, and
hydroxide? What is the pH?
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Water in Equilibrium with Both CaCO3 and CO2
-If we allow both CaCO3 and CO2 to come into contact with water, we have yet a more complex
system. Calcium carbonate is basic, carbon dioxide is acidic. Acids react with bases. Therefore,
calcium carbonate will be more soluble in water in contact with CO3.
-We simplify all the possible chemistry by combining reactions (1) through (4), and include the
dissociation of water:
H3O+(aq) + OH-(aq)  2H2O(l)
(6)
+
-14
Kw = [H3O ][OH ] = 1.0 x 10
-The net reaction is:
CaCO3(s) + CO2(g) + H2O(l)  Ca+2(aq) + 2HCO3-(aq)
(7)
-The overall equilibrium constant is:
K sp K b K H K a1
[Ca 2 ][ HCO 3 ]2
K7 
 1.5 106 
Kw
PCO 2
Example
Find the concentrations of calcium and bicarbonate in this system, and find the pH.
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Alkalinity Indices for Natural Waters
-Alkalinity is the measure of a solution's ability to react with H+ ions.
-pH is an inadequate measure of a solution's ability to absorb H+ because if the solution is
buffered (as natural water is) the pH depends on the ratio of acid to base, not the total
concentration of either.
-In natural waters, the carbonate/bicarbonate system is the most important buffer system. As
such, we should look at the predominance diagram for this system:
1.00
HCO 3-
H 2 CO 3
CO 32-
Fraction
0.75
0.50
0.25
pK 2
pK1
0.00
0
2
4
6
8
10
12
14
16
pH
-Above a pH of 10.3, the main form is carbonate. Between 6.3 and 10.3, bicarbonate is the
primary species, which is the pH range of most natural waters. Below a pH of 6.3, carbonic acid
predominates.
-Total alkalinity is the number of moles (or millimoles) needed to react with all the carbonate,
bicarbonate, and hydroxide in one liter of sample. The endpoint is based on the color change in
the indicator methyl orange, which changes color from yellow (above pH 4.4) to red (below
pH 3.1).
total alk.  2[CO32- ]  [HCO 3- ]  [OH- ]  [H  ]
-As you can see, when the methyl orange turns red below a pH of 4, upon addition of H+, all the
carbonate and bicarbonate will have been converted to carbonic acid.
-Can include minor bases such as ammonia (NH3), borate (B(OH)4-), trihydrogen silate
(H3SiO4-), hydrogen phosphate (HPO4-2), hydrogen sulfide (HS-), and organic matter. These ions
are usually present in concentrations low enough to ignore.
-A high alkalinity value is indicative of fertile lakes. Algae metabolize CO2, which removes
bicarbonate and forms carbonate, CO3-2, which then precipitates out as calcium carbonate. The
pH in such a system is maintained.
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Do Problem 9-18
Hardness Index
-Water in contact with limestone (CaCO3) or dolomitic limestone (CaMg(CO3)2) can have large
concentrations of Ca+2 and Mg+2 ions. The bicarbonate salts are quite soluble, and as we have
seen above, at neutral pH levels, bicarbonate is the dominant form of the species.
-Hard water is common in ground water.
-High concentrations of Ca+2 and Mg+2 have two major consequences:
1. Soaps are made from the sodium salts of the conjugate bases of fatty acids, NaO2C-R.
-Calcium and magnesium salts of these fatty acids are insoluble in water.
-Get soap scum--generally easily removed with vinegar or similar cleaners.
2. When heated, water with high concentrations of bicarbonate undergoes the following:
2HCO3-  CO3-2 + H2CO3(aq)
H2CO3(aq)  H2O(l) + CO2(g)
CO3-2 + Ca+2  CaCO3(s)
-In this case, one gets lime scale
-A nuisance in cooking, but a pipe clogger in plumbing.
-Hardness is not a health hazard.
-Hardness index is given by:
hardness = [Ca+2] + [Mg+2]
-Uses EDTA in a titration.
-Hardness reported as an equivalent amount of CaCO3 in mg/L
Do Problem 9-21
References
Iowa Groundwater Basics: http://www.igsb.uiowa.edu/GWBASICS/
USGS Water: http://water.usgs.gov/
USGS Explanation of Hardness: http://water.usgs.gov/owq/Explanation.html
WUSTL: Water Hardness:
http://www.chemistry.wustl.edu/~edudev/LabTutorials/Water/FreshWater/hardness.html
GWPC: State Report Sheets: http://www.gwpc.org/gwreport/states.htm
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