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Alkalinity Except for waters having high pH (greater than about 9.50) and some others having unusual chemical The alkalinity of a solution may be defined as the composition, especially water associatedwith petroleum capacity for solutesit contains to react with and neutralize and natural gas or water having much dissolved organic acid. The property of alkalinity must be determined by carbon, the alkalinity of natural waters can be assigned titration with a strong acid, and the end point of the entirely to dissolved bicarbonat’eand carbonate without titration is the pH at which virtually all solutes contrib- serious error. The important contribution of short-chain uting to alkalinity have reacted. The end-point pH that aliphatic acid anions to titratable alkalinity in water from should be used in this titration is a function of the kinds certain oil fields was pointed out by Willey and others of solute speciesresponsible for the alkalinity and their (1975). concentrations. However, the correct titration end point Sources of Alkalinity for a particular solution can be identified from the experi- mental data when the speciesinvolved are unknown. It is The principal source of carbon dioxide speciesthat the point at which the rate of change of pH per added produce alkalinity in surface or ground water is the CO2 volume of titrant (dpH/dV,,,d) is at a maximum. As dis- gas fraction of the atmosphere, Ior the atmospheric gases sociation constants in table 33 show, the ratio [HCOs]: present in the soil or in the unsaturated zone lying [HzCO~] will be near 1OO:l at pH 4.4, and the ratio between the surface of the land and the water table. The [HC03J:[C032-] will have a similar value at pH 8.3 at CO2 content of the atmosphere is near 0.03 percent by temperatures near 20°C. The best values for the end volume. Soil-zone and unsaturated-zoneair can be sub- points for a particular sample depend on ionic strength stantially enriched in carbon dioxide, usually owing to and temperature. Analytical procedures may specify a respiration by plants and the oxidation of organic matter. pH value between 5.1 and 4.5, or that of the methyl- In some natural systems there may be sources of orange end point (about pH 4.0-4.6). Sometimes, how- carbon dioxide other than dissolution of atmospheric or ever, an alkalinity above the phenolphthalein end point soil-zone CO2. Possible major local sourcesinclude bio- (about pH 8.3) is also specified. Thus one may find terms logically mediated sulfate reduction and metamorphism such as “methyl-orange alkalinity,” or its equivalent, of carbonate rocks. In some areas,outgassing from rocks “total alkalinity,” and “phenolphthalein alkalinity.” in the mantle 15 km or more below the surface has been Dilute solutions such as rainwater require special pro- suggested(Irwin and Barnes, 1980). Indications of source cedures for this determination (Stumm and Morgan, can sometimes be obtained from stable isotope (S13C) 1981, p. 226-229). data. Several different solute species contribute to the From studies of 613Cvalues in dissolved HC03. in alkalinity of water as defined above, and titration with 15 oil and gas fields, Carothers and Kharaka (1980) acid does not specifically identify them. The property of concluded that the decarboxylation of acetate and other alkalinity can be expressedin quantitative terms in various short-chain aliphatic acids was an important CO2 source ways. The most common practice is to report it in terms in these waters. This processalso produces methane and of an equivalent amount of calcium carbonate. It could other hydrocarbon gases. also be expressed in milliequivalents per liter, where Carbon dioxide species are important participants meq/L is l/50 times mg/L CaC03. in reactions tht control the pH of natural waters. Various In almost all natural waters the alkalinity is produced aspectsof this fact were discussedin the section on pH. by the dissolved carbon dioxide species,bicarbonate and Reactions among the alkalinity-related species,aqueous carbonate, and the end points mentioned above were COZ, HzCOs(aq), HC03-, and COs2-, and directly pH- selected with this in mind. Analyses in this book, and related species,OH- and H’, are relatively fast and can most others in current geochemical literature, follow the be evaluated with chemical equilibrium models. Rates of convention of reporting titrated alkalinity in terms of the equilibration between solute species and gaseous CO2 equivalent amount of bicarbonate and carbonate. across a phase boundary are slower, and water bodies The more important noncarbonate contributors to exposed to the atmosphere may not be in equilibrium alkalinity include hydroxide, silicate, borate, and organic with it at all times. The oceans are a major factor in ligands, especially acetate and propionate. Rarely, other maintaining atmospheric CO2 contents. It may be of speciessuch as NHdOH or HS- may contribute signifi- interest to note that carbonic acid, HzC03, is convention- cantly to alkalinity. If alkalinity is expressed in milli- ally used to represent all the dissolved undissociated equivalents per liter, or as CaC03, the contributions carbon dioxide. In actuality, only about 0.01 percent of from thesespecies will affect the cation-anion balance of the dissolved carbon dioxide is present in this form. We the analysis only if some of them are determined by other will usethe H2CO3convention in discussingthese systems, methodsand are thus included in the balancecomputation however, as the choice of terminology has no practical in two places. effect on final results. 106 Study and Interpretation of the Chemical Characteristics of Natural Water Relationships among the dissolved carbon dioxide was considered, and activity coefficients were assumed speciesand pH are summarized in figure 19, which is a to be unity. These and other simplifications limit the CO2 speciesdistribution diagram. The lines on this graph practical usefulnessof the diagram to some extent, but were computed from the first and second dissociation modified forms can be prepared using equilibrium con- equilibrium expressions, stants for other temperaturesand including calculated or assumedionic strengths. Diagrams of this type are useful FLCO,-1 for summarizing species’ pH dependenceand for other =K, [H+]-’ purposes(Butler, 1964, p. 120). W&O,1 Figure 19 indicates that carbon dioxide speciescan contribute small amounts to alkalinity down to pH 4.0. and The value of the HCOa-:H2COa ratio changeswith tem- mm perature and ionic strength. Barnes (1964) showed that =K, [H+]-‘, the correct titration end point pH may rangebetween 4.4 WC%1 and 5.4 and recommended that the titration be done at the sample collection site. Similar variation can occur in and an assumption that the total alkalinity is the sum of the carbonate:bicarbonateend point. The diagram shows the carbonate and bicarbonate activities. Values for K1 why small concentrations of carbonate cannot be deter- and KZ at various temperatures are given in table 33 mined very accurately by titration. The pH at which (appendix). The contribution of hydroxide to alkalinity carbonate constitutes 1 percent of the total dissolved can become significant above about pH 10, where the carbon dioxide species,about 8.3, is where the titration activity of OH- is about 1.7 mg/L. end point for carbonate would generally be placed. This The ratio of molar activities in the dissociation is a low enough pH that about 1 percent of the total now equations is a function of pH, and it is not necessaryto also is in the form of HaCOa. If a water contains much know the total amountsof the dissolvedspecies. However, bicarbonate and only a little carbonate, the overlapping in practice it is easierto usethe percentagecomposition, of the two steps in the vicinity of pH 8.3 may make it and the calculations for the graph were made using an impossible to determine the carbonate even to the nearest arbitrary total of 100for activities of the dissolved carbon milligram per liter. Becauseof the overlap, the change in dioxide species.The graph showsthe effect of temperature pH as acid is added may be gradual rather than abrupt at from 0°C to 50°C at 1 atmospherepressure; no gasphase this end point. Usually, if the carbonate concentration is EXPLANATION 1 Temperature 80 - . * . - 50%. > k - 25oC. HCO; J I 4.0 5.0 6.0 7.0 8.0 9.0 10.0 11.0 12.0 13.0 PH Figure 19. Percentages of dissolved carbon dioxide species activities at 1 atmosphere pressure and various tempera- tures as a function of pH. Significance of Properties and Constituents Reported in Water Analyses 107 small compared with the bicarbonate concentration, a commonly predicted is a rise in average surface tempera- value for carbonate can be calculated from the equilibrium ture of the Earth owing to the so-called greenhouse equations more accurately than it can be measured by effect. Carbon dioxide absorbs infrared radiation from titration. the Earth’s surface and prevents t he escape of some of the As noted in the earlier discussions of pH and calcium Sun’s energy that would otherwise be lost (Hileman, carbonate equilibria, a measurement of pH and of total 1982). alkalinity provides enough data to calculate activities of In a summary article Lieth ( 1963) gave some figures both the dissociated and undissociated carbon dioxide on productivity, defined as the amount of carbon dioxide species. A rigorous discussion of the chemical principles converted into organic matter per unit land or water area involved in evaluating alkalinity and acidity was given per year. In a middle-latitude forest, the estimated rate by Kramer (1982). was 15 metric tons of COZ per hectare per year. A The Carbon Cycle tropical forest was estimated to have a rate 2% times as high.

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