Alkalinity or AT is a measure of the ability of a solution to neutralize acids to the equivalence point of carbonate or bicarbonate. Alkalinity is closely related to the acid neutralizing capacity (ANC) of a solution and ANC is often incorrectly used to refer to alkalinity. However, the acid neutralizing capacity refers to the combination of the solution and solids present (e.g., suspended matter, or aquifer solids), and the contribution of solids can dominate the ANC (see carbonate minerals below).
The alkalinity is equal to the stoichiometric sum of the bases in solution. In the natural environment carbonate alkalinity tends to make up most of the total alkalinity due to the common occurrence and dissolution of carbonate rocks and presence of carbon dioxide in the atmosphere. Other common natural components that can contribute to alkalinity include borate, hydroxide, phosphate, silicate, nitrate, dissolved ammonia, the conjugate bases of some organic acids and sulfide. Solutions produced in a laboratory may contain a virtually limitless number of bases that contribute to alkalinity. Alkalinity is usually given in the unit mEq/L (milliequivalent per liter).
Alkalinity is sometimes incorrectly used interchangeably with basicity. For example, the pH of a solution can be lowered by the addition of CO2. This will reduce the basicity; however, the alkalinity will remain unchanged (see example below).
Theoretical treatment of alkalinity
AT = [HCO3−]T + 2[CO3−2]T + [B(OH)4−]T + [OH−]T + 2[PO4−3]T + [HPO4−2]T + [SiO(OH)3−]T − [H+]sws − [HSO4−]
(Subscript T indicates the total concentration of the species in the solution as measured. This is opposed to the free concentration, which takes into account the significant amount of ion pair interactions that occur in seawater.)
Alkalinity can be measured by a sample with a strong acid until all the buffering capacity of the aforementioned ions above the pH of bicarbonate or carbonate is consumed. This point is functionally set to pH 4.5. At this point, all the bases of interest have been protonated to the zero level species, hence they no longer cause alkalinity. For example, the following reactions take place during the addition of acid to a typical seawater solution:
- HCO3− + H+ → CO2 + H2O
- CO3−2 + 2H+ → CO2 + H2O
- B(OH)4− + H+ → B(OH)3 + H2O
- OH− + H+ → H2O
- PO4−3 + 2H+ → H2PO4−
- HPO4−2 + H+ → H2PO4−
- [SiO(OH)3−] + H+ → [Si(OH)40]
It can be seen from the above protonation reactions that most bases consume one proton (H+) to become a neutral species, thus increasing alkalinity by one per equivalent. CO3−2 however, will consume two protons before becoming a zero level species (CO2), thus it increases alkalinity by two per mole of CO3−2. [H+] and [HSO4−] decrease alkalintiy, as they act as sources of protons. They are often represented collectively as [H+]T.
Alkalinity is typically reported as mg/L as CaCO3. This can be converted into milliEquivalents per Liter (mEq/L) by dividing by 50 (the approximate MW of CaCO3/2).
Sum of contributing species
The following equations demonstrate the relative contributions of each component to the alkalinity of a typical seawater sample. Contributions are in μmol/kg-H2O and are obtained from A Handbook of Methods for the analysis of carbon dioxide parameters in seawater ","(Salinity = 35, pH = 8.1, Temp = 25°C).
AT = [HCO3−]T + 2[CO3−2]T + [B(OH)4−]T + [OH−]T + 3[PO4−3]T + [HPO4−2]T + [SiO(OH)3−]T − [H+] − [HSO4−] − [HF]
Phosphates and Silicate, being nutrients are typically negligible. At pH = 8.1 [HSO4−] and [HF] are also negligible. So,
AT = [HCO3-]T + 2[CO3−2]T + [B(OH)4−]T + [OH−]T − [H+]
AT = 1830 + 2*270 + 100 + 10 − 0.01
AT = 2480 μmol/kg−H2O
Addition of CO2
The addition (or removal) of CO2 to a solution does not change the alkalinity. This is because the net reaction produces the same number of equivalents of positively contributing species (H+) as negative contributing species (HCO3- and/or CO3--).
At neutral pH's:
CO2 + H2O → HCO3− + H+
At high pH's:
CO2 + H2O → CO3−2 + 2H+
Dissolution of carbonate rock
Addition of CO2 to a solution in contact with a solid can affect the alkalinity, especially for carbonate minerals in contact with groundwater or seawater . The dissolution (or precipitation) of carbonate rock has a strong influence on the alkalinity. This is because carbonate rock is composed of CaCO3 and its dissociation will add Ca+2 and CO3−2 into solution. Ca+2 will not influence alkalinity, but CO3−2 will increase alkalinity by 2 units.
- Alkali soils
- Biological pump
- Global Ocean Data Analysis Project
- Ocean acidification
Carbonate system calculators
The following packages calculate the state of the carbonate system in seawater (including pH):
- CO2SYS, a stand-alone executable (also available in a version for Microsoft Excel/VBA)
- seacarb, a R package for Windows, Mac OS X and Linux (also available here)
- CSYS, a Matlab script
- Holmes-Farley, Randy. "Chemistry and the Aquarium," Advanced Aquarist's Online Magazine. Alkalinity as it pertains to salt-water aquariums.
- DOE (1994) ","Handbook of methods for the analysis of the various parameters of the carbon dioxide system in sea water. Version 2, A. G. Dickson & C. Goyet, eds. ORNL/CDIAC-74ca:Alcalinitat
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