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). Commercially, as in the pool industry, alkalinity might also be given in the unit ppm or parts per million.
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).
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 titrating 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:
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).
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−soln-1
At neutral pH's:
CO2 + H2O → HCO3− + H+
At high pH's:
CO2 + H2O → CO3−2 + 2H+
Natural sources of acid neutralizing capacity in low alkalinity lakes of the Precambrian Shield. (Experimental Lakes Area, Ontario)
May 16, 1986; Natural Sources of Acid Neutralizing Capacity in Low alkalinity Lakes of the Precambrian Shiled IN ACID-SENSITIVE LAKES,...