Editing Alkalinity (section)

==Theoretical treatment==
In typical [[groundwater]] or [[seawater]], the measured total alkalinity is set equal to:

: A<sub>T</sub> = [{{chem|HCO|3|-}}]<sub>T</sub> + 2[{{chem|CO|3|2-}}]<sub>T</sub> + [{{chem|B(OH)|4|-}}]<sub>T</sub> + [OH<sup>−</sup>]<sub>T</sub> + 2[{{chem|PO|4|3-}}]<sub>T</sub> + [{{chem|HPO|4|2-}}]<sub>T</sub> + [{{chem|SiO(OH)|3|-}}]<sub>T</sub> − [H<sup>+</sup>]<sub>sws</sub> − [{{chem|HSO|4|-}}]

(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&nbsp;4.5. At this point, all the bases of interest have been protonated to the zero level species, hence they no longer cause alkalinity. In the carbonate system the bicarbonate ions [{{chem|HCO|3|-}}] and the carbonate ions [{{chem|CO|3|2-}}] have become converted to carbonic acid [H<sub>2</sub>CO<sub>3</sub>] at this pH. This pH is also called the CO<sub>2</sub> equivalence point where the major component in water is dissolved CO<sub>2</sub> which is converted to H<sub>2</sub>CO<sub>3</sub> in an aqueous solution. There are no strong acids or bases at this point. Therefore, the alkalinity is modeled and quantified with respect to the CO<sub>2</sub> equivalence point. Because the alkalinity is measured with respect to the CO<sub>2</sub> equivalence point, the dissolution of CO<sub>2</sub>, although it adds acid and dissolved inorganic carbon, does not change the alkalinity. In natural conditions, the dissolution of basic rocks and addition of ammonia [NH<sub>3</sub>] or organic amines leads to the addition of base to natural waters at the CO<sub>2</sub> equivalence point. The dissolved base in water increases the pH and titrates an equivalent amount of CO<sub>2</sub> to bicarbonate ion and carbonate ion. At equilibrium, the water contains a certain amount of alkalinity contributed by the concentration of weak acid anions. Conversely, the addition of acid converts weak acid anions to CO<sub>2</sub> and continuous addition of strong acids can cause the alkalinity to become less than zero.<ref>Benjamin. Mark M. 2015. ''Water Chemistry''. 2nd Ed. Long Grove, Illinois: Waveland Press, Inc.</ref> For example, the following reactions take place during the addition of acid to a typical seawater solution:

: {{chem|B(OH)|4|−}} + H<sup>+</sup> → B(OH)<sub>3</sub> + H<sub>2</sub>O

: OH<sup>−</sup> + H<sup>+</sup> → H<sub>2</sub>O

: {{chem|PO|4|3−}} + 2 H<sup>+</sup> → {{chem|H|2|PO|4|−}}

: {{chem|HPO|4|2−}} + H<sup>+</sup> → {{chem|H|2|PO|4|−}}

: [{{chem|SiO(OH)|3|-}}] + H<sup>+</sup> → [Si(OH)<sub>4</sub>]

It can be seen from the above protonation reactions that most bases consume one proton (H<sup>+</sup>) to become a neutral species, thus increasing alkalinity by one per equivalent. {{chem|CO|3|2-}} however, will consume two protons before becoming a zero-level species (CO<sub>2</sub>), thus it increases alkalinity by two per mole of {{chem|CO|3|2-}}. [H<sup>+</sup>] and [{{chem|HSO|4|-}}] decrease alkalinity, as they act as sources of protons. They are often represented collectively as [H<sup>+</sup>]<sub>T</sub>.

Alkalinity is typically reported as mg/L ''as'' CaCO<sub>3</sub>.  (The conjunction "as" is appropriate in this case because the alkalinity results from a mixture of ions but is reported "as if" all of this is due to CaCO<sub>3</sub>.)  This can be converted into milliequivalents per Liter (meq/L) by dividing by 50 (the approximate [[Molar mass|MW]] of CaCO<sub>3</sub> divided by 2).