Alkalinity

Definition

Alkalinity is the name given to the quantitative capacity of water to neutralize an acid to the equivalence point of carbonate or bicarbonate (Water Quality Association, 2000). The total alkalinity of sea water was defined by Dickson (1981) as “. . . the number of moles of hydrogen ion equivalent to the excess of proton acceptors (bases formed from weak acids with a dissociation constant K ≤ 10^−4.5 at 25 °C and zero ionic strength) over proton donors (acids with K > 10^−4.5) in 1 kg of sample.” For the compounds found in water, the total alkalinity (A_T) is expressed as:

$$ \begin{array}{c}{\mathrm{A}}_{\mathrm{T}} = \left[{{\mathrm{H}\mathrm{CO}}_3}^{-}\right]+2\left[{{\mathrm{CO}}_3}^{-2}\right]+\left[\mathrm{B}{{\left(\mathrm{O}\mathrm{H}\right)}_4}^{-}\right]+\left[{\mathrm{OH}}^{-}\right]\\ {}+\left[{{\mathrm{H}\mathrm{PO}}_4}^{-2}\right]+2\left[{{\mathrm{PO}}_4}^{-3}\right]+\left[{\mathrm{H}}_3{{\mathrm{SiO}}_4}^{-}\right]+\left[{\mathrm{NH}}_3\right]\\ {}+\left[{\mathrm{H}\mathrm{S}}^{-}\right]-{\left[{\mathrm{H}}^{+}\right]}_{\mathrm{F}}-\left[{{\mathrm{H}\mathrm{S}\mathrm{O}}_4}^{-}\right]-\left[\mathrm{H}\mathrm{F}\right]-\left[{\mathrm{H}}_3{\mathrm{PO}}_4\right]\end{array} $$

where [H⁺]_F is the free concentration of the hydrogen ion (Dickson, 2010).

In natural waters, carbonate alkalinity, A_C = [HCO₃^–] + 2[CO₃⁻²], tends to comprise most of the A_T due to the common occurrence and dissolution of carbonate rocks and presence of carbon dioxide in the atmosphere. Other common natural components of A_T are borate, hydroxide, phosphate, silicate, nitrate, ammonia, sulfide, and the conjugate bases of some organic acids. In anoxic conditions the relative role of sulfide, ammonia, and phosphate components of A_T increases (Volkov et al., 1998). In coastal regions, especially estuaries, dissolved organic matter can significantly contribute to A_T (Kim and Lee, 2009).

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 (i.e., total titratable alkalinity). 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.

An addition (or removal) of CO₂ to a solution does not change the alkalinity. 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 alkalinity. In open ocean waters, alkalinity can be connected with salinity and temperature with a functional dependence (Lee et al., 2006). Rivers can act as either a source or a sink of alkalinity.

The actual units for the alkalinity titration are moles or equivalents per volume (mol L⁻¹ or Eq L⁻¹). They can be converted to mol kg⁻¹ or, in terms of calcium carbonate, to mg CaCO₃ L⁻¹.

Cross-references

• pH

• Water Quality