Alkalinity

In 1884, Professor Wilhelm (William) Dittmar of Anderson College, now the University of Strathclyde, analysed 77 pristine seawater samples from around the world brought back by the Challenger expedition. He found that in seawater the major ions were in a fixed ratio, confirming the hypothesis of Johan Georg Forchhammer, that is now known as the Principle of Constant Proportions. However, there was one exception. Dittmar found that the concentration of calcium was slightly greater in the deep ocean, and named this increase alkalinity.

Also in 1884, Svante Arrhenius submitted his PhD theses in which he advocated the existence of ions in solution, and defined acids as hydronium ion donors and bases as hydroxide ion donors. For that work, he received the Nobel Prize in Chemistry in 1903. See also Svante Arrhenius#Ionic disassociation.

Alkalinity roughly refers to the molar amount of bases in a solution that can be converted to uncharged species by a strong acid. For example, 1 mole of HCO⁻
₃ in solution represents 1 molar equivalent, while 1 mole of CO²⁻
₃ is 2 molar equivalents because twice as many H⁺ ions would be necessary to balance the charge. The total charge of a solution always equals zero.[8] This leads to a parallel definition of alkalinity that is based upon the charge balance of ions in a solution.

${\displaystyle \sum ({\text{cations}})=\sum ({\text{anions}})}$

Certain ions, including Na⁺, K⁺, Ca²⁺, Mg²⁺, Cl⁻, SO²⁻
₄, and NO⁻
₃ are "conservative" such that they are unaffected by changes in temperature, pressure or pH.[8] Others such as HCO⁻
₃ are affected by changes in pH, temperature, and pressure. By isolating the conservative ions on one side of this charge balance equation, the nonconservative ions which accept or donate protons and thus define alkalinity are clustered on the other side of the equation.

${\displaystyle {\begin{aligned}&\sum ({\text{conservative cations}})-\sum ({\text{conservative anions}})=\\&\quad [\mathrm {HCO_{3}^{-}} ]+2[\mathrm {CO_{3}^{2-}} ]+[\mathrm {B(OH)_{4}^{-}} ]+[\mathrm {OH^{-}} ]+[\mathrm {HPO_{4}^{2-}} ]+2[\mathrm {PO_{4}^{3-}} ]+[\mathrm {H_{3}SiO_{4}^{-}} ]+[\mathrm {NH_{3}} ]+[\mathrm {HS^{-}} ]-[\mathrm {H^{+}} ]-[\mathrm {HSO_{4}^{-}} ]-[\mathrm {HF} ]-[\mathrm {H_{3}PO_{4}} ]-[\mathrm {HNO_{2}} ]\end{aligned}}}$

This combined charge balance and proton balance is called total alkalinity.[9] Total alkalinity is not (much) affected by temperature, pressure, or pH, and is thus itself a conservative measurement, which increases its usefulness in aquatic systems. All anions except HCO⁻
₃ and CO²⁻
₃ have low concentrations in Earth's surface water (streams, rivers, and lakes). Thus carbonate alkalinity, which is equal to [HCO⁻
₃] + 2[CO²⁻
₃] is also approximately equal to the total alkalinity in surface water.[8]

Alkalinity measures the ability of a solution to neutralize acids to the equivalence point of carbonate or bicarbonate, defined as pH 4.5 for many oceanographic/limnological studies.[10] The alkalinity is equal to the stoichiometric sum of the bases in solution. In most Earth surface waters carbonate alkalinity tends to make up most of the total alkalinity due to the common occurrence and dissolution of carbonate rocks and other geological weathering processes that produce carbonate anions. Other common natural components that can contribute to alkalinity include borate, hydroxide, phosphate, silicate, dissolved ammonia, and the conjugate bases of organic acids (e.g., acetate). Solutions produced in a laboratory may contain a virtually limitless number of species that contribute to alkalinity. Alkalinity is frequently given as molar equivalents per liter of solution or per kilogram of solvent. In commercial (e.g. the swimming pool industry) and regulatory contexts, alkalinity might also be given in parts per million of equivalent calcium carbonate (ppm CaCO₃). Alkalinity is sometimes incorrectly used interchangeably with basicity. For example, the addition of CO₂ lowers the pH of a solution, thus reducing basicity while alkalinity remains unchanged (see example below).

A variety of titrants, endpoints, and indicators are specified for various alkalinity measurement methods. Hydrochloric and sulfuric acids are common acid titrants, while phenolpthalein, methyl red, and bromocresol green are common indicators.[11]

In typical groundwater or seawater, the measured total alkalinity is set equal to:

A_T = [HCO⁻
₃]_T + 2[CO²⁻
₃]_T + [B(OH)⁻
₄]_T + [OH⁻]_T + 2[PO³⁻
₄]_T + [HPO²⁻
₄]_T + [SiO(OH)⁻
₃]_T − [H⁺]_sws − [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 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 [HCO⁻
₃] and the carbonate ions [CO²⁻
₃] have become converted to carbonic acid [H₂CO₃] at this pH. This pH is also called the CO₂ equivalence point where the major component in water is dissolved CO₂ which is converted to H₂CO₃ 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 CO₂ equivalence point. Because the alkalinity is measured with respect to the CO₂ equivalence point, the dissolution of CO₂, 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 [NH₃] or organic amines leads to the addition of base to natural waters at the CO₂ equivalence point. The dissolved base in water increases the pH and titrates an equivalent amount of CO₂ 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 CO₂ and continuous addition of strong acids can cause the alkalinity to become less than zero.[12] For example, the following reactions take place during the addition of acid to a typical seawater solution:

B(OH)⁻
₄ + H⁺ → B(OH)₃ + H₂O

OH⁻ + H⁺ → H₂O

PO³⁻
₄ + 2 H⁺ → H
₂PO⁻
₄

HPO²⁻
₄ + H⁺ → H
₂PO⁻
₄

[SiO(OH)⁻
₃] + H⁺ → [Si(OH)₄]

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. CO²⁻
₃ however, will consume two protons before becoming a zero-level species (CO₂), thus it increases alkalinity by two per mole of CO²⁻
₃. [H⁺] and [HSO⁻
₄] decrease alkalinity, as they act as sources of protons. They are often represented collectively as [H⁺]_T.

Alkalinity is typically reported as mg/L as CaCO₃. (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 CaCO₃.) This can be converted into milliequivalents per Liter (meq/L) by dividing by 50 (the approximate MW of CaCO₃ divided by 2).

In the ocean, alkalinity is completely dominated by carbonate and bicarbonate plus a small contribution from borate.[4]

Thus the chemical equation for alkalinity in seawater is:

A_T = [HCO₃^-] + 2[CO₃^2-] + [B(OH)₄^-]

There are many methods of alkalinity generation in the ocean. Perhaps the most well known is the dissolution of calcium carbonate to form Ca²⁺ and CO²⁻
₃ (carbonate). The carbonate ion has the potential to absorb two hydrogen ions. Therefore, it causes a net increase in ocean alkalinity. Calcium carbonate dissolution occurs in regions of the ocean which are undersaturated with respect to calcium carbonate.

The increasing carbon dioxide level in the atmosphere, due to carbon dioxide emissions, results in increasing absorption of CO₂ from the atmosphere into the oceans.[13] This does not affect the ocean's alkalinity[14]^: 2252 but it does result in a reduction in pH value (called ocean acidification).[15] Ocean alkalinity enhancement has been proposed as one option to add alkalinity to the ocean and therefore buffer against pH changes.[16]^: 181

Biological processes have a much greater impact on oceanic alkalinity on short (minutes to centuries) timescales.[17] Aerobic respiration of organic matter can decrease alkalinity by releasing protons.[17] Denitrification and sulfate reduction occur in oxygen-limited environments. Both of these processes consume hydrogen ions (thus increasing alkalinity) and release gases (N₂ or H₂S), which eventually escape into the atmosphere. Nitrification and sulfide oxidation both decrease alkalinity by releasing protons as a byproduct of oxidation reactions.[18]

• Alkali soils
• Base (chemistry)
• Biological pump
• Dealkalization of water
• Global Ocean Data Analysis Project

• 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-74.