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Sea surface alkalinity (from the GLODAP climatology).

Alkalinity or A_T measures the ability of a solution to neutralize acids to the equivalence point of carbonate or bicarbonate. 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 swimming 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).

Theoretical treatment of alkalinity[link]

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

A_T = [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:

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 alkalinity, as they act as sources of protons. They are often represented collectively as [H⁺]_T.

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

Example problems[link]

Sum of contributing species[link]

The following equations demonstrate the relative contributions of each component to the alkalinity of a typical seawater sample. Contributions are in μmol^.kg−soln^-1 and are obtained from A Handbook of Methods for the analysis of carbon dioxide parameters in seawater "[1],"(Salinity = 35, pH = 8.1, Temperature = 25°C).

A_T = [HCO₃⁻]_T + 2[CO₃⁻²]_T + [B(OH)₄⁻]_T + [OH⁻]_T + 2[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,

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

A_T = 1830 + 2*270 + 100 + 10 − 0.01

A_T = 2480 μmol^.kg−soln^-1

[edit] 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 CO₃^2-).

At neutral pH values:

CO₂ + H₂O → HCO₃⁻ + H⁺

At high pH values:

CO₂ + H₂O → CO₃^2- + 2H⁺

Dissolution of carbonate rock[link]

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.

Oceanic Alkalinity[link]

Processes that increase alkalinity[link]

There are many methods of alkalinity generation in the ocean. Perhaps the most well known is the dissolution of CaCO₃ (calcium carbonate, which is a component of coral reefs) to form Ca²⁺ and CO₃^2- (carbonate). The carbonate ion has the potential to absorb two hydrogen ions. Therefore, it causes a net increase in ocean alkalinity. Calcium carbonate dissolution is an indirect result of ocean acidification. It can cause great damage to coral reef ecosystems, but seems to have a relatively low effect on the total alkalinity (A_T) in the ocean. ^[1]

Anaerobic degradation processes, such as denitrification and sulfate reduction, have a much greater impact on oceanic alkalinity. Denitrification and sulfate reduction occur in the deep ocean, where there is an absence of oxygen. Both of these processes consume hydrogen ions and releases quasi-inert gases (N₂ or H₂S), which eventually escape into the atmosphere. This consumption of H⁺ increases the alkalinity. It has been estimated that anaerobic degradation could be as much as 60% of the total oceanic alkalinity. ^[2]

Processes that decrease alkalinity[link]

Anaerobic processes generally increase alkalinity. Conversely, aerobic degradation can decrease A_T. This process occurs in portions of the ocean where oxygen is present (surface waters). It results in dissolved organic matter and the production of hydrogen ions. ^[3] An increase in H⁺ clearly decreases alkalinity. However, the dissolved organic matter may have base functional groups that can consume these hydrogen ions and negate their effect on alkalinity. Therefore, aerobic degradation has a relatively low impact on the overall oceanic alkalinity. ^[4]

All of these aforementioned methods are chemical processes. However, physical processes can also serve to affect A_T. The melting of polar ice caps is a growing concern that can serve to decrease oceanic alkalinity. If the ice were to melt, then the overall volume of the ocean would increase. Because alkalinity is a concentration value (mol/L), increasing the volume would theoretically serve to decrease A_T. However, the actual effect would be much more complicated than this. ^[5]

Global temporal variability[link]

Researchers have shown oceanic alkalinity to vary over time. Because A_T is calculated from the ions in the ocean, a change in the chemical composition would alter alkalinity. One way this can occur is through ocean acidification. However, oceanic alkalinity is relatively stable, so significant changes can only occur over long time scales (i.e. hundreds to thousands of years). ^[6] As a result, seasonal and annual variability is generally very low. ^[7]

Spatial variability[link]

Researchers have also shown alkalinity to vary depending on location. Local A_T can be affected by two main mixing patterns: current and river. Current dominated mixing occurs close to the shore in areas with strong water flow. In these areas, alkalinity trends follow current and have a segmented relationship with salinity. ^[8]

River dominated mixing also occurs close to the shore; it is strongest close to the mouth of a large river (i.e. the Mississippi or Amazon). Here, the rivers can act as either a source or a sink of alkalinity. A_T follows the outflow of the river and has a linear relationship with salinity. This mixing pattern is most important in late winter and spring, because snow melt increases the river’s outflow. As the season progresses into summer, river processes are less significant, and current mixing can become the dominant process. ^[9]

Oceanic alkalinity also follows general trends based on latitude and depth. It has been shown that A_T is often inversely proportional to sea surface temperature (SST). Therefore, it generally increases with high latitudes and depths. As a result, upwelling areas (where water from the deep ocean is pushed to the surface) also have higher alkalinity values. ^[10]

Measurement Data Sets[link]

Throughout recent history, there have been many attempts to measure, record, and study oceanic alkalinity. Some of the larger data sets are listed below.

GEOSECS (Geochemical Ocean Sections Study)

TTO/NAS (Transient Tracers in the Ocean/North Atlantic Study)

JGOFS (Joint Global Ocean Flux Study)

WOCE (World Ocean Circulation Experiment)

CARINA (Carbon dioxide in the Atlantic Ocean)

References[link]

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See also[link]

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

Carbonate system calculators[link]

The following packages calculate the state of the carbonate system in seawater (including pH):

• CO2SYS, available as a stand-alone executable, Excel spreadsheet, or MATLAB script.
• seacarb, a R package for Windows, Mac OS X and Linux (also available here)
• CSYS, a Matlab script

External links[link]

• Holmes-Farley, Randy. "Chemistry and the Aquarium: What is Alkalinity?," Advanced Aquarist's Online Magazine. Alkalinity as it pertains to salt-water aquariums.
• DOE (1994) "[2],"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.
• GEOSECS data set [3]
• JGOFS data set [4]
• WOCE data set [5]
• CARINA data set [6]