# Alkalinity

Alkalinity (from Arabic: القلوي, romanizedal-qaly, lit.'ashes of the saltwort')[1] is the capacity of water to resist acidification.[2] It should not be confused with basicity, which is an absolute measurement on the pH scale.

Sea surface alkalinity (from the GLODAP climatology).

Alkalinity is the strength of a buffer solution composed of weak acids and their conjugate bases. It is measured by titrating the solution with an acid such as HCl until its pH changes abruptly, or it reaches a known endpoint where that happens. Alkalinity is expressed in units of concentration, such as meq/L (milliequivalents per liter), μeq/kg (microequivalents per kilogram), or mg/L CaCO3 (milligrams per liter of calcium carbonate).[3] Each of these measurements corresponds to an amount of acid added as a titrant.

Although alkalinity is primarily a term used by oceanographers,[3] it is also used by hydrologists to describe temporary hardness. Moreover, measuring alkalinity is important in determining a stream's ability to neutralize acidic pollution from rainfall or wastewater. It is one of the best measures of the sensitivity of the stream to acid inputs.[4] There can be long-term changes in the alkalinity of streams and rivers in response to human disturbances such as acid rain generated by SOx and NOx emissions.[5]

## History

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 Svante_Arrhenius#Ionic_disassociation.

## Simplified summary

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 equal zero.[6] 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.[6] 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.[7] 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.[6]

## Detailed description

Alkalinity or measures the ability of a solution to neutralize acids to the equivalence point of carbonate or bicarbonate, defined as pH 4.5 for oceanographic/limnological studies.[8] 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 presence of carbon dioxide in the atmosphere. Other common natural components that can contribute to alkalinity include borate, hydroxide, phosphate, silicate, dissolved ammonia, the conjugate bases of some organic acids (e.g., acetate), and sulfate. Solutions produced in a laboratory may contain a virtually limitless number of bases that contribute to alkalinity. Alkalinity is usually given as molar equivalents per liter or kilogram of solution. Commercially, as in the swimming pool industry, 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.[9]

## Theoretical treatment

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

AT = [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.[10] 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
2
PO
4
HPO2−
4
+ H+H
2
PO
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).

## Example problems

### Sum of contributing species

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] Archived 2011-10-25 at the Wayback Machine," salinity = 35 g/kg, pH = 8.1, temperature = 25 °C.

AT = [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+] − [HSO
4
] − [HF]

Phosphates and silicate, being nutrients, are typically negligible. At pH = 8.1, [HSO
4
] and [HF] are also negligible. So,

 AT = [HCO−3]T + 2[CO2−3]T + [B(OH)−4]T + [OH−]T − [H+] = 1830 + 2 × 270 + 100 + 10 − 0.01 = 2480 μmol.kg−soln−1

Addition (or removal) of CO2 to a solution does not change its alkalinity, since the net reaction produces the same number of equivalents of positively contributing species (H+) as negative contributing species (HCO
3
and/or CO2−
3
). Adding CO2 to the solution lowers its pH, but does not affect alkalinity.

At all pH values:

CO2 + H2O ⇌ HCO
3
+ H+

Only at high (basic) pH values:

HCO
3
+ H+CO2−
3
+ 2 H+

### Dissolution of carbonate rock

Addition of CO2 to a solution in contact with a solid can (over time) 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 Ca2+ and CO2−
3
into solution. Ca2+ will not influence alkalinity, but CO2−
3
will increase alkalinity by 2 units. Increased dissolution of carbonate rock by acidification from acid rain and mining has contributed to increased alkalinity concentrations in some major rivers throughout the eastern U.S.[5] The following reaction shows how acid rain, containing sulfuric acid, can have the effect of increasing river alkalinity by increasing the amount of bicarbonate ion:

2 CaCO3 + H2SO4 → 2 Ca2+ + 2 HCO
3
+ SO2−
4

Another way of writing this is:

CaCO3 + H+ ⇌ Ca2+ + HCO
3

The lower the pH, the higher the concentration of bicarbonate will be. This shows how a lower pH can lead to higher alkalinity if the amount of bicarbonate produced is greater than the amount of H+ remaining after the reaction. This is the case since the amount of acid in the rainwater is low. If this alkaline groundwater later comes into contact with the atmosphere, it can lose CO2, precipitate carbonate, and thereby become less alkaline again. When carbonate minerals, water, and the atmosphere are all in equilibrium, the reversible reaction

CaCO3 + 2 H+ ⇌ Ca2+ + CO2 + H2O

shows that pH will be related to calcium ion concentration, with lower pH going with higher calcium ion concentration. In this case, the higher the pH, the more bicarbonate and carbonate ion there will be, in contrast to the paradoxical situation described above, where one does not have equilibrium with the atmosphere.

## Oceanic alkalinity

### Processes that increase alkalinity

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 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 is an indirect result of ocean pH lowering. It can cause great damage to coral reef ecosystems, but has a relatively low effect on the total alkalinity (AT) in the ocean.[11] Lowering of pH due to absorption of CO2 actually raises the alkalinity by causing dissolution of carbonates.

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.[11]

### Processes that decrease alkalinity

Anaerobic processes generally increase alkalinity. Conversely, aerobic degradation can decrease AT. 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.[11] 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.[12]

All of these aforementioned methods are chemical processes. However, physical processes can also serve to affect AT. 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 AT. However, the actual effect would be much more complicated than this.[13]

### Global temporal variability

Researchers have shown oceanic alkalinity to vary over time. Because AT 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).[14] As a result, seasonal and annual variability is generally very low.[11]

### Spatial variability

Researchers have also shown alkalinity to vary depending on location. Local AT 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.[15]

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. AT 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 snowmelt increases the river's outflow. As the season progresses into summer, river processes are less significant, and current mixing can become the dominant process.[11]

Oceanic alkalinity also follows general trends based on latitude and depth. It has been shown that AT 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.[16]

### Measurement data sets

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

1. ^ "alkali". Dictionary.com Unabridged (Online). n.d. Retrieved 2018-09-30.
2. ^ "What is alkalinity?". Water Research Center. 2014. Retrieved 5 February 2018.
3. ^ a b Dickson, Andrew G. (1992). "The development of the alkalinity concept in marine chemistry". Marine Chemistry. 40 (1–2): 49–63. doi:10.1016/0304-4203(92)90047-E.
4. ^ "Total Alkalinity". United States Environment Protection Agency. Retrieved 6 March 2013.
5. ^ a b Kaushal, S. S.; Likens, G. E.; Utz, R. M.; Pace, M. L.; Grese, M.; Yepsen, M. (2013). "Increased river alkalinization in the Eastern U.S.". Environmental Science & Technology. 47 (18): 10302–10311. doi:10.1021/es401046s. PMID 23883395.
6. ^ a b c Drever, James I. (1988). The Geochemistry of Natural Waters, Second Edition. Englewood Cliffs, NJ: Prentice Hall. ISBN 0-13-351396-3.
7. ^ Wolf-Gladrow, Dieter A.; Zeebe, Richard E.; Klaas, Christine; Körtzinger, Arne; Dickson, Andrew G. (July 2007). "Total alkalinity: The explicit conservative expression and its application to biogeochemical processes". Marine Chemistry. 106 (1–2): 287–300. doi:10.1016/j.marchem.2007.01.006.
8. ^ Dickson, A.G. (June 1981). "An exact definition of total alkalinity and a procedure for the estimation of alkalinity and total inorganic carbon from titration data". Deep Sea Research Part A. Oceanographic Research Papers. 28 (6): 609–623. Bibcode:1981DSRA...28..609D. doi:10.1016/0198-0149(81)90121-7.
9. ^ "2320 alkalinity", Standard Methods For the Examination of Water and Wastewater, Standard Methods for the Examination of Water and Wastewater, American Public Health Association, 2017-08-27, doi:10.2105/smww.2882.023, retrieved 2022-12-01
10. ^ Benjamin. Mark M. 2015. Water Chemistry. 2nd Ed. Long Grove, Illinois: Waveland Press, Inc.
11. Thomas, H.; Schiettecatte, L.-S.; et al. Enhanced Ocean Carbon Storage from Anaerobic Alkalinity Generation in Coastal Sediments. Biogeosciences Discussions. 2008, 5, 3575–3591
12. ^ Kim, H.-C., and K. Lee (2009), Significant contribution of dissolved organic matter to seawater alkalinity, Geophys. Res. Lett., 36, L20603, doi:10.1029/2009GL040271
13. ^ Chen, B.; Cai, W. Using Alkalinity to Separate the Inputs of Ice-Melting and River in the Western Arctic Ocean. Proceedings from the 2010 AGU Ocean Sciences Meeting, 2010, 22-26.
14. ^ Doney, S. C.; Fabry, V. J.; et al. Ocean Acidification: The Other CO2 Problem. Annu. Rev. Mar. Sci., 2009, 69-92. doi:10.1146/annurev.marine.010908.163834
15. ^ Cai, W.-J.; Hu, X. et al. Alkalinity Distribution in the Western North Atlantic Ocean Margins. Journal of Geophysical Research. 2010, 115, 1-15. doi:10.1029/2009JC005482
16. ^ Millero, F. J.; Lee, K.; Roche, M. Distribution of alkalinity in the surface waters of the major oceans. Marine Chemistry. 1998, 60, 111-130.