Alkalinity

Not to be confused withalkaliorbase (chemistry).

Sea surface alkalinity (from theGLODAPclimatology).

Alkalinity(fromArabic:القلوية,romanized:al-qaly,lit. 'ashes of thesaltwort')is the capacity ofwaterto resistacidification.It should not be confused withbasicity,which is an absolute measurement on thepHscale. Alkalinity is the strength of abuffer solutioncomposed ofweak acidsand theirconjugate bases.It is measured bytitratingthesolutionwith an acid such asHCluntil its pH changes abruptly, or it reaches a known endpoint where that happens. Alkalinity is expressed in units of concentration, such as meq/L (milliequivalentsperliter), μeq/kg (microequivalents per kilogram), or mg/L CaCO₃(milligrams per liter ofcalcium carbonate).Each of these measurements corresponds to an amount of acid added as atitrant.

Infreshwater,particularly those on non-limestoneterrains, alkalinities are low and involve a lot of ions. In the ocean, on the other hand, alkalinity is completely dominated bycarbonateandbicarbonateplus a small contribution fromborate.

Although alkalinity is primarily a term used bylimnologistsandoceanographers,it is also used byhydrologiststo describetemporary hardness.Moreover, measuring alkalinity is important in determining a stream's ability to neutralize acidic pollution fromrainfallorwastewater.It is one of the best measures of the sensitivity of the stream to acid inputs.There can be long-term changes in the alkalinity of streams and rivers in response to human disturbances such as acid rain generated by SO_xand NO_xemissions.

History

In 1884, ProfessorWilhelm (William) Dittmarof Anderson College, now theUniversity of Strathclyde,analysed 77 pristine seawater samples from around the world brought back by theChallenger expedition.He found that in seawater the major ions were in a fixed ratio, confirming the hypothesis ofJohan 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 Arrheniussubmitted his PhD theses in which he advocated the existence ofionsin solution, and defined acids ashydroniumion donors and bases ashydroxideion donors. For that work, he received theNobel Prize in Chemistryin 1903.See alsoSvante 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 ofHCO⁻
₃in solution represents 1 molar equivalent, while 1 mole ofCO²⁻
₃is 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.This leads to a parallel definition of alkalinity that is based upon the charge balance of ions in a solution.

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

Certain ions, including Na⁺,K⁺,Ca²⁺,Mg²⁺,Cl⁻,SO²⁻
₄,andNO⁻
₃are "conservative"such that they are unaffected by changes in temperature, pressure or pH.Others such asHCO⁻
₃are 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.

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

This combined charge balance and proton balance is calledtotal alkalinity.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 exceptHCO⁻
₃andCO²⁻
₃have low concentrations in Earth's surface water (streams, rivers, and lakes). Thuscarbonate alkalinity,which is equal to[HCO⁻
₃] + 2[CO²⁻
₃]is also approximately equal to the total alkalinity in surface water.

Detailed description

Alkalinity measures the ability of a solution to neutralizeacidsto theequivalence pointof carbonate or bicarbonate, defined as pH 4.5 for many oceanographic/limnological studies.The alkalinity is equal to thestoichiometricsum of thebasesin solution. In most Earth surface waterscarbonate alkalinitytends to make up most of the total alkalinity due to the common occurrence and dissolution ofcarbonaterocks and other geologicalweatheringprocesses that produce carbonate anions. Other common natural components that can contribute to alkalinity includeborate,hydroxide,phosphate,silicate,dissolvedammonia,and theconjugate basesoforganic 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 withbasicity.For example, the addition of CO₂lowers the pH of a solution, thus reducing basicity while alkalinity remains unchanged (see example below).

A variety oftitrants,endpoints, andindicatorsare specified for various alkalinity measurement methods.Hydrochloricandsulfuricacids are common acid titrants, whilephenolpthalein,methyl red,andbromocresol greenare common indicators.

Theoretical treatment

In typicalgroundwaterorseawater,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 ofion pairinteractions 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 CO₂equivalence point where the major component in water is dissolved CO₂which is converted to H₂CO₃in an aqueous solution. There are no strong acids or bases at this point. Therefore, the alkalinity is modeled and quantified with respect to the CO₂equivalence point. Because the alkalinity is measured with respect to the CO₂equivalence 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 CO₂equivalence point. The dissolved base in water increases the pH and titrates an equivalent amount of CO₂to 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 CO₂and continuous addition of strong acids can cause the alkalinity to become less than zero.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⁻
₃however, will consume two protons before becoming a zero-level species (CO₂), thus it increases alkalinity by two per mole ofCO²⁻
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/LasCaCO₃.(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 approximateMWof CaCO₃divided by 2).

Carbon dioxide interactions

Addition of CO₂

Addition (or removal) of CO₂to 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⁻
₃and/orCO²⁻
₃). Adding CO₂to 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 CO₂to 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 CaCO₃and its dissociation will add Ca²⁺andCO²⁻
₃into solution. Ca²⁺will not influence alkalinity, butCO²⁻
₃will 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.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⁺+ 2HCO⁻
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.

Changes to oceanic alkalinity

See also:Ocean § Alkalinity,pH § pH in the ocean,Ocean acidification,andEffects of climate change on oceans

In the ocean, alkalinity is completely dominated bycarbonateandbicarbonateplus a small contribution fromborate.

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 theocean.Perhaps the most well known is the dissolution of calcium carbonate to form Ca²⁺andCO²⁻
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 increasingcarbon dioxide level in the atmosphere,due tocarbon dioxide emissions,results in increasing absorption of CO₂from the atmosphere into the oceans.This does not affect the ocean's alkalinitybut it does result in a reduction in pH value (calledocean acidification).Ocean alkalinity enhancementhas been proposed as one option to add alkalinity to the ocean and therefore buffer against pH changes.

Biological processes have a much greater impact on oceanic alkalinity on short (minutes to centuries) timescales.Aerobic respirationof organic matter can decrease alkalinity by releasing protons.Denitrificationandsulfate reductionoccur in oxygen-limited environments. Both of these processes consume hydrogen ions (thus increasing alkalinity) and release gases (N₂or H₂S), which eventually escape into the atmosphere.Nitrificationandsulfide oxidationboth decrease alkalinity by releasing protons as a byproduct of oxidation reactions.

Global temporal and spatial variability

The ocean's alkalinity varies over time, most significantly over geologic timescales (millennia). Changes in the balance between terrestrialweatheringand sedimentation of carbonate minerals (for example, as a function of ocean acidification) are the primary long-term drivers of alkalinity in the ocean.Over human timescales, mean ocean alkalinity is relatively stable.Seasonal and annual variability of mean ocean alkalinity is very low.

Alkalinity varies by location depending on evaporation/precipitation, advection of water, biological processes, and geochemical processes.

River dominated mixing also occurs close to the shore; it is strongest close to the mouth of a large river. Here, the rivers can act as either a source or a sink of alkalinity. A_Tfollows the outflow of the river and has a linear relationship with salinity.

Oceanic alkalinity also follows general trends based on latitude and depth. It has been shown that A_Tis often inversely proportional tosea 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.

There are many programs to measure, record, and study oceanic alkalinity, together with many of the other characteristics of seawater, like temperature and salinity. These include: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).

Interventions to add alkalinity

This section is an excerpt fromOcean acidification § Carbon removal technologies which add alkalinity.[edit]

Somecarbon removaltechniques add alkalinity to the ocean and therefore immediately buffer pH changes which might help the organisms in the region that the extra alkalinity is added to. The two technologies that fall into this category are ocean alkalinity enhancement and electrochemical methods.Eventually, due to diffusion, that alkalinity addition will be quite small to distant waters. This is why the termlocal ocean acidification mitigationis used. Both of these technologies have the potential to operate on a large scale and to be efficient at removing carbon dioxide.However, they are expensive, have many risks and side effects and currently have a lowtechnology readiness level.

See also

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

External links

• DOE (1994) "[1],"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.

Carbonate system calculators

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

• CO2SYSArchived2011-10-14 at theWayback Machine,available as a stand-aloneexecutable,Excelspreadsheet, orMATLABscript.
• seacarb,aR packageforWindows,Mac OS XandLinux(also availablehere)
• CSYS,aMatlab script