How is seawater alkaline?

If when water and carbon dioxide react they form carbonic acid, carbonate, and bicarbonate, how does seawater still have a pH of around 8? Doesn't a compound need an hydroxide ion to be a base?




The short answer to the title question is *the ubiquitous presence of geologic and biogenic calcium carbonate $(\mathrm{p}K\_\mathrm{a} = 9)$*. The oceans as a whole can largely be thought of as residing on beds of calcium carbonate. The following are some excerpts from [this Wikipedia page](https://en.wikipedia.org/wiki/Calcium_carbonate):



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> Eggshells, snail shells and most seashells are predominantly calcium carbonate.
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> Carbonate is found frequently in geologic settings and constitute an enormous carbon reservoir. Calcium carbonate occurs as aragonite, calcite and dolomite. The carbonate minerals form the rock types: limestone, chalk, marble, travertine, tufa, and others.
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> Calcium carbonate contributors, including plankton (such as coccoliths and planktic foraminifera), coralline algae, sponges, brachiopods, echinoderms, bryozoa and mollusks, are typically found in shallow water environments where sunlight and filterable food are more abundant. Cold-water carbonates do exist at higher latitudes but have a very slow growth rate.
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Regarding your question



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> Doesn't a compound need an hydroxide ion to be a base?
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Here you go:


$$\ce{CO3^2- (aq) + H2O (l) <=> HCO3- (aq) + OH-(aq)}$$


Although calcium carbonate is only sparingly soluble $(K\_\mathrm{sp} = 3.3×10^{−9})$, as stated above the oceans are simply rife with it.


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Fellow collegaes,


The Question: how and why is seawater alkaline? Is very important.


Forget what you have ever learned and read " how to understand acid-base" by Peter Stewart.


Trying to break down his message:
1. First Consider pure H2O (without minerals) and it relation with temperature.
PH of pure H2O changes from 7.00 at 25 degrees celcius to 7.47 at 0 degrees celcius. H2O starts dissociating less as temperature goes down. From 25 to 0 degrees both [H+] and [OH] decrease and pH goes up, although the water remains neutral ([H+]=[OH-])


2. Second: the only independent factors (apart from temperature) capable of changing pH are: 1. Strong ion difference (alkalising) and 2. PCO2 (acidifying)


Consider all cations and anions in seawater.
[Na+]+[Mg]+[Ca2+]+[K+]+[rest cations]+[H+] = [Cl]+[SO4]+[HCO3]+[rest anions]+[OH-]
- All cations are "strong" (meaning fully dissociated)
- Not all anions are "strong" but a minority is "weak" meaning for a small part in equilibrium with [OH-], for instance:
1. HCO3- + H2O <-> H2CO3 + OH- or
2. HPO4- + H2O <-> H2PO4 + OH-.
H2CO3 and H2PO4 are unable to give of H+ due to a positive SID.


C/ the Strong Ion Difference between cations and anions (=SID) is what makes the ocean alkaline.


CO2 acidicifies but is unable to counter the SID


Consider:
1.If H2O evaporates: SID increases and pH goes up If H2O is added due to rains: SID decreases and pH goes down
2. Stewart demonstrated that a positive SID is the most potent buffer system and determines all because electroneutrality determines all. Formation of CaCO3 is only possible due to a positive SID. Any reasoning the other way around (e.q that CaCO3 could buffer the ocean) is wrong.


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The ultimate source of alkalinity in the oceans is the basic character of many of the rocks on Earth. As explained [here](https://earthscience.stackexchange.com/a/25266/20607), many rocks contain alkali-metal and alkaline-earth-metal-bearing silicates, in which the silicate ions are basic thus convert carbon dioxide to bicarbonate. In the case of the ocean the latter species is the predominant (weakly) basic ion in ocean waters. Alkaline-earth-metal silicates are especially important in the ocean environment because much of the rock in oceanic crust is basalt, which contains predominantly mafic minerals such as calcium-rich plagioclase feldspar and pyroxenes (which are rich in both calcium and magnesium).


The minerals in basalt are also susceptible to forming basic hydroxides at high temperature even in the absence of carbon dioxide. [Serpentization](https://en.wikipedia.org/wiki/Serpentinization), a high-temperature between magnesium- and iron-rich silicates and water, occurs in hot spots including the geologically active mid-ocean ridges. The overall chemistry of this reaction is complex, but one of the major results is that the magnesium-bearing components in the rock produce basic products such as brucite (magnesium hydroxide). Where calcium is present in the rock, it woukd similarly produce calcium hydroxide; this dissolves and reacts further as a base (Wikipedia gives dissolved calcium ions as the product in such cases).


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The chemistry is complex given the organic content of sea water and the influence of sunlight. For example, [here is an article](https://pubs.acs.org/doi/full/10.1021/acsomega.7b00453) detailing the influence of sunlight on oceans containing dissolved organic matter, to quote:



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> Solar radiation mineralizes dissolved organic matter (DOM) to dissolved inorganic carbon through photochemical reactions (DIC photoproduction) that are influenced by iron (Fe) and pH [...] Fe raised the rate of photobleaching and steepened the spectral slopes of CDOM in low pH but resisted the slope steepening in neutral to alkaline pH.
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where apparently acidic pH is much more effect implying an asymmetric effect with the consumption of H+. 


Similarly, other reactions involving iron or other transition metals consume H+ also:


2 Mn(II)/Fe(II) + 1/2 O2 + H+ --> 2 Mn(III)/Fe(III) + OH- 


Bottom line, asymmetric pH reactions can in heterogeneous systems found in large bodies of water, rich in organic matter and subject to irradiation, can result in a pH shift, including into the alkaline range.


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