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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?

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    $\begingroup$ The chemistry of the oceans make for a pretty complex situation. But for starters, think about what could be in the ocean and the geology of the materials "containing" the ocean. The adsorption of carbon dioxide of course affects the pH, but it is far from the only process at play here. What kind of minerals could dissolve to give a pH in the 8.3 range? $\endgroup$ – airhuff Feb 22 '17 at 2:51
  • $\begingroup$ It's true that I made a questionable assumption, but as I see it, none of the common salts in the sea could've affected the pH, so I would think you would need to dissolve some atmospheric gases in the ocean, or find some organic process to explain the alkalinity. $\endgroup$ – smaude Feb 22 '17 at 2:56
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    $\begingroup$ How about carbonate mineralogy and carbonate based sea shells? You are right that carbon dioxide is an acidic gas, and as the atmospheric concentration increases, more of it dissolves into the oceans, slightly increasing their acidity, but also dissolving more carbonates, acting somewhat as a buffer system. $\endgroup$ – airhuff Feb 22 '17 at 3:14
  • $\begingroup$ yeah, I guess that would be the thing. I'd just really like to figure out how it works. $\endgroup$ – smaude Feb 22 '17 at 3:26
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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:

Eggshells, snail shells and most seashells are predominantly calcium carbonate.

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.

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.

Regarding your question

Doesn't a compound need an hydroxide ion to be a base?

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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  • $\begingroup$ Also, sea water is less than $25^o C$, so the $k_w$ is less in sea water; therefore, the $H^+$ concentration is also lower, leading to a higher pH. $\endgroup$ – Cyclopropane Jan 22 '19 at 12:18
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    $\begingroup$ @DrPepper You are not serious, are you? $\endgroup$ – andselisk Jan 22 '19 at 12:20
  • $\begingroup$ @andselisk why do you think i'm joking? $\endgroup$ – Cyclopropane Jan 22 '19 at 14:42
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    $\begingroup$ @DrPepper First, temperature of the sea water also exceeds 25 °C in tropical areas, which, following your logic, should result in slightly acidic water near equator, which is not true. Second, yes, $K_w$ depends on temperature, but it doesn't depend on concentrations: with more released protons there is going to be an equal amount of hydroxide-ions in the solution; formally, you get a slightly higher pH, but also the same increase in pOH and it doesn't affect alkalinity at all (as the title mentions it). $\endgroup$ – andselisk Jan 22 '19 at 14:51
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    $\begingroup$ @DrPepper Third, there is no way to increase the formal pH value by lowering the temperature to even merely reach the pH 8 mentioned in the question. Even at 0 °C pH = pOH = 7.47! $\endgroup$ – andselisk Jan 22 '19 at 14:52
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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-])

  1. 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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    $\begingroup$ Strictly speaking, magnesium ion does show some acidity due to hydrolysis and calcium ion, in the presence of bicarbonate, can consume hydroxide ions via $\ce{Ca^{2+} + HCO3- + OH- <=> CaCO3(s) + H2O}$. But the former is much less than the alkaline bicarbonate hydrolysis and the latter requires the seawater to become alkaline first, so the basic reasoning passes muster. $\endgroup$ – Oscar Lanzi Jan 22 '19 at 14:16
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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 detailing the influence of sunlight on oceans containing dissolved organic matter, to quote:

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.

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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to advance the discussion: consider the pH of water is temperature dependent enter image description here

for pure H2O (without any minerals): 1.[H+]=[OH-] (always neutral) and 2.from 0 degrees to 60 degrees [H+] and [OH-] increase simultaneously. The pH of pure and neutral water drops from 7.47 to 6.14 from 0 to 100 degrees. For discussion purpose lets assume the standard temperature of 25 degrees (pH=7.0).

Doesn't a compound need an hydroxide ion to be a base? First realize pH (and pOH) concerns a compound in water; second, every aquatic/H2O solution has both [OH-] and [H+], therefore Stewart (stewart textbook of acid-base, 2.2 definitions) simply defines: a base as a solution wherein [OH-]>[H+].

Due to the salts in solution: [H+] > [OH-] (acidic) or [OH-] > [H+] (alkaline).

Salts quantatively determine [H+] and [OH-] because 1. electroneutrality is a general principle ([cations]=[anions]) and 2. concentrations of salts (order of 600 meq/l = 6e102) far outweigh [H+] and [OH]- (order of neq/l =10 -9) and 3. [H+]*[OH-] is Kw.

CO2 pushes an equilibrium (CO2 <-> H2CO3 <-> H+ and HCO3-) already determined by cations (Na+, Mg2+, Ca2+, K+, etc) and anions (Cl-, SO4 2-, HCO3-, etc).

the reason seawater is basic is because HCO3- is pushed into equilibrilium with CO2 (aq) and OH- by the cations.

for example a solution with 600 mmol NaCl and 2 mmol NaHCO3 (calculated with Aqion online Ph calculator). Mainly the dissociation of HCO3- into CO2 and OH- contributes to pH 7.71. If the same composition is analysed at 4 degrees the pH is 8.07 (calculated with phreeq online pH calculator).

enter image description here

the aquous species are:

enter image description here

In seawater the strong cations sodium, magnesium, calcium and potassium are all contributing to [OH-] being larger than [H+]

An example of cation and anion composition in mmol/l with equilibrium of monovalent and bivalent ions. (mM concentration is calculated for 20 degrees) Calculated miliequivalents are: total cations 620,56meq/l and total anions 620,305 meq/l.

enter image description here

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  • $\begingroup$ While this link may answer the question, it is better to include the essential parts of the answer here and provide the link for reference. Link-only answers can become invalid if the linked page changes. - From Review $\endgroup$ – A.K. Jan 22 '19 at 13:41
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    $\begingroup$ If you start with pure water, you can make it alkaline by adding a compound that contains hydrogen but does not contain oxygen (e.g. the base $\ce{NH3}$. This counter example demonstrates that alkalinity is not strictly related to the proportion of oxygen and hydrogen in a sample. To give you another counter example: Dissolving elemental oxygen in pure water does not make it alkaline, even though I now more than the 1:2 ratio of total oxygen atoms to total hydrogen atoms in the sample. $\endgroup$ – Karsten Theis Jan 22 '19 at 16:35
  • $\begingroup$ I have undeleted this because a large improvement was made, but I have no way to validate its content, so I'll leave that up to the community to decide. $\endgroup$ – jonsca Jan 28 '19 at 2:51
  • $\begingroup$ Oxygen is both the medium and more electronegative than nitrogen. Oxygen gladly donates one of its protons+ to NH3 leaving hydroxide-. Nitrogens concentration within seawater however is minute and existing as NH4(1+) and its oxidized forms NO2(1-) and NO3 (1-). Oxygen is far more prevelant than nitrogen both is the earths crust and in the ocean. $\endgroup$ – HW Willemsen Mar 22 at 10:58

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