<p>Solubility. The question is quite chemical, in fact.</p>
<p>Solubility works somewhat counter-intuitively: if two ions that are not soluble together are present in the solution, they find each other and form a solid. You can mix e.g. sodium carbonate (soluble) and calcium chloride (soluble) and get calcium carbonate (insoluble).</p>
<p>This is why oceans don't contain much of calcium - they contain quite a deal of carbonate.</p>
<p>If one makes the ocean more acidic, the carbonate ion will be much less and we could indeed have calcium in the solution.</p>
<p>We can't have much iron either - iron ions are insoluble because they precipitate with hydroxide ions and you need quite an acidic solution in order to have barely soluble iron.</p>
<p>The same phenomenon explains the weathering - the rainwater and aquifiers wash away ions depending on their own initial ion content (including acidity).</p>
<p>What's important with NaCl? Both its ions are quite soluble in the presence of almost any other ion. This is why they are easy to leach from the stones and rocks and in fact we are almost done with this. This is why rivers don't bring much salt into the oceans these days - there is no much of them left.</p>
<p>P.s. in regard to iron - long ago, before the first bacteria learned to produce oxygen, the ocean was full of iron. It was Fe<sup>2+</sup> - this is the form stable when no oxygen is present. Fe<sup>2+</sup> is much more soluble than the more oxidized form (Fe<sup>3+</sup>). This is how some large iron ore deposits formed.</p>

CC BY-SA 4.0
2023-01-25T14:30:35.360Z
+1 This answer gives some intuition about how solubility works and an example of how it depends on the ambient conditions (Fe2+ in the oceans of the early Earth).
WaterMolecule
CC BY-SA 4.0
2023-01-25T15:37:23.713Z
This is an amazingly informative answer, thanks so much. I wish I could accept both your answer and the one from David Hammen.
Fraser Orr
CC BY-SA 4.0
2023-01-25T17:24:44.233Z
There is a YouTube channel devoted to refining precious metals. The guy who runs the channel will test a sample for silver in solution by putting a drop of HCl in it. The results are instant and unmistakable when silver is present.
EvilSnack
CC BY-SA 4.0
2023-01-25T22:48:08.003Z
Although making the ocean more acidic would in general reduce CO3-- concentration, combination of CO2 and H2O to form H2CO3 (carbonic acid) would have the opposite effect.
supercat
CC BY-SA 4.0
2023-01-26T06:38:54.443Z
@supercat extra CO2 works to dissolve CaCO3 exactly because of its acidity. This is how some caves are formed.
fraxinus
CC BY-SA 4.0
2023-01-26T12:16:08.093Z
Great answer. The upshot is that it would be more useful to compare not solubilities of salts, but rather ... _precipitabilities(?)_ of ions.
leftaroundabout
CC BY-SA 4.0
2023-01-26T18:09:33.463Z
@fraxinus: Could a weak H2CO3 solution dissolve CaCO3 in the absence of other ions to associate with the Ca++ and CO3--?  I can see how exposing CaCO3 to e.g. acetic acid could result in a single-exchange reaction forming soluble calcium acetate and carbonic acid, but such a "reaction" between calcium carbonate with carbonic acid would yield calcium carbonate and carbonic acid.
supercat
CC BY-SA 4.0
2023-01-26T19:21:47.823Z
@supercat pretty much yes. https://en.wikipedia.org/wiki/Karst
fraxinus
CC BY-SA 4.0
2023-01-26T21:53:09.080Z
[Potassium](https://en.wikipedia.org/wiki/Potassium) is [not that different from sodium](https://en.wikipedia.org/wiki/Potassium_chloride). Does it have a solubility heel? Is sodium stealing chlorine from it?
Peter Mortensen
CC BY-SA 4.0
2023-01-27T00:20:03.143Z
@PeterMortensen I suspect this is because potassium feldspars are more stable than sodium feldspars. I found some papers that talk about ion exchange between KCl and albite (NaAlSi₃O₈), but not sure I understand this right. https://www.sciencedirect.com/science/article/abs/pii/S0892687518303984
leftaroundabout
CC BY-SA 4.0
2023-01-27T08:15:58.903Z
Oceans contain a great deal of Potassium as well. No, it does not "steal" anything. When dissolved, these ions are fully independent from their negative counterparts as long as they don't form an insoluble substance. And Potassium itself does not, just like Sodium. What makes Potassium less prominent in oceans and easilly found in soils is its higher ability to "stick" to clays - related to @leftaroundabout comment. This ability does not prevent water from washing it down, but the process is slower than with Sodium.
fraxinus
CC BY-SA 4.0
2023-01-27T17:26:25.200Z
As they say, if you're not part of the solution, you're part of the precipitate.
Robert Columbia
CC BY-SA 4.0
2023-01-25T14:23:45.983Z
2023-01-25T20:30:47.857Z
2023-01-25T20:30:47.857Z
fraxinus
<p>Fluoride salts tend to be not particularly soluble in water. Chloride salts are. The same goes for salts containing sodium versus those containing calcium. Sodium chloride is ridiculously easy to dissolve.</p>
<p>Regarding your second question, it is geological forces that keep salinity more or less constant. People formerly argued that the Earth can't be more than a few hundred million years old because otherwise the river waters running into the oceans would eventually result in an insanely high salinity. It turns out that the Earth's oceans are young (young compared to the 4.5 billion year age of the Earth). The vast majority of oceanic crust is less than 100 million years old. We see huge salt deposits sprinkled across the world because those are the dried up remnants of former seas and oceans. Salt is also drawn into the Earth at subduction zones, where it combines chemically with basalt.</p>

CC BY-SA 4.0
2023-01-25T15:36:19.687Z
Thanks for the great answer -- I had a look and apparently NaCl is twenty times more soluble in water than, for example, KCl.
Fraser Orr
CC BY-SA 4.0
2023-01-25T05:01:11.193Z
2023-01-25T05:01:11.193Z
David Hammen
<p>Why is sodium chloride far and away the most abundant salt dissolved in ocean water? Its two constituent ions do have a very high frequency in the crust of the earth, but they are far from the most common. Chlorine is (according to <a href="https://en.wikipedia.org/wiki/Chlorine" rel="noreferrer">Wikipedia</a>) the 21st most abundant element, and <a href="https://en.wikipedia.org/wiki/Sodium" rel="noreferrer">sodium</a> 6th.</p>
<p>I certainly understand that a combination of their solubility and reasonably high frequency would lead one to expect them to be abundant in sea water, but they are hyper abundant, completely dominating all other salt ions. Iron, for example, is twice as abundant, and potassium only a little less abundant, and fluorine more abundant than chlorine.</p>
<p>Moreover, if the salts are deposited in the ocean through weathering of rocks and deposition via rivers, why does the salinity not simply grow and grow? I understand that some is lost due to tectonic activity, but it seems extraordinarily unlikely that these two forces should be equally balanced, and so we would see a significant change in average salinity over time.</p>
<p>(Please note I am migrating this question from the Chemistry SE at their recommendation.)</p>

<oceanography><geochemistry>
Why is NaCl so hyper abundant in the ocean?
Fraser Orr
# Why is NaCl so hyper abundant in the ocean?

(Asked by: username_0 on 2023-01-25)

Why is sodium chloride far and away the most abundant salt dissolved in ocean water? Its two constituent ions do have a very high frequency in the crust of the earth, but they are far from the most common. Chlorine is (according to [Wikipedia](https://en.wikipedia.org/wiki/Chlorine)) the 21st most abundant element, and [sodium](https://en.wikipedia.org/wiki/Sodium) 6th.

I certainly understand that a combination of their solubility and reasonably high frequency would lead one to expect them to be abundant in sea water, but they are hyper abundant, completely dominating all other salt ions. Iron, for example, is twice as abundant, and potassium only a little less abundant, and fluorine more abundant than chlorine.

Moreover, if the salts are deposited in the ocean through weathering of rocks and deposition via rivers, why does the salinity not simply grow and grow? I understand that some is lost due to tectonic activity, but it seems extraordinarily unlikely that these two forces should be equally balanced, and so we would see a significant change in average salinity over time.

(Please note I am migrating this question from the Chemistry SE at their recommendation.)

**Tags:** oceanography, geochemistry


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## 2 answers

### Answer by username_8 on 2023-01-25. Score: 50 | Accepted answer

Fluoride salts tend to be not particularly soluble in water. Chloride salts are. The same goes for salts containing sodium versus those containing calcium. Sodium chloride is ridiculously easy to dissolve.

Regarding your second question, it is geological forces that keep salinity more or less constant. People formerly argued that the Earth can't be more than a few hundred million years old because otherwise the river waters running into the oceans would eventually result in an insanely high salinity. It turns out that the Earth's oceans are young (young compared to the 4.5 billion year age of the Earth). The vast majority of oceanic crust is less than 100 million years old. We see huge salt deposits sprinkled across the world because those are the dried up remnants of former seas and oceans. Salt is also drawn into the Earth at subduction zones, where it combines chemically with basalt.

> **Comments:**
> **username_0**: Thanks for the great answer -- I had a look and apparently NaCl is twenty times more soluble in water than, for example, KCl.  

---

### Answer by username_1 on 2023-01-25. Score: 44

Solubility. The question is quite chemical, in fact.

Solubility works somewhat counter-intuitively: if two ions that are not soluble together are present in the solution, they find each other and form a solid. You can mix e.g. sodium carbonate (soluble) and calcium chloride (soluble) and get calcium carbonate (insoluble).

This is why oceans don't contain much of calcium - they contain quite a deal of carbonate.

If one makes the ocean more acidic, the carbonate ion will be much less and we could indeed have calcium in the solution.

We can't have much iron either - iron ions are insoluble because they precipitate with hydroxide ions and you need quite an acidic solution in order to have barely soluble iron.

The same phenomenon explains the weathering - the rainwater and aquifiers wash away ions depending on their own initial ion content (including acidity).

What's important with NaCl? Both its ions are quite soluble in the presence of almost any other ion. This is why they are easy to leach from the stones and rocks and in fact we are almost done with this. This is why rivers don't bring much salt into the oceans these days - there is no much of them left.

P.s. in regard to iron - long ago, before the first bacteria learned to produce oxygen, the ocean was full of iron. It was Fe2+ - this is the form stable when no oxygen is present. Fe2+ is much more soluble than the more oxidized form (Fe3+). This is how some large iron ore deposits formed.

> **Comments:**
> **username_2**: +1 This answer gives some intuition about how solubility works and an example of how it depends on the ambient conditions (Fe2+ in the oceans of the early Earth).  
> **username_0**: This is an amazingly informative answer, thanks so much. I wish I could accept both your answer and the one from username_8 username_8.  
> **username_3**: There is a YouTube channel devoted to refining precious metals. The guy who runs the channel will test a sample for silver in solution by putting a drop of HCl in it. The results are instant and unmistakable when silver is present.  
> **username_4**: Although making the ocean more acidic would in general reduce CO3-- concentration, combination of CO2 and H2O to form H2CO3 (carbonic acid) would have the opposite effect.  
> **username_1**: @username_4 extra CO2 works to dissolve CaCO3 exactly because of its acidity. This is how some caves are formed.  
> **username_5**: Great answer. The upshot is that it would be more useful to compare not solubilities of salts, but rather ... _precipitabilities(?)_ of ions.  
> **username_4**: @username_1: Could a weak H2CO3 solution dissolve CaCO3 in the absence of other ions to associate with the Ca++ and CO3--?  I can see how exposing CaCO3 to e.g. acetic acid could result in a single-exchange reaction forming soluble calcium acetate and carbonic acid, but such a "reaction" between calcium carbonate with carbonic acid would yield calcium carbonate and carbonic acid.  
> **username_1**: @username_4 pretty much yes. https://en.wikipedia.org/wiki/Karst  
> **username_6**: [Potassium](https://en.wikipedia.org/wiki/Potassium) is [not that different from sodium](https://en.wikipedia.org/wiki/Potassium_chloride). Does it have a solubility heel? Is sodium stealing chlorine from it?  
> **username_5**: @username_6 I suspect this is because potassium feldspars are more stable than sodium feldspars. I found some papers that talk about ion exchange between KCl and albite (NaAlSi₃O₈), but not sure I understand this right. https://www.sciencedirect.com/science/article/abs/pii/S0892687518303984  
> **username_1**: Oceans contain a great deal of Potassium as well. No, it does not "steal" anything. When dissolved, these ions are fully independent from their negative counterparts as long as they don't form an insoluble substance. And Potassium itself does not, just like Sodium. What makes Potassium less prominent in oceans and easilly found in soils is its higher ability to "stick" to clays - related to @username_5 comment. This ability does not prevent water from washing it down, but the process is slower than with Sodium.  
> **username_7**: As they say, if you're not part of the solution, you're part of the precipitate.  

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