Why are potassium organic salts ever preferable to the sodium equivalents?

$\ce{KOH}$, $\ce{KCN}$, $\ce{KNO3}$, $\ce{K2CO3}$, and plenty of other potassium salts are frequently used in industry. But, as far as I can see, the sodium equivalents are also widely used, often interchangeably. Since Na weighs only 60% of K it seems like the sodium salts should always be preferred because they carry more of whatever base is attached per unit weight (i.e., the sodium salts have "more of what you want"). $\ce{Na}$ is also slightly more abundant (on earth) than K.


So what other factors favor the industrial use of K-salts? E.g., are K-salts just more abundant in natural sources? Are they more reactive because the ionization energies of K are lower than those of Na? Are the behaviors of these simpler salts actually different between the K- and Na- variants, and if so, is there some other elemental characteristic that suggests this?


Going the other way, it seems like if K is preferable to Na, Rb should be even more so (because it appears that in every elemental characteristic in which K exceeds Na, Rb exceeds K). Would it be if it weren't so relatively scarce?




The differences in the properties of the alkali cations are subtle, and the rationale for choosing one over the other will depend heavily on the application. I suspect you're right, that to some extent natural abundance of, e.g., potassium salts relative to sodium ones in certain mineral contexts explains the industrial preference for one or the other -- but, really, variations in abundance are just further examples of these subtle differences in the properties of the cations!


One specific example that I recently came across is the [Hofmeister series](http://www1.lsbu.ac.uk/water/hofmeister_series.html) (see also the [Wikipedia article](https://en.wikipedia.org/wiki/Hofmeister_series)), where the choice of cation makes a demonstrable difference in the solubility properties of proteins. Speaking also from personal experience, the choice of, e.g., $\ce{Na2SO4}$ versus $\ce{K2SO4}$ versus $\ce{(NH4)2SO4}$ as a component of the supporting electrolyte in electrodeposition of copper from acidic $\ce{CuSO4}$ matters significantly to the properties of the deposit obtained.


**Addendum, 17 Apr 2023:** Another example of a significant difference in properties is the aggressiveness of $\geq 1\mathrm{M}$ aqueous hydroxide solutions on aluminum and its alloys. In a $1\mathrm{M}\ \ce{KOH}$ solution, [6061 Al](https://www.azom.com/article.aspx?ArticleID=6636) reacts slowly, releasing small traces of fine bubbles. In a $1\mathrm{M}\ \ce{NaOH}$ solution, the reaction is vigorous, with significant gassing and relatively rapid consumption of the metal.


 10

There's a couple of reasons with which I'm familiar:


1. **Degree of hydration:** Some sodium salts are readily available as their hydrates, but the anhydrous form is more expensive to produce due to how hygroscopic it is. On the other hand, the potassium salt of the same anion may not form hydrates; for example, sodium molybdate is commonly found as the dihydrate while potassium molybdate does not form hydrates. In these cases, the potassium salt may sometimes be preferred if a water-free procedure is required, or if the degree of hydration is high enough that the hydrated sodium salt would actually contain less of the attached anion(s) than the anhydrous potassium salt.
2. **Working with sodium-sensitive organisms:** For example, working with organisms that are used to growing in soil. Many soils have much higher potassium than sodium and therefore soil-based life is often well-adapted to potassium - many plants even require a lot of potassium as a major macronutrient. However, the same organisms will experience significant salt stress if placed into an environment with high sodium. Hence, potassium salts are often used in such setups; for example, phosphate buffer is often prepared with KH2PO4 and K2HPO4 rather than with the sodium equivalents to avoid the aforementioned salt stress (e.g. <https://www.usbio.net/media/N1000/Nematode%20Growth%20Medium/data-sheet>). Similarly, a lot of plant growth regulators are solubilized using potassium hydroxide instead of sodium hydroxide (<https://phytotechlab.com/media/documents/ProductLiterature/PlantGrowthRegulators.pdf> - notice the number of entries in the solvent column saying "KOH" but none saying "NaOH").
3. **Double-displacement reactions where you want to precipitate the potassium salt of something else:** As a very general rule of thumb, most potassium salts are less soluble than their sodium equivalents. If you were conducting a double-displacement reaction where you wanted to separate the two products, it might make sense to select the potassium salt as one of the reactants so the potassium byproduct can be precipitated.
4. **Differing physical properties of the products:** For example, saponification with sodium hydroxide typically results in a much firmer, solid soap bar while saponification with potassium hydroxide usually results in softer soaps or even liquid soap. I'd expect there to be other situations where use of the sodium compound results in a product with significantly different physical properties from the potassium compound.


As for rubidium - it varies. Rubidium could be used over potassium in 3 (precipitation of potassium side-product) if it were more common while it might also be used in 1 (degree of hydration) on a case-by-case basis. I'm a bit iffy on whether rubidium could replace potassium in 4 (physical properties) as it might be the case that the physical properties of the potassium-based product are exactly what is sought after (e.g. liquid soap) and use of rubidium would result in an unusable product. I doubt rubidium-based compounds would replace potassium in biological applications\* (use 2); experiments where rubidium was used as a potassium substitute in lab animal diets resulted in the deaths of said animals (<https://doi.org/10.1152/ajplegacy.1943.138.2.246>).


\*: Though, in an alternate timeline where rubidium was really common on Earth, any biological life that developed would probably have ways of using or detoxifying rubidium, the same way most terrestrial life has learned to utilise or detoxify elements like potassium, copper, or selenium when they're encountered at the levels which life is accustomed to. In that alternate timeline, rubidium might be used when working with organisms in the same way potassium is today, but it's unlikely rubidium would replace potassium in such applications e.g. if we came across a big chunk of rubidium in space and mined it.


 3

Sodium and potassium salts often can be used interchangeably - it these cases the difference in price and accessibility can be important. For example one can easily buy cheap sodium chloride in shop, but not potassium chloride, which isn't as cheap or accessible as NaCl. Other differences can be however very important.


There are some sodium and potassium salts that strongly differ in water solubilities. For example antimonates - [potassium hexahydroxoantimonate (V)](http://www.sigmaaldrich.com/catalog/product/fluka/60500?lang=en&region=US) is even used to detect sodium via precipitation. On the other hand solubility of [potassium perchlorate](https://en.wikipedia.org/wiki/Potassium_perchlorate) is about 100 times lower than that of analogous sodium salt.


Difference between sodium and potassium salts is big for living organisms. For example [potassium chloride](https://en.wikipedia.org/wiki/Potassium_chloride) is rather bitter than salty and it's very dangerous when injected - can easily cause heart attack.


 3

One difference between $\ce){Na^+}$and$\ce{K^+}$ is the former, more compact cation is more likely to attract water molecules and form hygroscopic salts. This is a disadvantage in analytical applications, as [Wikipedia](https://en.wikipedia.org/wiki/Sodium_permanganate) reports in the case of sodium permanganate:



> 
> Sodium permanganate behaves similarly to potassium permanganate. It dissolves readily in water to give deep purple solutions, evaporation of which gives prismatic purple-black glistening crystals of the monohydrate $\ce{NaMnO4·H2O}$. The potassium salt does not form a hydrate. Because of its hygroscopic nature, it is less useful in analytical chemistry than its potassium counterpart.
> 
> 
> 


Another difference, in the case of permanganates, is that the common method of producing [potassium permanganate](https://en.wikipedia.org/wiki/Potassium_permanganate) does not work with sodium as the cation:



> 
> Potassium permanganate is produced industrially from manganese dioxide, which also occurs as the mineral pyrolusite. In 2000, worldwide production was estimated at 30,000 tonnes.[1] The MnO2 is fused with potassium hydroxide and heated in air or with another source of oxygen, like potassium nitrate or potassium chlorate.[1] This process gives potassium manganate:
> 
> 
> $\ce{2 MnO2 + 4 KOH + O2 -> 2 K2MnO4 + 2 H2O}$
> 
> 
> (With sodium hydroxide, the end product is not sodium manganate but an Mn(V) compound, which is one reason why the potassium permanganate is more commonly used than sodium permanganate. Furthermore, the potassium salt crystallizes better.[1]).
> 
> 
> 


So preparation of sodium permangnate requires more extensive and more indirect preparation giving, as noted above, a less desirable salt for analytical uses.


**Cited Reference**


1. Miessler GL, Fischer PJ, Tarr DA (2014). *Inorganic Chemistry* (5th ed.). Pearson. p. 430. ISBN 978-0321811059.


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