It's normally assumed that a higher pH of solution would slow nucleation rates and hence lead to larger crystal formation. I've found literature that it is due to solid-liquid inter-facial tension. Which is a concept that I do not fully understand.

In my experiments, higher pH solutions led to the formation of smaller crystallites. I can't make sense of why its contradicting the literature. Here are the images from my experiments:

Images of crystals at different pH

My samples are precipitates produced from mixing aqueous $\ce{CaCl2}$ +$\ce{MgCl2}$ and $\ce{Na2CO3}$ solutions. The elevated pH images were produced through the addition $\ce{NaOH}$ to the previously mentioned mixture.

If you have any comments or ideas, it would be greatly appreciated. I am a engineering student who is trying his hand at inorganic chemistry.

Thanks in advance.

  • $\begingroup$ Pure interfaces reason do not take into consideration possible common ions effects, where common means whenever there is an equilibrium that links precipitation and pH. $\endgroup$
    – Alchimista
    Commented Apr 13, 2019 at 7:46

2 Answers 2


At low pHs there is little "free" $\ce{CO3^{2-}}$ to precipitate the $\ce{Mg^{2+}}$ and $\ce{Ca^{2+}}$ cations. Most of the carbonate species are dissolved $\ce{CO2}$, $\ce{H2CO3}$ and $\ce{HCO3-}$. Thus the precipitate forms slowly and you get relatively large crystals.

At high pHs there is a lot "free" $\ce{CO3^{2-}}$ to precipitate the $\ce{Mg^{2+}}$ and $\ce{Ca^{2+}}$ cations. Thus the precipitate forms relatively rapidly and you get relatively small crystals.

I have no idea what literature would lead you to believe otherwise.


Aside of influence on the crystallization rate, you have also consider the direct chemical influence.

The sodium carbonate hydrolyzes in not enough alkaline solutions. $$\ce{Na2CO3 + H2O <=> NaHCO3 + NaOH}$$

Precipitation of carbonate is interfered by solubility of Bicarbonates. $$\ce{CaCO3 v + H2O + HCO3- <=> Ca(HCO3)2 + OH-} $$

Addition of the hydroxide affects both processes. It shifts the equilibrium towards precipitation and also speeds up the precipitation.

The same for magnesium.

For magnesium, there is possibility of a side reaction of precipitation the insoluble hydroxide. But lower carbonate solubility will probably overrule this.

$$\ce{MgCl2 + 2 NaOH -> Mg(OH)2 v + 2 NaCl}$$


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