Can potassium carbonate be used to calibrate hygrometers? Context:
In the process of trying to calibrate a number of hygrometers, with Sensirion MEMS electronic sensors, I came across the method of using wet salt: an airtight container will stabilize at a predictable relative humidity over time, given a particular salt and temperature.
For a list of salts and their RH values at particular temperatures see this page on the website of Vernier, a maker of STEM education equipment. The various salts suggested for this method repeatably produce a range of low to high relative humidities, also allowing for two-point calibration.
I require my hygrometers be calibrated close to 40% RH. With a predictable 43% RH that makes potassium carbonate the ideal candidate. However, when I wettened potassium carbonate with water it underwent an moderate exothermic reaction. Googling what was going on, I read &quot;potassium carbonate dissociates completely in water into potassium (K+) and carbonate ions (CO32-).&quot;
Which makes me wonder:
Can potassium carbonate be used at all for the wet salt calibration method, or does its swift dissociation render it useless for this application?
Note that all salts, used in humidity calibration, dissociate to their respective hydrated cations and anions when dissolved.
From the solution crystallizes the sesquihydrate $\ce{K2CO3·​\frac 32 H2O}$ (&quot;potash hydrate&quot;). Heating this solid above 200 °C gives the anhydrous salt (Wikipedia).
It seems you have anhydrous carbonate, that hydrates itself exothermically. You can still use it, but you need to prepare rather the saturated solution with excess of carbonate, that would integrate the crystal water into itself.
Or, let the prepared saturated solution partially evaporate and crystalize.

What is important is maintained equilibrium.
$\text{Air 43% humidity }\ce{ &lt;=&gt; K2CO3} \text{ saturated solution } \ce{&lt;=&gt; K2CO3 . 3/2 H2O}$
Anhydrous $\ce{K2CO3}$, wetted by water, rather just absorbs water into its crystall lattice. If you add too little water, no solution would left, similarly what would happen to the plaster of Paris.
You need the carbonate in hydrate form, together with its saturated solution. Take some water, add there carbonate to dissolve as much of as it can, and then some. Let is settle.
If you had the hydrate, you could just moisten it. But you do not have it.
There is the third option - moistening  it enough so some liquid phase always remains. For $\pu{100 g }\ce{K2CO3}$, you need about $\ce{20 mL}$ of water just for the crystal water, so it becomes the part of the solid.
The carbonate solubility is $\pu{110.3 g/100 mL}$, so I would carefully mix some $$\pu{100 g}\  \ce{K2CO3} + \pu{50 mL}\ \ce{H2O}$$
or recalculated to  more suitable amounts.