How to extract pure potassium carbonate from ash?

Question

In our village, lots of ash is produced which is thrown off. I want to extract pure potassium carbonate from that ash. I am able to extract the potassium compound and a soluble compound by recrystallizing and evaporating but I found impurities. The first test I performed is electrolysis. This test confirmed the presence of chlorine gas which is able to change the litmus paper to white. The second test I performed is flame test. During the flame test, I got a yellowish orange color while my expectation was bluish violet. So, I assume there is sodium. How can I extract the pure compound?

To conclude, there is a small amount of potassium chloride, sodium chloride, sodium carbonate which are the impurities.

Answer 1

Historically speaking, potassium derives from pot ash and now potash when -- if there were no geologic sources of $\ce{K2CO3}$ -- woods were burnt in a pot. The addition of water will dissolve $\ce{K2CO3}$ and other high soluble matter, but not the ash. According to wikipedia, the addition of amine may lower the amount of KCl, but basically then, as by Poutnik's numbers, it is fractional crystallization.

Additions:

• You may find temperature dependent solubility of $\ce{K2CO3}$ as sole solute in water, e.g. by Schultz et al. in Ullman's Encyclopedia of Industrial Chemistry (doi 10.1002/14356007.a22_039):

\begin{array}{c|c} \mathbf{Temperature (^\circ C)} & \mathbf{Solubility(g~K_2CO_3/ 100~g~H_2O)} \\\hline \text{0} & \text{105.5}\\ \text{10} & \text{108.0} \\ \text{20} & \text{110.5}\\ \text{30} & \text{113.7}\\ \text{40} & \text{116.9}\\ \text{50} & \text{121.9}\\ \hline \text{60} & \text{126.8}\\ \text{70} & \text{133.1}\\ \text{80} & \text{139.8}\\ \text{90} & \text{147.5}\\ \text{100} & \text{155.7}\\\end{array}

The caveat however is that the precipitate obtained from these solutions in this temperature range isn't $\ce{K2CO3}$, but the sesquihydrate, $\ce{K2CO3 * 1.5 H2O}$. Like $\ce{K2CO3}$, it has to be stored in a dry environment or otherwise, deliquesces. It looses water if heated to $\pu{130-160 ^\circ{}C}$.

• Consider the economic and ecologic perspectives: Synthetic $\ce{K2CO3}$ may be much cheaper than collecting and washing the ashes, and subsequent steps of purification. It may be purer, with less salination problems for farmers, too. E.g., how do you will treat waste waters containing other salts than $\ce{K2CO3}$? How do you generate the heat to recrystallize the material? Don't get into the rut by yourself or / and your friends to collect wood to heat the dishes, which generates new ashes, which again will be leached ... leading quickly into a disaster. If put into execution, see how solar energy (e.g., evaporation under the sun to preconcentrate, solar ovens to boil) may be used instead.

Answer 2

The composition of the wood ashes varies with the nature and the origin of the wood. In the average, it is a mixture of $40$ % - $70$ % $\ce{CaCO3}$, $5$% - $10$% $\ce{MgCO3}$, $5$% - $10$% $\ce{K2CO3}$, $5$% - $10$% $\ce{Na2CO3}$, $2$% - $5$% $\ce{SiO2}$, $2$% - $5$% $\ce{Ca3(PO4)2}$, $0.5$% -$2$% $\ce{Al2O3}$, <$1$% $\ce{NaCl}$, <$1$% $\ce{KCl}$, plus traces of manganese and iron. So it is difficult to purify, and/or to extract any of these compounds, as a lot of these compounds are insoluble in water: $\ce{CaCO3, MgCO3}$, $\ce{SiO2, Ca3(PO4)2, Al2O3}$. The only soluble compounds are $\ce{K2CO3}$ and $\ce{Na2CO3}$, plus traces of $\ce{NaCl}$ and $\ce{KCl}$, and they form less than $20$% of the total mass of the ashes. The solution obtained from washing the wood ashes is a mixture of sodium and potassium carbonate. The problem is how to separate those two carbonates. Fractional crystallization does not seem easy. I would rather transform them into nitrates by adding nitric acid : at $0°$C, $\ce{KNO3}$ is not very soluble ($5$ g/$100$ mL), which is much less than $\ce{NaNO3}$ ($73$ g/$100$ mL). So obtaining $\ce{KNO3}$ from the mixture by cooling a hot solution should be relatively easy.

Ref.: J. Hartwig, Ann. der Chem. u. Pharm. XLVI, p. 97. A. Wurtz, Dictionnaire de chimie pure et appliquée, Vol. 1, p. 641 (1891).

Answer 3

It seems that by partial crystalization, it can be somewhat purified as the most soluble salt, but cannot be easily purified to high purity degree. It may be andvantageous to crystalized near water freezing point.

Solubility in $\pu{g/100 mL}$ for $\pu{0^{\circ}C, 20^{\circ}C, 100 ^{\circ}C}$

$\ce{KCl}$:28, 34.1, 56.3
$\ce{K2CO3}$: 105, 111, 156
$\ce{NaCl}$:35.7, 35.9, 39
$\ce{Na2CO3}$: 7, 21.5, 45.5

Potassium salts and Sodium salts

Answer 4

The trick is to convert the soluble K2CO3 to the more insoluble KHCO3. Extract the ash with warm water; cool, filter to remove CaCO3 and other insolubles. Get together with the beer and wine makers and bubble their CO2 in and precipitate KHCO3, Filter, dry and heat to get K2CO3. You must refine the process to get good yields. It might be practical to recover Ca and Mg carbonates by treating a slurry of the filtered solids with CO2 since their bicarbonates are reasonably soluble. Boil the solution to precipitate the carbonates. BTW your tests for purity are specious. There are very low levels of sodium, chlorine in fresh water vegetation; that is why herbivores look for salt licks.