If sodium hydrogen carbonate (baking soda, $\ce{NaHCO3}$) is heated to greater than $50~\mathrm{^\circ C}$, it will release carbon dioxide and water to form sodium carbonate:

$$\ce{2NaHCO3 -> Na2CO3 + H2O + CO2}$$

And when the sodium carbonate is heated further to about $850\ \mathrm{^\circ C}$, it will release more carbon dioxide and form sodium oxide:

$$\ce{Na2CO3 -> Na2O + CO2}$$

And according to Wikipedia:

Sodium oxide (SOX) is a chemical compound with the formula $\ce{Na2O}$. It is used in ceramics and glasses, though not in a raw form. It is the base anhydride of sodium hydroxide, so when water is added to sodium oxide $\ce{NaOH}$ is produced.

So it should be possible to produce sodium hydroxide from baking soda and water, with some heat. I heated some baking soda in a gas flame. I am not sure exactly what the temperature of the flame was, but it was enough for the stainless steel spoon that I used to glow red. (If it is useful, the gas where I live is approximately $50~\%$ hydrogen, $25~\%$ methane and the rest $\ce{CO2}$ and $\ce{CO}$, and it burns completely with a blue flame. I suspect this flame to be at least $1000\ \mathrm{^\circ C}$, based on available information about flame temperatures of different fuels.)

However, $\ce{NaOH}$ is produced industrially via the chloralkali process, so there must be problems with using this method. What are they, and are they significant if trying to produce a small amount of sodium hydroxide for a minor experiment?


You are correct: Taking baking soda and heating it to $850~\mathrm{^\circ C}$ will generate sodium oxide and liberate carbon dioxide and water. To generate sodium hydroxide, all you would need to do is wait for the heated soda to cool down and then add water. Beware! The reaction of sodium oxide and water is strongly exothermic. React the two carefully!

A red-glowing spoon is a certain indicator that the temperature of the flame was larger than $500~\mathrm{^\circ C}$. A blue flame is usually an indicator for complete combustion. In the context of bunsen burners, that is typically larger than $1000~\mathrm{^\circ C}$. I cannot say whether your gas flame reached that temperature or not.

Note that once you have created sodium hydroxide you should handle it carefully. Especially do not attempt drinking it and keep it off your hands. Proper safety equipment (goggles, gloves) would be preferable.

Concerning the industrial process, the question is usually not so much what the problems with a certain method are, but rather what the benefit of a certain method is — especially if that benefit relates to cash. Electrolysis of a sodium chloride solution is a very cheap process with sodium chloride and water both available at almost no cost and electricity also typically cheap. Acquiring sodium carbonate is probably already more expensive than the required sodium chloride and heating to $850~\mathrm{^\circ C}$ is certainly much more expensive than letting a current flow through the solution. Hence why the chloralkali process is preferred in industry.


That 850°C figure may be just some arbitrarily chosen temperature for "decomposition begins". Or perhaps what you saw was that Na2CO3 melts at 854°C but is listed as "decomposing" (i.e. there's more than just a simple phase change to contend with). But remember that thermal decomposition is an equilibrium reaction, not a phase change that occurs at a sharply defined temperature.

What I've found while similarly chasing down reaction data is https://www.solvay.us/en/binaries/HeatEffects_of_the_TronaSystem-237230.pdf that shows the equilibrium pCO2 at 850°C at something between 10 and 12mmHg. That's going to be a slow process, but at least it will occur. Now remember that the local pCO2 in your gas flame is going to be much higher than 400ppm-of-100kPa, and according to that table on page 2 you're going to need well in excess of 1200°C just for the reaction to progress forwards. Red heat isn't going to cut it here.



First, the chlor-alkali process is surely cheaper, all the more so by producing other marketable products (chlorine, hydrogen).

Second, in the lab I would consider making the sodium hydroxide in solution by reacting sodium bicarbonate or sodium carbonate with calcium hydroxide.

  • $\begingroup$ For your first point, what do you mean? I used dry baking soda, heated it directly, and then poured it into a beaker. Then, I added water quickly, since I thought it might absorb carbon dioxide from the air. $\endgroup$ May 8 '16 at 18:07
  • $\begingroup$ Edited that out, misunderstood the first time. $\endgroup$ May 8 '16 at 18:17

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