Why is sodium carbonate used to absorb sulfur dioxide in the synthesis of sodium sulfite, instead of sodium hydroxide?

You’ve analyzed it quite comprehensively; in my opinion

There should also be the consideration of oxidation by oxygen in the air; the use of sodium carbonate produces carbon dioxide, which reduces the oxygen level. Both sodium sulfite and sodium bisulfite will be oxidized. I’m just talking nonsense’

Thank you/I forgot to consider oxidation. If oxidation occurs, using a mixture of sodium hydroxide and sodium carbonate will cause oxidation, right? Theoretically, sulfur dioxide will first react with sodium hydroxide, with no carbon dioxide being produced; and even if some carbon dioxide is generated, the amount will be less compared to using pure sodium carbonate. Oxidation still occurs in that case. If pure sulfur dioxide can be used, then sodium hydroxide can be employed, right? I’ve read that the temperature needs to reach 80 degrees for sulfur dioxide to react – higher temperatures result in faster reactions. Online, it’s said that at lower temperatures, the reaction is not very significant. So, what temperature should be maintained for the sodium carbonate solution in winter? If pure sulfur dioxide is used, the temperature can be room temperature. Is it necessary to raise the temperature when using sodium hydroxide? Sodium sulfite has a low solubility; will it precipitate during the absorption process? If it does precipitate, could this cause blockages in the absorption tower? I’ve seen online that an absorption tower is used in combination with a reactor – is it to separate the precipitated sodium sulfite and sodium bicarbonate first, allowing only the clear liquid to circulate? Once absorption is complete and hydrogen sulfite is formed, its solubility increases, so no more precipitation occurs. I’ve also seen that there seems to be a decolorization step – is this done after hydrogen sulfite is formed, or after sodium hydroxide is added to convert it into sodium sulfite?

Sodium carbonate is a weak base, and its absorption effect is inferior to that of sodium hydroxide.

I checked the market prices of soda ash and caustic soda; soda ash costs a little over 2000, while caustic soda costs over 3000. Therefore, soda ash is the more cost-effective option. Moreover, its purity is relatively higher, and it poses less risk. The price of ion-exchange membrane caustic soda is less than 1,000, but its concentration is 32%; when adjusted accordingly, the price is roughly the same. As for the effect of air oxidation on sodium bisulfite, I don’t think it is a major factor; the reactor definitely needs to be sealed, and it can also be protected with nitrogen.

{:3_49:} It definitely has nothing to do with oxidation; the gas is fed continuously, so the oxygen is removed long ago. Your process should be sodium carbonate + sulfur dioxide — sodium bisulfite + sodium hydroxide — sodium sulfite. Sodium hydroxide isn’t needed at the beginning; it’s likely related to the heat of reaction – sodium hydroxide releases heat very intensely. The use of sodium hydroxide later on is intended to facilitate rapid neutralization, stabilize the pH, and obtain a pure product. Using sodium carbonate for neutralization makes it easy to add too much of it. . . After all, both salts are weak acid salts

We have a system using sodium hydroxide to absorb sulfur dioxide, capable of producing 100,000 tons of anhydrous sodium sulfite per year; it has been in operation for five or six years now.

Thermal effect. Sodium hydroxide absorbs heat so intensely that the water from chickenpox turns into steam.