# Why don't decomposition reactions form pure elements as products?

For example, if I had a reaction like this:

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

• Why does it not break down all the way down to its elements?
• What makes it form such "intermediate" products, stopping at water, carbon dioxide, etc.?
• If you heat sodium bicarbonate strongly you can convert it to $\ce{Na2O}$. In general the reaction favors the most thermodynamically stable chemical. – MaxW Feb 21 '17 at 1:28

Using your example, you simply have not given the system enough energy to do so. Water and $\ce{CO2}$ are very stable compounds under ordinary conditions of temperature and pressure, and thus are the preferred products under standard conditions for that reaction.

If you were to heat the system further, further decomposition would occur. For example, at higher temperatures you will convert $\ce{Na2CO3}$ into $\ce{NaO}$ and more $\ce{CO2}$. Under extreme conditions, i.e. thousands of degrees K, you will ultimately have enough energy to break all of the chemical bonds and then you will be left with the elements.

• Quibble - I'd explicitly note that "extreme conditions" mean that the whole sample will be atomized to the gaseous elements. – MaxW Feb 21 '17 at 6:41
• Just one disagreement with this answer. Not elements, rather atom of those elements. – CupC_56 Feb 21 '17 at 6:55
• Or a plasma I guess. – airhuff Feb 21 '17 at 6:57
• @CupC_56, my plasma comment was a response to Max. Another good point though. And for that matter we could go nuclear with our energy source to get to subatomic particles, but I guess at that point we have to move the discussion to Physics.SE ;) – airhuff Feb 21 '17 at 7:02