I realised the other day that I'm not sure why the reactions of strong acids with metals are exothermic. Any reaction is exothermic if there's more energy released from bonds being made than is used in breaking bonds, of course (using the word 'bonds' in a sufficiently loose sense).
In these reactions, metallic bonds are broken; covalent bonds are formed in hydrogen, but the hydrogen is in a gaseous state, so its molecules must be less attracted to each other than the hydrogen ions were to each other in the liquid. I'd vaguely guessed that the latter fact would make this an endothermic reaction (it works for sodium carbonate, after all) but evidently not!
Is there a rigorous way of calculating the total enthalpy change, or an easy explanation that I'm missing?
I guess I can calculate the enthalpy change of every one of the state changes involved, but it's not obvious that that's always going to come out negative (which I gather it does) so I feel like there's probably a less mathematical way of explaining this observation.