I understand that energy is NEEDED to BREAK bonds. But don't chemical bonds already have some form of potential energy in them already? So wouldn't it make sense that even though energy is required to break them, their potential energy is also RELEASED once they are broken?

In other words, shouldn't the bond breaking process be considered simultaneously endothermic AND exothermic?

Side note: I'm studying grade 12 level chemistry right now, so I don't have too much knowledge on the nitty gritty details of this concept. That being said, It would be greatly appreciated if anyone could explain this in very simple terms.

  • $\begingroup$ A reaction overall cannot be simultaneously endo- and exothermic, the values would cancel out such that the reaction would be overall endothermic or overall exothermic, taking into account all of the bonds being made/broken. $\endgroup$ – NotEvans. Dec 28 '17 at 22:24
  • $\begingroup$ Think of "exothermic" and "endothermic" as representing the difference in energy from start to finish. Yes, energy may need to be added to trigger a reaction (e.g. lighting thermite, which takes a lot), but if, at the end, more energy was produced than added, it is exothermic. $\endgroup$ – DrMoishe Pippik Dec 28 '17 at 22:53
  • 2
    $\begingroup$ Bond breaking in itself does not release energy!! $\endgroup$ – orthocresol Dec 29 '17 at 0:48

Bond breaking does NOT release energy!!

In the course of a chemical reaction, you typically break some bonds and form some new bonds. The bond breaking is always endothermic. The formation of new bonds is exothermic, so depending on whether the old bonds or the new bonds were stronger, the reaction overall can be either endothermic or exothermic.

When somebody says chemical bonds “contain potential energy” it is a HIGHLY misleading statement. You justifiably interpreted this as: when you break bonds, they release the potential energy that was stored within. This is, unfortunately, wrong. That phrase should honestly be banned, but somehow it has always persisted.

What that phrase really means is more nuanced. It means that some chemical bonds can be broken in a reaction which subsequently produces more energy via bond formation.

Bond breaking is always endothermic!

For further reading see: Is Bond Formation "Strictly" Exothermic?

  • $\begingroup$ en.wikipedia.org/wiki/Helium_dimer#Molecular_ions ? $\endgroup$ – Karl Dec 29 '17 at 16:41
  • $\begingroup$ Impressive counterexample. “Always” should therefore be replaced with “almost always”, then. $\endgroup$ – orthocresol Dec 29 '17 at 16:56
  • $\begingroup$ Typical case when the original question was not well put, here had this misunderstood/sloppy "potential energy" thing in it. Great how this makes you think about long learned stuff again. :-) $\endgroup$ – Karl Dec 29 '17 at 19:18

A bond is always some local energy minimum. You definitely need to add some energy to get out of there, but when you separate the remaining parts to infinity, you might get more energy back than you put into it. Certainly not if you split water or sodium chloride.

So you're principally correct, but the two (or more, breaking a bond will usually induce further rearrangement in the parts) processes cannot be separated. Can't break a bond but not remove the parts from each other.

  • $\begingroup$ While I understand what you mean, I think it doesn’t hit the nail on the head. If you decompose an azide the bond breaking is still endothermic. Subsequent bond formation may well drive the reaction (in this case the formation of the N-N triple bond) but OP is only asking about bond breaking, not a reaction as a whole. The problem with OP is that they have been taught something about how chemical bonds “contain potential energy”, which is obviously not true because of your first sentence. $\endgroup$ – orthocresol Dec 29 '17 at 0:45
  • $\begingroup$ @orthocresol The azide example might be wrong, and definitely too complex, right. My thinking was this: Two particles at infite distance are at energy level zero. You bring them closer, and they will first repel each other (unless it's an anion and cation), so E > 0. When (if) a bond forms, the energy level goes down again, but not necessarily below zero. $\endgroup$ – Karl Dec 29 '17 at 12:38
  • $\begingroup$ (I notice I ignored London etc. interaction. Hm.) $\endgroup$ – Karl Dec 29 '17 at 12:58

Not the answer you're looking for? Browse other questions tagged or ask your own question.