Is the Conventional Wisdom of Energy and Bond Formation Always Accurate?

[QuestionMarks]

Just double-checking myself on something:

Typically at the high-school level, we say that it takes energy to break a bond, but energy is released to form a bond. This has always bothered me because even with a naive understanding of enthalpies we can see formations with different signs. To me then, this conventional wisdom is a decent generalization only for stable bonds, but for molecules that spontaneously break, this wisdom is not so sound.

Sound alright?

 

[Simon Bridge]

By definition, a bond involves a lower energy state the unbound state.
It follows that you need to supply energy to break a bond.

Some bonds are easily broken because the energy required is very low.
For instance, some combinations of atoms are unstable at room temperature because the hea in the room quickly supplies the energy needed to break the bond. It is also possible for a configuration to be finely balanced so the energy vs distance graph has a local minima on top of a global maxima ... thus, a small amount of energy input means that the configuration cannot come back together again.

Can you come up with an example of a molecule that spontaneously breaks without input of energy?

 

[Borek]

Sound alright?

No. Every bond means a depression in the potential energy landscape, to leave the depression you need to deliver the energy to jump from the depression. Unstable bonds have the depression very shallow, and as energy is not distributed uniformly between molecules it often happens that part of the population of molecules have enough energy for the unstable bond to be broken - but it doesn't mean it breaks without energy.

 

[QuestionMarks]

My problem then seems a poor choice of semantics as I'm willing to imagine a "molecule" effectively held together by some externally applied force (which would spontaneously break upon removal of that external energy)...but we would not call this a molecule.