Just to repeat what's already been explained to you; your statement implying only "heat" and light are sources of (external) energy is really misguided/wrong. Especially in the context of chemistry. To name two other sources of external "energy", collisions and electricity. Energy, as is obvious once you've learned some Freshman chemistry and physics, is an abstract concept, I can't point to something and say "That's energy" - it's something else which contains energy. Since atoms obey quantum mechanics, and since quantum mechanics is about probabilities, a reaction in which an A-B bond with a near-by C may change to (say) an A-C bond, and will if the A-C species is (has) lower energy, even near absolute zero, in the dark, etc. OTOH, in general a reaction is more likely to occur as the difference in energy between reactants and products increases and as the activation energy decreases (activation energy is the difference in energy between the transition state and the reactants - see Wikipedia). But lets take an example which might be what you were thinking about. Let's imagine a molecule with a X-Y bond (while my example here is conceptual, there are plenty of examples of it). If the "internal degrees of freedom" have sufficient energy to break the X-Y bond, then it may come to pass that the bond breaks. So, what do I mean by "internal degrees of freedom"? Well, conservation of energy and momentum hold (almost perfectly), but say there is a bond in the molecule which is vibrating (all bonds vibrate, even at 0 K, although the "energy" of vibration of a bond in its lowest energy state can't be used). If that energy is above the lowest possible for that vibration, it can "couple" to the X-Y electronic energy and increase its vibration, possibly to the point where the bond breaks. What other "internal degrees of freedom" does a molecule have? Spin, steric strain (conformation), it also has velocity (momentum) and angular momentum, but those are conserved. A molecule (or molecular fragment) may already have enough energy internally to spontaneously decompose. The rate this happens can be fast (boom) or take centuries (diamond converting to graphite). Repeating myself: even a molecule "sitting" on an atomic lattice at 0 K which is near to an atom which tends to substitute can "make the switch", BUT this is typically very slow to happen. Hydrogen is the principle atom which finds this kind of quantum mechanical "tunneling" to be easy. Responding to a different part of your question (is it because...) Well, I suppose you could look at it that way for some reactions. It's better, imho, to look at it as nature's tendency to attain the lowest energy state. So if a "highly reactive" element were to replace a less reactive element BUT the product were even MORE highly reactive, then the reaction wouldn't be likely to occur spontaneously. So it's not about "reactivity" per se, but about lowering the system's energy. Electronegativity is a concept which is well defined for elements in their ground state but what is the electronegativity of Fe(II) compared to Co(III) ? or of Cl(-) compared to OH(-). Chemistry is almost (but not quite) exclusively about electronic (and steric) changes, so all of it (almost) will be "about" electrons and charge (electron density) differences. It may be helpful in some situations to look at it that way. But in others it will not lead to obvious predictions/conclusions. That is, in some situations it will be useful, in others not so much. But that's nothing new. Most of science is like that.