Take hydrogen peroxide decomposition for example. I get that it’s spontaneous and the products have less energy in the bonds than the reactants. How is the formation of these new bonds exactly increasing the kinetic energy of surrounding molecules? Is it some sort of light energy colliding with the molecules?
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The energy originally stored in the bonds of the peroxide has to go somewhere. Part of it forms the new bonds of the decomposition products. The rest gets converted to heat. Oxygen oxygen single bonds in peroxides not only contain a decent amount of energy, they also really don’t want to keep on existing and will take the first opportunity to react.
Say the peroxide has 2 J of bond energy and the decomposition products have 1 J of bond energy. That other 1 J has to go somewhere since you can’t create or destroy energy. That 1 J becomes heat.
Note: the numbers I gave are for example purposes only and in no way the actual bond energy numbers.
As to why it happens, things tend towards the most stable lower energy state they can be in. Why doesn’t everything just spontaneously react? The difference in the energy of the two states matters. Very stable to even more stable is less likely to happen. Sometimes, not all the atoms required for a reaction are present and also activation energy (how much of a nudge you have to give for the reaction to start) is high.
Think of the atoms like little magnets. When those magnets snap together, they move at a decent speed. When they collide it’s usually pretty violent, from all of that kinetic energy they picked up as they flew together.
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