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Along the x-axis, the progress of the reaction is observed, while along the y-axis, the potential energy of the reactants/products are observed. They keep rising until they attain the required activation energy for breaking of older bonds, and formation of new ones. But what exactly is this activation energy? Is it a form of potential energy (stored within the chemical bonds), or is it a form of kinetic energy (vibrations of the particles)? I don't really understand why, if the activation energy is just potential energy, it results in collisions (without any motion) and thus forms new compounds, given that, we all know, potential energy is the energy of a body at rest.

If it is not so, then what is the role of chemical potential energy into all of this, at all?

I am still a student, stuck within concepts, so please, keep the answer as brief as possible...

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It is just potential energy: the potential energy needed for the reaction to happen.

This potential energy obviously needs to come from somewhere, and that's where kinetic energy comes in. If the molecules have high enough kinetic energy at the point of the collision for that reaction to happen, then part of that kinetic energy is converted to potential energy, and the reaction proceeds.

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This is also why most reactions get faster as you increase the temperature. Higher temperature means more kinetic energy, meaning more successful reactive collisions.

Edit: I should add that there seems to be some sort of misconception in your question. There's nothing along a potential energy graph that would result in a collision.

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  • $\begingroup$ But what exactly is the role of Potential Energy in a "collision". This kind of sounds weird, because potential energy is solely the energy at rest, whereas collisions come from Kinetic energy itself. Aren't both of them contradictory to each other?? $\endgroup$
    – Aniket
    Jun 30 at 12:11
  • $\begingroup$ I don't understand the question. The potential energy plays no role in a collision, apart from the fact that the total energy needs to be conserved. You start with a given amount of total energy, distributed between kinetic and potential energies, and those two can interchange into each other, as long as the sum is conserved. The activation energy tells you that how much of the total energy needs to be in potential energy at least for the reaction to happen. $\endgroup$
    – Szgoger
    Jun 30 at 12:13
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The potential energy is the energy due to the positions of the atoms and their electrons have in a molecule that is a sort of combination of reactant and product, i.e. a transient structure from the path from reactants to products, some times called a 'transition state'. This barrier can be overcome by collisional (kinetic) energy between the reactant and product but also by thermal energy since by the Boltzmann distribution there is a small chance that a molecule may have, for a short period of time, many times the average energy expected at that temperature. Thus reactants may collide many millions of times before a reaction occurs, if the barrier is high, or only a few times if the barrier is low relative to average thermal energy. Consequently we observe rate constants in different types of reaction that vary from $10^{12}$/sec to less that 1 per hour.

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Maybe this additional perspective will help. Reactions involve bond breaking and bond making. Breaking bonds results in an increase in potential energy (there is less attraction between atoms; you must add energy to break bonds, therefore, increase in PE). Bond Making results in a decrease in potential energy (there is more attraction between atoms; less obvious, energy is released when bonds form). So, in the first part of the reaction bond breaking dominates and the potential energy is increasing. In the latter stages of the reaction bond making dominates and potential energy is decreasing. Thus, the graph goes up and then it goes down. This is a very simplistic view, but it will give you a better perspective.

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