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I don't just mean reactions that require heat to proceed, storing surplus energy in chemical bonds. I wonder about strongly endothermic reactions that suck heat out of environment.

You take some substance A (e.g. Ammonium Nitrate), and some substance B (e.g. water), both at room temperature. You mix them together in a beaker which is room temperature, all performed in room temperature air. As the reaction begins, the two substances binding into substance C, the beaker cools down quite a bit below room temperature.

This is what we observe on macroscopic scale.

I wonder what happens on the microscopic scale - chemical particles, atoms, elementary particles, their energy. Normally entropy would suggest everything would remain in equilibrium but suddenly we have a higher energy concentration within the substance at cost of energy of the environment. It is not easily reversed. What happens that makes particles "want" to bind so much that they specifically "steal energy" from the environment just so that they can react?

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The actual subject you want here is thermodynamic/statistical mechanics/physical chemistry. –  dmckee Feb 15 '13 at 18:00
    
@dmckee: Not really - I don't really want the maths of the averages of real thermodynamics but the micro-scale view of what happens with and within the atoms, how they absorb stray photons, how the electrons skip energy levels etc. –  SF. Feb 15 '13 at 18:21
    
SF, Yeah, really. Those averages tell you about the entropy/free-energy costs of having the system in various macroscopic states. Then you can ask for a microscopic explanation of each term. Why that work? Why that pressure? Why that chemical potential? And so on... But you have to ask about each term for each process. When you ask about photons you introduce a third level of consideration. Quantum processes occur because they are not forbidden and the question is always "At what rate?". –  dmckee Feb 15 '13 at 18:34

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I don't just mean reactions that require heat to proceed, storing surplus energy in chemical bonds. I wonder about strongly endothermic reactions that suck heat out of environment

A reaction that requires heat to proceed, a reaction that sucks heat out of the environment, and an endothermic reaction are all the same. These are all just descriptions of reactions that occur because entropy (S) increases, despite that fact that enthalpy (H) also increases.

$G = H -TS$

A reaction is favorable if Gibbs free energy decreases. The reaction can still be favorable, despite enthalpy (H) increasing, if entropy (S) increases enough.

You take some substance A (e.g. Ammonium Nitrate), and some substance B (e.g. water)... I wonder what happens on the microscopic scale

  1. you have to separate molecules/ions/atoms of A from others of A. This could involve breaking apart ionic interactions in a crystal, or intermolecular interactions for example. This takes energy.
  2. you have to seperate molecules/ions/atoms of B from others of B. For a solvent, this would involve breaking apart intermolecular interactions such as dipole-dipole interactions. This takes energy.
  3. new interactions between A and B form. This releases energy.

If the energy released in step 3 is less than the energy required for steps 1 and 2, the process is endothermic.

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