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In a Redox reaction, a valence electron is transferred/shared from one atom to another. The Electronegativity of the atom will dictate its tendency to attract shared electrons.

More specifically for this question, Combustion is an exothermic redox reaction that reduces the enthalpy of the system. The reaction can be quantified by the Bond Enthalpy of the chemical bond.

Now the question: In an exothermic redox reaction what exactly is happening that the new molecule gains momentum? Is it the acceleration of the electron through the electrostatic field created between the now charged ions and its subsequent capture by the adjacent atom thereby imparting momentum onto the new molecule as a whole? Or is it the attraction from the ions to each other until they "collide" and bond? If so, how do we calculate the electric potential created between the two atoms and how this imparts energy into the resulting molecule?

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  • $\begingroup$ Your question is very curious. Although I do not have a canonical answer to your question, I would like to emphasize one thing. The electric field caused by the electrostatic potential has the property of conservative force, so the electrostatic potential cannot be the candidate of the continuous energy source of the redox reaction. $\endgroup$
    – HEMMI
    Commented Sep 25 at 2:25
  • $\begingroup$ @TheTiler do you suggest a more appropriate tag? Quantum chemistry? $\endgroup$
    – JustAnotherUser
    Commented Sep 28 at 23:35
  • $\begingroup$ I'm not suggesting anything, I just thought of this reaction:en.wikipedia.org/wiki/Aluminothermic_reaction to weld the rails where the iron becomes liquid.youtube.com/watch?v=5uxsFglz2ig $\endgroup$
    – The Tiler
    Commented Sep 29 at 9:09

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While electron transfer is central to redox reactions, the energy release and momentum gain in an exothermic reaction like combustion are more complex than just electron acceleration. It's about the rearrangement of chemical bonds and the associated changes in potential energy.

Breaking bonds requires energy: To initiate combustion, you need to break the existing bonds in the reactants (e.g., the bonds between carbon and hydrogen in a fuel molecule and the bonds between oxygen atoms in oxygen molecules).   Forming bonds releases energy: The formation of new bonds in the products (e.g., the bonds between carbon and oxygen in carbon dioxide and the bonds between hydrogen and oxygen in water) releases energy.   Exothermic reactions: In an exothermic reaction, the energy released in forming new bonds is greater than the energy required to break the initial bonds.

This excess energy is released as EM radiation, or transferred to the vibrational, rotational, and translational modes of the product molecules. These internal modes of energy can also decay to lower energy levels emitting radiation. The photons emitted in all directions work as a radiative pressure when colliding with the molecules, with the net effect of this EM is to increasing the molecule's momentums.

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  • $\begingroup$ I appreciate the answer, but it's vague. How does electron transfer create em radiation? Photons were not captured. How does the energy get transferred to vibrational and rotational energy? Did the molecules get dragged together by the electrostatic potential? Are there articles, diagrams, or any other sources describing this? $\endgroup$
    – JustAnotherUser
    Commented Oct 7 at 13:41
  • $\begingroup$ Sometimes excited states of molecules spontaneously decay to a lower-energy state and emit a photon. This is called spontaneous emission. If you insist on trying to think classically (not quantumly), you can get a picture for the transfer to vibrational modes as happening because two point particles are pulled together by an electrostatic force, accelerating them as they go, so that when they are close enough to bond, they have momentum and energy to keep spinning around each other or oscillating back and forth. But the true story is quantum, so such a picture is very coarse/qualitative. $\endgroup$
    – Rokas Veitas
    Commented Oct 10 at 13:51

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