I came across a textbook that stated a combination process that required heat. There was no explanation whether it was a prerequisite for a combination process to have heat or it was just an example just that it was grouped under "heat process chemical changes".

Was doing a test paper that asked about 2 differences between thermal decomposition and combination processes and one of the difference was that "combination processes do not need heat but thermal decomposition processes do".

Would like to ask which interpretation is correct. From preliminary searches on Google it seems that the consensus is that any reaction that simply involves two separate substances combining in a reaction to become a single product is considered to be a combination reaction regardless of whether there is heat involved or not.

Would like to know what is the correct interpretation.

  • $\begingroup$ By combination process, you mean like: $$\ce{A + B -> C}$$ $\endgroup$ Jun 1 '17 at 3:51
  • $\begingroup$ Hi! Thanks for the reply. Yes, I do mean: A+B⟶C. However, I would like to know if there needs to be heat involved in the reaction by definition. $\endgroup$
    – truetoall
    Jun 1 '17 at 3:53

Definitely. Although combination processes may be exothermic, they do have activational energy. As an example, consider the reaction:

$$\ce{H2 + I2 -> 2HI}$$

While the formation of two $\ce{H-I}$ bonds is quite energy releasing, you still have to go through breaking a $\ce{H-H}$ and a $\ce{I-I}$ bond. You need to supply heat to give thermal agitation to allow the molecules with sufficient energy to break the bonds. A typical reaction like this one would have this energy-reaction coordinate graph:

enter image description here

The reaction proceeds into a high-energy transition state, and then comes down to the low energy products. To reach the transition state however, you must supply heat.

Edit: As @porphyrin mentioned, there are some reactions that have essentially zero activational energy. Combination of radicals may be an example:

$$\ce{Cl. + Cl. -> Cl2}$$

This one has no activational energy because there is no unstable transition state involved. Another popular example is the basic neutralization reaction:

$$\ce{H+ + OH- <=> H2O}$$

Here again, there isn't any highly unstable transition state.

  • $\begingroup$ Thank you for the answer! Appreciate the detailed explanation! $\endgroup$
    – truetoall
    Jun 1 '17 at 5:36
  • $\begingroup$ @truetoall If this answer completely satisfies you, do click the ✅ button to accept it. This lets others know this question is solved. $\endgroup$ Jun 1 '17 at 5:54
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    $\begingroup$ As it happens the example you chose is a bit misleading as this reaction is not a simple bimolecular one but involves H and I radicals. You could point out also that endothermic reactions have an activation energy and some reactions effectively have zero activation energy. $\endgroup$
    – porphyrin
    Jun 1 '17 at 10:15

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