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Enthalpy and activation energy statements

  • #1

Homework Statement


For a particular reversible reaction, the forward process is exothermic and the reverse process is endothermic. Which of the following statements must be true about this reaction?
(A) The forward reaction will be spontaneous under standard conditions
(B) The reverse reaction will be spontaneous under standard conditions
(C) The activation energy will be greater for the forward reaction than for the reverse reaction
(D) The activation energy will be greater for the reverse reaction than for the forward reaction.

Homework Equations


NA

The Attempt at a Solution


I eliminated B because if the reverse is endothermic, this will not be spontaneous. I chose the right answer (D) but I wanted to be sure my logic is on the right track. So, I assumed D must be correct because endothermic reactions require much more energy to force a reaction, whereas the forward exothermic reaction would require less.
 

Answers and Replies

  • #2
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  • #3
A is wrong because it usually takes some energy input to make the reactants reach activation energy, which will then lead to products
 
  • #4
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"A is wrong," is a correct statement, but the activation energy argument doesn't hold water. There are reactions which may be strongly exothermic, but the Gibb's free energy will still be positive due to entropy of the reaction, and which will not be spontaneous (condensation of water vapor, freezing of liquid water, other such things). Enthalpy is a good clue for spontaneity, but is not sufficient of itself to guarantee spontaneity.

"B" you will have deduced can be spontaneous for an argument opposite to what I just gave you for "A," and since we're looking for "must be true," is also discarded.

So, now, what can you tell me about "C?"
 

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