# Relationship between enthalpy change and endo/exothermic reactions

• Something that confused me, is there a relationship between enthalpy change and endo/exothermic? Does region "x" represent the enthalpy change?
• What represents enthalpy change, kinectic energy, activation energy and potential energy on the graph below?
• But why must, for an endothermic reaction, the value of activation energy be higher than the enthalpy change? Then, how about the backwards in equillibrium, if the forward reaction is exothermic, is it the same?
• And why is, for an exothermic reaction, the value of the activation energy equal to the enthalpy change? Is that a must? Can the activation be energy greater than the enthalpy change? I will attempt to answer all your questions however I don't completely understand some of them. Could you perhaps reword them or explain them a bit more?

Question 1

is there a relationship between enthalpy change and endo/exothermic

Of course there is! What determines if a reaction is endothermic or exothermic is by looking at the sign of the $\Delta \mathrm{H}$ (enthalpy change). If $\Delta \mathrm{H}$ is negative, then the reaction is exothermic and when $\Delta \mathrm{H}$ is positive, the reaction is endothermic.

does x represent the enthalpy change

That is also correct as enthalpy change is a state function. In other words, it doesn't mater what happens in the middle of the graph, the enthalpy change is just the final value minus initial value.

Question 2

What represents enthalpy change, kinetic energy, activation energy and potential energy on the graph below?

Enthalpy change- as already discussed, x represents the enthalpy change

Activation Energy- can be seen as the difference between the maximum potential energy and initial potential energy. Hence it is w.

Potential energy- that is the y value at every point on the graph

I am not sure how you can find the kinetic energy. I suppose you can find the change in kinetic energy if assume that the potential energy completely converts to kinetic energy.

Question 3

why must for an endothermic reaction the value of activation energy be higher than the enthalpy change?

That is true for any reaction. That is because as I stated above, the activation energy is the difference between the maximum potential energy and initial potential energy. Hence there is no way the enthalpy change, the energy difference between the final and initial state, can be larger than the activation energy.

Then, how about the backwards in equilibrium, if the forward reaction is exothermic, is it the same?

I don't really understand what you mean here. Could you perhaps explain this more.

Question 4

And why is, for an exothermic reaction, the value of the activation energy equal to the enthalpy change? Is that a must? Can the activation be energy greater than the enthalpy change?

Like I have said above several times, the value of the activation energy is always greater than the enthalpy change.

• Thanks a lot , for question 3 , i mean that if there is an equillibrium , such that A + B ⇌ C -49kJ mol-1, (backward reaction is endothermic ), is the value of activation energy be higher than the enthalpy change?But after looking your detail explaination , i get it , thx :) sorry for my poor english – Yin Ting Ng Mar 10 '16 at 13:28
• I said the same thing, and I was the first one to post your answer – user1825567 Mar 10 '16 at 17:17
• @YinTingNg No worries. I am glad I was able to help. – Nanoputian Mar 10 '16 at 19:47

Yes the region x represents Enthalpy change.The relation between Enthalpy change and endothermic reaction is that if a reaction is endothermic then Enthalpy change is +ve and if the reaction is exothermic then the Enthalpy change is -ve.

In an endothermic reaction the amount of energy you supply is lesser that the amount of energy released when product is formed so the amount of energy you supply to transform the reactants to activated complex will be more than the energy released when the activated complex transforms to product.

In the graph you have provided the reaction in forward direction is exothermic and in backward direction is endothermic. Whichever direction you take activation energy will always be more than the Enthalpy change because Enthalpy change is the difference in the Enthalpy of the products and the reactants whereas the activation energy is the difference in the Enthalpy of the activated complex and the reactants. The Enthalpy of the activated complex is always more than the Enthalpy of the products and the reactants so both these quantities cannot be equal.

For both exothermic as well as endothermic reactions it not a necessary condition for the activation energy to be equal to the Enthalpy change as activation energy is always greater than the Enthalpy change.