Entropy of a chemical reaction

[erisedk]

Homework Statement

The entropy change of surrounding is greater than that of the system in a/an
(A) exothermic process
(B) endothermic process
(C) both (A) and (B) are correct
(D) none of these are correct

Homework Equations

The Attempt at a Solution

ΔS_tot = ΔS_sys + ΔS_surr
For a spontaneous process, ΔS_tot > 0
For an exothermic process, ΔH < 0 and for an endothermic process ΔH > 0
I don't really know what to do now. Please help.

 

[Ygggdrasil]

Try looking up a formula that relates $\Delta S_{surr}$ to $\Delta H$. Remember that for chemical reactions performed at constant pressure, the change in enthalpy tells you the amount of heat transferred between the system and surroundings.

 

[erisedk]

ΔG = ∆H_sys-T∆S_sys
At equilibrium ∆G = 0
So ∆S_sys = ∆H_sys/T
For exothermic reactions, ∆H_sys < 0 so ∆S_sys < 0
For ∆S_sys + ∆S_surr > 0
∆S_surr > 0
Since a positive number is always greater than a negative number, ∆S_surr > ∆S_sys
So A is correct.
For endothermic reactions, ∆S_sys > 0 as ∆H_sys > 0
Now for ∆S_sys + ∆S_surr > 0
We can have ∆S_surr as
1. positive and larger than system entropy
2. Positive and smaller than system entropy
3. Negative but smaller in magnitude than system entropy.

So which of these would be appropriate?

 

[haruspex]

Since a positive number is always greater than a negative number, ∆Ssurr > ∆Ssys

The question asks which entropy change is the greater. I would interpret that in terms of magnitude.

 

[Ygggdrasil]

A useful formula here is to remember that the entropy change of the surroundings depends on the amount of heat transferred to/from the system: $\Delta S_{surr} = - \Delta H/T$

With this in mind (and the fact that for any spontaneous process $\Delta G < 0$), how do the magnitudes of $\Delta S_{sys}$ and $\Delta S_{surr}$ compare in the two cases?