Calculating Enthalpy Changes at Different Temperatures
@Chet already gave the qualitative answer but the actual equation it is based on is:

$$
\Delta H^\circ(T)=\Delta H{^\circ}_{T_1} + \int_{T_1}^{T} \Delta C_p{^\circ}(T^\prime ) \,\mathrm{d}T^\prime$$

If you know the enthalpy at a reference temperature and the heat capacity as a function of temperature you can calculate the enthalpy change at any other temperature.

Actually d) and e) together are not wrong. We all know these equations:
$$\Delta G= - RT\ln (K) \qquad \Delta G= \Delta H - T\Delta S $$

With d) you can calculate $\Delta G$ and with e) you can calculate $\Delta H$.

Lets look deeper into: $\Delta G^\circ$, $\Delta G$, $K$ and $Q$. 
To answer your second question completely I would recommend to look again into a textbook of your choice. Especially the definitions of activity and chemical potential.

If you know the chemical potential $\mu^\circ$ of a pure substance you can calculate the chemical potential of mixtures if you know the activities:
$$\mu= \mu^\circ + RT \ln(a)$$

Now we make two assumptions: 


We have constant temperature and pressure. ($\mu = \Delta G$)
We can calculate the activities simply over the partial pressures or concentrations.


This leads then to your equation:
$$
\Delta G= \Delta G^\circ + RT \ln(Q)
$$
where $Q$ is the reaction quotient.
Since we assume equilibrium,  we know that $\Delta G=0$ and $Q = K$. This leads to:
$$ 0= \Delta G^\circ + RT \ln(K) \Rightarrow \Delta G^\circ = -RT \ln(K) $$

The last equation is the handy one you use for calculating equilibrium constants from energies and vice versa.
So to be really precise one has to write:
$$ \Delta G^\circ= \Delta H^\circ - T\Delta S^\circ $$

But one usually skips the superscript $^\circ$.
At equilibrium $Q$ and $K$ are essentially the same.