When doing constant volume calorimetry, the enthalpy can be calculated as follows: ΔH = ΔU + Δ(PV) ΔH = w + q + Δ(PV) ΔH = PΔV + q + Δ(PV) and at constant volume: ΔH = q + VΔP which I've then see people rewrite using the ideal gas law as follows: ΔH = q + (Δn)RT where Δn is the change in the moles of gas and T is constant. This is what I don't understand. Why is T constant? If you're doing calorimetry, the temperature is changing. Why are we now assuming that it is constant? Before looking it up, I originally had the following: ΔH = q + nRΔT Why isn't it this? Or even ΔH = q + ΔnRΔT How do you know when to keep moles constant and when to keep the temperature constant? Thanks!