In physical chemistry, it is often stated that "at constant pressure, the change in enthalpy gives the amount of heat transferred into the system" and the justification for this claim is
dH = d(U+PV) = dU + PdV = \delta q + \delta w + PdV = \delta q
because \delta w = -PdV (we are also assuming only reversible p-V work). I understand this argument and why it must be true for a gas system in which the (well-defined) total pressure is being held constant (even if partial pressures of the individual components are changing). However, my textbook (Moore Physical Chemistry) makes use of this fact in discussing thermochemistry of reactions involving solids and liquids as well as gases. It is stated that in the lab, such reactions commonly occur under constant atmospheric pressure and therefore, we have q_P = \Delta H. I am unsure how to interpret this constant pressure condition for liquids and solids. For liquids, I agree that the pressure at the surface must equal the atmospheric pressure, but liquids generally have much greater pressure variation with depth than gases and we cannot, in general, assume that the pressure is relatively uniform. Furthermore, I don't see how we can even speak of the pressure of a solid as pressure is usually defined as force per unit area exerted by a fluid. I am hoping someone can explain how to carefully interpret pressure and enthalpy for reactions involving solids and liquids. Thank you in advance!
