Define enthalpy change (ΔH) and explain the sign conventions for exothermic and endothermic reactions.

Define enthalpy change (ΔH) and explain the sign conventions for exothermic and endothermic reactions.

• A. ΔH is heat transfer at constant pressure; negative for exothermic, positive for endothermic.
• B. ΔH is heat transfer that occurs only at constant volume.
• C. ΔH is total energy change, independent of pressure.
• D. ΔH is enthalpy change that always decreases for exothermic reactions.

Answer: (A)

The main idea being tested is how enthalpy change is defined and how its sign relates to heat flow in a reaction. Enthalpy, H, is a measure of the heat content of a system that, for processes at constant pressure, corresponds to the heat exchanged with the surroundings. Therefore, the enthalpy change ΔH tells you how much heat is absorbed or released when the system goes from one state to another at constant pressure. If the system releases heat to the surroundings, the heat flow q at constant pressure is negative, so ΔH is negative. If the system absorbs heat, q is positive, so ΔH is positive. That’s why stating that ΔH is heat transfer at constant pressure and that its sign is negative for exothermic and positive for endothermic reactions is the most accurate description. A few details help connect this to the bigger picture: since H = U + pV, the change in enthalpy ΔH can be related to changes in internal energy and in the pV term. At constant pressure, the heat added to the system equals ΔH, ignoring other work terms, which is why q_p = ΔH in that case. This is a practical reason scientists measure enthalpy changes in calorimetry experiments conducted at roughly constant pressure. The other statements aren’t quite right because they either limit heat transfer to constant volume, mischaracterize what enthalpy change represents, or mishandle the sign interpretation. Heat transfer at constant volume is tied to the internal energy change, not ΔH. Enthalpy change is not the total energy change independent of pressure; it includes the pV contribution and depends on pressure conditions. And saying exothermic enthalpy changes always decrease is misleading—the sign indicates heat flow direction, not a monotonically decreasing value over time.

The main idea being tested is how enthalpy change is defined and how its sign relates to heat flow in a reaction. Enthalpy, H, is a measure of the heat content of a system that, for processes at constant pressure, corresponds to the heat exchanged with the surroundings. Therefore, the enthalpy change ΔH tells you how much heat is absorbed or released when the system goes from one state to another at constant pressure. If the system releases heat to the surroundings, the heat flow q at constant pressure is negative, so ΔH is negative. If the system absorbs heat, q is positive, so ΔH is positive. That’s why stating that ΔH is heat transfer at constant pressure and that its sign is negative for exothermic and positive for endothermic reactions is the most accurate description.

A few details help connect this to the bigger picture: since H = U + pV, the change in enthalpy ΔH can be related to changes in internal energy and in the pV term. At constant pressure, the heat added to the system equals ΔH, ignoring other work terms, which is why q_p = ΔH in that case. This is a practical reason scientists measure enthalpy changes in calorimetry experiments conducted at roughly constant pressure.

The other statements aren’t quite right because they either limit heat transfer to constant volume, mischaracterize what enthalpy change represents, or mishandle the sign interpretation. Heat transfer at constant volume is tied to the internal energy change, not ΔH. Enthalpy change is not the total energy change independent of pressure; it includes the pV contribution and depends on pressure conditions. And saying exothermic enthalpy changes always decrease is misleading—the sign indicates heat flow direction, not a monotonically decreasing value over time.