Final answer:
The statement is false; while the net heat exchange with the surroundings is intended to be zero in a calorimeter, it does not imply that the total enthalpy is unchanged -- rather that heat changes within the system are balanced by the calorimeter's contents.
Step-by-step explanation:
The statement that the total enthalpy of the calorimeter plus contents is unchanged during your experiment, thus deltaH = 0, is False. In the context of calorimetry, deltaH represents the change in enthalpy or total heat content of the system during a chemical reaction. According to the principles of calorimetry, the amount of heat released or absorbed by the reaction (the "system") will be equal in magnitude but opposite in sign to the heat absorbed or released by the surroundings, such as the water in the calorimeter. Therefore, while the net heat exchange with the surroundings is intended to be zero to prevent losses, this does not mean that the total enthalpy of the system and the calorimeter (surroundings) remains unchanged; it simply means that the heat released or absorbed by the system is accounted for by an equal and opposite change in the calorimeter's contents.
An ideal calorimeter is designed to minimize any energy exchange with the external environment, meaning that it is insulated well enough to prevent heat transfer with its surroundings outside of the calorimeter. In practice, some heat may be gained or lost in an actual calorimeter, but this is often negligible. The important aspect is that within the isolated system of the calorimeter, energy conservation is observed: the heat produced by the reaction (qreaction) plus the heat absorbed by the solution (qsolution) equals zero (qreaction + qsolution = 0).