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In the textbook, it says "Because heat is not a state function, we should not speak of a system as possessing a certain amount of heat." (the same claim is said for work). But yet when we describe reactions, we assign it a quantitative value such as +750 kJ or 5 J, etc. Can someone explain this reasoning, to me it seems counterintuitive.
I believe the value assigned is the enthalpy of the reaction, which is a state function, and is different from heat. Also, joules and kilojoules are units of energy, which gives way to understanding that we are tracking the change in energy of the function, which is the enthalpy.
q is the general symbol we use for heat and it is a path dependent function. H is the general symbol for enthalpy and it has the definition of H = U + PV which is derived from the 1st law of thermodynamics. H is a state function while q is a path dependent function. This is why in many of the problems we are asked to calculate delta H but in almost none of the problems are we asked to calculate delta q.
Adding to the other posts, when we calculate the change in internal energy of a system, ∆U = Q + W, Q represents the energy that the system either loses or gains as heat, not the heat the system possesses.
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