Change in Enthalpy vs Q: State Functions Explained

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Q represents heat transfer and is classified as a non-state function, while the change in enthalpy (dH) is a state function. The equation for heat transfer, Q = P*dV + dE, indicates that Q depends on the path taken, whereas dH = P*dv + dE reflects a property of the system that depends only on its current state. The confusion arises from the relationship between Q and dH, as they can be equal under certain conditions, yet their classifications differ. It is important to understand that while enthalpy itself is a state function, the change in enthalpy (dH) is also a state function, as it is determined solely by the initial and final states of the system. Thus, dH is indeed a state function, despite the non-state nature of Q.
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as Q(heat transfer)=P*dV + dE
Q is a non state function,
while the change of enthalpy,
dH=P*dv+dE,
i suppose it is equal to the Q above but why dH is a state function while Q is not since they are the same?
 
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The enthalpy, H, is a state function. Is that necessarily true of change of enthalpy?
 
so u mean dH is not a state function?
 

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