Difference between 'q' and 'ΔH' in thermochemistry?

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The discussion clarifies the distinction between 'q' (heat transfer) and 'ΔH' (change in enthalpy) in thermochemistry. 'q' specifically refers to the amount of heat transferred, while 'ΔH' encompasses the total energy change in a system, including heat, pressure/volume work, and entropy. For ΔH to equal q, conditions must be met where there is no change in pressure or volume and entropy is neglected. Additionally, the relationship between Gibbs' free energy (ΔG) and these terms is highlighted, with ΔG representing the useful energy change of the system under constant temperature and pressure, calculated as ΔH minus entropy.
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What is the basic difference between 'q' and 'ΔH' in thermochemistry? I get confused between them!
Is there any criteria for ΔH to become equal to q?
 
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Abdul Quadeer said:
What is the basic difference between 'q' and 'ΔH' in thermochemistry? I get confused between them! Is there any criteria for ΔH to become equal to q?

Q is an amount of heat being transferred and only heat. The change in enthalpy (ΔH), is the change in total energy of the system. That includes heat, but also pressure/volume work and entropy.

Gibbs' free energy, ΔG, is a measure of the change of the useful (i.e. work-producing) energy of the system, given no change in temperature or pressure. So it's simply the enthalpy minus the entropy.

So the heat transferred to a system in a reaction, Q, equals ΔH only if there is no change in pressure/volume or entropy. (You also neglect how the change in temperature caused by the heat from ΔH changes ΔH itself)
 
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