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

In summary, the basic difference between 'q' and 'ΔH' in thermochemistry is that q is an amount of heat being transferred, while ΔH is the change in total energy of the system. ΔH also includes pressure/volume work and entropy, while q does not. Additionally, Gibbs' free energy, ΔG, is a measure of the useful energy of the system and is equal to ΔH minus the entropy. Q equals ΔH only when there is no change in pressure/volume or entropy.
  • #1
zorro
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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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  • #2


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)
 

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

1. What is the difference between 'q' and 'ΔH' in thermochemistry?

Q is the symbol used to represent heat, which is a form of energy. It is measured in Joules (J) or calories (cal). ΔH (delta H) represents the change in enthalpy, which is the total heat content of a system. It is also measured in Joules or calories.

2. How are 'q' and 'ΔH' related to each other in thermochemistry?

Q and ΔH are related through the equation Q = ΔH + ΔnRT, where Δn is the change in the number of moles of gas molecules, R is the gas constant, and T is the temperature in Kelvin. This equation shows that q is a component of ΔH and that they are both influenced by the same factors.

3. Are 'q' and 'ΔH' always positive in thermochemistry?

No, q and ΔH can be positive, negative, or even zero in thermochemistry. A positive value for q or ΔH indicates that the system is gaining heat or increasing in enthalpy, respectively. A negative value indicates that the system is losing heat or decreasing in enthalpy. A value of zero means that there is no change in heat or enthalpy.

4. How do 'q' and 'ΔH' affect a chemical reaction in thermochemistry?

The values of q and ΔH can tell us whether a reaction is endothermic or exothermic. An endothermic reaction absorbs heat from the surroundings, resulting in a positive value for q and ΔH. An exothermic reaction releases heat to the surroundings, resulting in a negative value for q and ΔH.

5. Can 'q' and 'ΔH' be used interchangeably in thermochemistry calculations?

No, q and ΔH cannot be used interchangeably. Q represents the heat exchanged between a system and its surroundings, while ΔH represents the overall change in enthalpy of a system. They have different units and are calculated using different equations, so they cannot be substituted for each other in calculations.

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