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

• zorro
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.
zorro
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?

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)

## 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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