When 1 mol of a fuel burns at constant pressure, it produces 3452 kJ of heat and does 11 kJ of work. What are ΔE and ΔH for the combustion of the fuel?

Verified step by step guidance

1. Identify the given values: heat produced (q) = -3452 kJ and work done (w) = -11 kJ. Note that both values are negative because the system is losing energy.
2. Recall the first law of thermodynamics, which states that the change in internal energy (ΔE) is given by ΔE = q + w.
3. Substitute the given values into the equation: ΔE = (-3452 kJ) + (-11 kJ).
4. Recognize that at constant pressure, the change in enthalpy (ΔH) is equal to the heat exchanged, so ΔH = q.
5. Substitute the given value for heat into the equation for ΔH: ΔH = -3452 kJ.

Key Concepts

First Law of Thermodynamics

The First Law of Thermodynamics states that energy cannot be created or destroyed, only transformed from one form to another. In the context of chemical reactions, this principle implies that the change in internal energy (ΔE) of a system is equal to the heat added to the system minus the work done by the system on its surroundings.

Enthalpy (ΔH)

Enthalpy (ΔH) is a thermodynamic quantity that represents the total heat content of a system at constant pressure. It accounts for the internal energy of the system plus the product of its pressure and volume. For reactions occurring at constant pressure, the change in enthalpy is equal to the heat absorbed or released during the reaction.

Work in Thermodynamics

In thermodynamics, work refers to the energy transferred when a force is applied over a distance. For chemical reactions, work can be done by the system on the surroundings, such as when gases expand against external pressure. The work done affects the internal energy change and must be considered when calculating the overall energy changes in a reaction.