Internal Energy

In a chemical process, the difference between the internal energies of reactants and products, represented by ΔU, is used to determine whether the system gained or lost energy during the reaction.

If ΔU is greater than zero, the final internal energy was higher than the initial internal energy and the system gained energy during the reaction.

If ΔU is less than zero, the final internal energy is lower than the initial internal energy, meaning the system lost energy.

According to the first law of thermodynamics, any change in a system’s energy must be balanced by an equal and opposite change in its surroundings.

Thus, the change in a system’s internal energy is equal to the energy transferred as heat, symbolized by “q”, plus the energy transferred as work symbolized by “w”, during the process.

In chemistry, the signs for heat and work depend on whether the system gains or loses energy.

Consider the transformation of carbon dioxide to elemental carbon and oxygen. The reactant has lower internal energy than the products, meaning ΔU is positive. Energy is transferred from the surroundings to the system increasing its internal energy.

Alternatively, during the reaction between sulfur and oxygen gas to make sulfur dioxide, energy is transferred to the surroundings. Here, the products have a lower internal energy than the reactants, and ΔU is negative.

Hence, ΔU only depends on the initial and final internal energy states of the system and the magnitude of heat and work exchanged with the surroundings.