Difference Between Enthalpy and Internal Energy

Enthalpy vs Internal Energy

For the study purposes in chemistry, we divide the universe into two as a system and surrounding. At any time, the part we are interested is the system, and the rest is surrounding. Enthalpy and internal energy are two concepts related to the first law of thermodynamics, and they describe the reactions taking place in a system and the surrounding.

What is Enthalpy?

When a reaction takes place, it may absorb or evolve heat, and if the reaction is carried at constant pressure, this heat is called the enthalpy of the reaction. Enthalpy of molecules cannot be measured. Therefore, change in enthalpy during a reaction is measured. The enthalpy change (∆H) for a reaction in a given temperature and pressure is obtained by subtracting the enthalpy of reactants from the enthalpy of products. If this value is negative, then the reaction is exothermic. If the value is positive, then the reaction is said to be endothermic. The change in enthalpy between any pair of reactants and products is independent of the path between them. Moreover, enthalpy change depends on the phase of the reactants. For example, when the oxygen and hydrogen gases react to produce water vapor, the enthalpy change is -483.7 kJ. However, when the same reactants react to produce liquid water, the enthalpy change is -571.5 kJ.

2H₂ (g) +O₂ (g) → 2H₂O (g); ∆H= -483.7 kJ

2H₂ (g) +O₂ (g) → 2H₂O (l); ∆H= -571.7 kJ

What is Internal energy?

Heat and work are two ways of transferring energy. In mechanical processes, energy may be transferred from one place to another, but the total quantity of energy is conserved. In chemical transformations, a similar principle applies. Consider a reaction such as the combustion of methane.

CH₄ + 2 O₂ → CO₂ + 2 H₂O

If the reaction takes place in a sealed container, all that happens is that heat is released. We could use this released enzyme to do mechanical work such as run a turbine or steam engine, etc. There is an infinite number of ways that the energy produced by the reaction could be divided up between heat and work. However, it is found that the sum of the heat evolved and the mechanical work done is always a constant. This leads to the idea that in going from reactants to products, there is some property called, the internal energy (U). The change of internal energy is denoted as ∆U.

∆U= q + w; where q is the heat and w is the work done

The internal energy is called a state function as its value depends on the state of the system and not how the system came to be in that state. That is, the change in U, when going from the initial state “i” to final state “f”, depends only on the values of U in the initial and final states.

∆U= U_f – U_i

According to the first law of thermodynamics, the internal energy change of an isolated system is zero. Universe is an isolated system; therefore, ∆U for the universe is zero.