Second Law of Thermodynamics

According to the first law of thermodynamics, the energy change in a system is equal and opposite to the energy change of the surroundings.

When an ice cube, the system, is added to a cup of hot tea, the surroundings, the ice melts while the tea becomes cooler. The heat gained by the ice cube is equal to the heat lost from the tea. Energy is conserved no matter the direction of heat transfer.

However, adding an ice cube would never make the tea hotter because the amount of heat transferred does not determine which way the heat flows.

The associated change in entropy must be considered to explain the direction of heat transfer and other spontaneous reactions.

The second law of thermodynamics states that the entropy of the universe, which is the total entropy of both the system and surroundings, increases for all spontaneous processes. This means that the ΔS of the universe, the difference between the entropy of the universe’s final and initial states, must be greater than zero.

As entropy is a measure of energy dispersal, a process where the energy of the universe is more dispersed in the final state than in the initial will be spontaneous.

When an ice cube melts, the water molecule