Energy Changes in Chemical and Physical Processes

In chemical reactions, energy is either absorbed or released. This branch of chemistry is known as Thermochemistry. The total heat content of a system is called Enthalpy (H), and we measure the change in enthalpy (∆H) during a process.

1. Exothermic and Endothermic Reactions

Exothermic Reactions (-∆H): These are reactions where heat is released to the surroundings. The temperature of the surroundings increases. Examples include:

• Reaction of Sodium (Na) or Potassium (K) with water.
• Dissolving Sodium Hydroxide (NaOH) in water.
• Combustion of fuels.
• Neutralization reactions between acids and bases.

Endothermic Reactions (+∆H): These are reactions where heat is absorbed from the surroundings. The temperature of the surroundings decreases. Examples include:

• Dissolving Ammonium Chloride (NH4Cl) in water.
• Thermal decomposition of Calcium Carbonate (CaCO3).
• Photosynthesis.

2. Entropy (∆S)

Entropy is a measure of the degree of disorder or randomness in a system. The more disordered a system is, the higher its entropy.

• Solids: High order, low entropy.
• Liquids: Moderate disorder, moderate entropy.
• Gases: High disorder, high entropy.

Entropy increases when a solid melts, a liquid evaporates, or when a salt dissolves in water (as ions become free-moving).

3. Gibbs Free Energy (∆G) and Spontaneity

The spontaneity of a reaction (whether it can occur on its own) is determined by the relationship between Enthalpy (∆H), Entropy (∆S), and Temperature (T).

The Gibbs-Helmholtz Equation: ∆G = ∆H - T∆S

• If ∆G < 0 (Negative): The reaction is spontaneous.
• If ∆G > 0 (Positive): The reaction is non-spontaneous.
• If ∆G = 0: The system is at equilibrium.

4. Driving Forces for Reactions

Nature tends towards minimum energy (negative ∆H) and maximum disorder (positive ∆S). A reaction is always spontaneous if it is exothermic (-∆H) and results in increased disorder (+∆S).