LU 3 THERMOCHEMISTRY LU 3: THERMOCHEMISTRY 3.1 Definitions Almost all chemical reactions absorbed or released energy. Heat is the form of energy most commonly transferred in the reactions. 1. Energy is the ability to do work or transfer heat. Energy can take many forms: mechanical, electrical, thermal, chemical or nuclear. It can be classified as kinetic energy or potential energy: Kinetic energy - energy available because of the motion of an object Potential energy - energy due to the position which is associated with forces of attraction and repulsion between objects. 2. All forms of energy can be changed from one to form another. When a ball starts to roll downhill, its potential energy is converted to kinetic energy. The total quantity of energy in the universe is thus assumed to remain constant and is neither be created nor destroyed. This statement is known as the law of conservation of energy.
3.2 Thermochemistry Thermochemistry (energetics) is the study of the relationship between energy (heat) and chemical reactions. Heat is exchanged between system and surroundings when a reaction occurred. There are two types of reactions as listed in Table 3.1: Exothermic reaction Endothermic reaction
A reaction where heat is given out to the A reaction where heat is absorbed from the surroundings which causes the temperature of the surroundings which causes the temperature of the reaction mixture to increase reaction mixture to decrease The enthalpy of the products is lower than reactants
The enthalpy of the products is higher than reactants
Energy C + D (products)
A + B (reactants) ∆H= negative C + D (products) A + B (reactants) ∆H= positive
The ∆H is negative Example 1: CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (l) H = - 890.4 kJ H2 (g) + O2 (g) H2O2 (g) ∆H = -186 kJ
The ∆H is positive Example 2: 6CO2 (g) + 6H2O (l) C6H12O6 (aq) + 6O2 (g) H = +2802.5 kJ H2O (l) H2 (g) + ½ O2 (g) ∆H = +287 kJ
Table 3.1: Exothermic and endothermic...