Thermochemistry

Thermochemistry

Thermochemistry is a branch of thermodynamics that deals with the relationship between chemical reactions and energy changes involving heat. It focuses on the study of energy transfer during physical and chemical processes.

Fundamental Concepts

Before diving into the specifics of thermochemistry, it's important to understand some fundamental concepts:

  • System and Surroundings: The part of the universe that is under study is called the system, and everything else is the surroundings.
  • Internal Energy (U): The total energy contained within a system.
  • Enthalpy (H): A measure of the total energy of a thermodynamic system, defined as ( H = U + PV ), where ( P ) is pressure and ( V ) is volume.
  • Heat (q): The transfer of energy due to a temperature difference.
  • Work (w): The transfer of energy that results from a force acting through a distance.

First Law of Thermodynamics

The first law of thermodynamics, also known as the law of energy conservation, states that energy cannot be created or destroyed in an isolated system. The change in internal energy of a system is equal to the heat added to the system minus the work done by the system on its surroundings.

[ \Delta U = q - w ]

Enthalpy Change

The enthalpy change (( \Delta H )) of a reaction is the heat change at constant pressure. It is a useful quantity in thermochemistry because most chemical reactions occur at atmospheric pressure.

  • Exothermic Reactions: Reactions that release heat (( \Delta H < 0 )).
  • Endothermic Reactions: Reactions that absorb heat (( \Delta H > 0 )).

Standard Enthalpy of Formation

The standard enthalpy of formation (( \Delta H_f^\circ )) is the heat change that results when one mole of a compound is formed from its elements in their standard states.

Hess's Law

Hess's Law states that the total enthalpy change for a reaction is the same, regardless of the number of steps in which the reaction is carried out. This allows us to calculate the enthalpy change for a reaction by summing the enthalpy changes of individual steps.

Calorimetry

Calorimetry is the measurement of heat flow. A calorimeter is a device used to measure the amount of heat involved in a chemical or physical process.

Table of Differences and Important Points

Property Exothermic Reaction Endothermic Reaction
Heat Flow Heat is released to the surroundings Heat is absorbed from the surroundings
Sign of ( \Delta H ) Negative (( \Delta H < 0 )) Positive (( \Delta H > 0 ))
Temperature Change Surroundings become warmer Surroundings become cooler
Example Combustion of gasoline Photosynthesis

Formulas in Thermochemistry

  • Enthalpy Change: ( \Delta H = H_{products} - H_{reactants} )
  • Heat Capacity (C): The amount of heat required to raise the temperature of a substance by 1 degree Celsius. ( q = C \Delta T )
  • Specific Heat (c): The heat capacity per unit mass. ( q = mc\Delta T )
  • Hess's Law: ( \Delta H_{reaction} = \sum \Delta H_{products} - \sum \Delta H_{reactants} )

Examples

Example 1: Exothermic Reaction

The combustion of methane (( CH_4 )) is an exothermic reaction:

[ CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(l) \quad \Delta H = -890 \, kJ/mol ]

The negative sign indicates that heat is released.

Example 2: Endothermic Reaction

The decomposition of calcium carbonate is an endothermic reaction:

[ CaCO_3(s) \rightarrow CaO(s) + CO_2(g) \quad \Delta H = +178 \, kJ/mol ]

The positive sign indicates that heat is absorbed.

Example 3: Using Hess's Law

To find the enthalpy change for the reaction:

[ N_2(g) + O_2(g) \rightarrow 2NO(g) ]

We can use the following reactions and their enthalpy changes:

  1. ( 1/2 N_2(g) + O_2(g) \rightarrow NO_2(g) \quad \Delta H_1 = +33.2 \, kJ )
  2. ( NO_2(g) \rightarrow NO(g) + 1/2 O_2(g) \quad \Delta H_2 = +56.6 \, kJ )

By Hess's Law:

[ \Delta H_{reaction} = \Delta H_1 + \Delta H_2 ] [ \Delta H_{reaction} = +33.2 \, kJ + 2(+56.6 \, kJ) ] [ \Delta H_{reaction} = +146.4 \, kJ ]

Thermochemistry is essential for understanding energy changes in chemical reactions, which is crucial for predicting reaction behavior, designing energy-efficient processes, and understanding environmental impact.