6.1 Energy in Chemical Reactions

Energy, often in the form of heat, is either evolved or absorbed during a chemical reaction. This occurs because chemical reactions involve the breaking of old chemical bonds and the formation of new ones.

• Bond Breaking: Always consumes energy.
• Bond Making: Always releases energy.

The net energy change depends on the balance between these two processes.

• If energy released from bond formation is greater than energy consumed by bond breaking, the reaction releases energy to the surroundings (exothermic).
• If energy consumed by bond breaking is greater than energy released from bond formation, the reaction absorbs energy from the surroundings (endothermic).

Thus, in chemical reactions, there is always an energy transfer between the system (the reaction itself) and the surroundings.

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6.1.1 Units of Thermal Energy

The standard unit of heat or thermal energy in the SI system is the Joule (J).

Another common unit is the calorie (cal).

• Definition: The thermal energy required to raise the temperature of one gram of water from $14.5^{\circ }\mathrm{C}$ to $15.5^{\circ }\mathrm{C}$.
• Conversion: $1\text{calorie}=4.184\text{Joules}$

For larger quantities, the kilocalorie (kcal) and kilojoule (kJ) are commonly used: $1\text{kcal}=1000\text{cal}=4184\text{J}=4.184\text{kJ}$

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6.1.2 Thermochemical Reactions

A thermochemical reaction is a chemical reaction that proceeds with the evolution or absorption of heat. The study of these heat changes is called thermochemistry.

A thermochemical equation is a balanced chemical equation that includes the heat change (enthalpy change, $\Delta H$).

Example: ${\mathrm{C}}_{(\mathrm{s})}+{\mathrm{O}}_{2(\mathrm{g})}\to \mathrm{C}{\mathrm{O}}_{2(\mathrm{g})}\Delta H^{\circ }=-393.5\text{kJ}$

This equation indicates that when one mole of solid carbon reacts with one mole of oxygen gas to form one mole of carbon dioxide gas, 393.5 kJ of heat is released.

(1) Exothermic Reactions

An exothermic reaction is a chemical reaction that proceeds with the evolution of heat ($\Delta H$ is negative). In these reactions, the system transfers energy to the surroundings.

Examples:

• Combustion of Carbon: ${\mathrm{C}}_{(\mathrm{s})}+{\mathrm{O}}_{2(\mathrm{g})}\to \mathrm{C}{\mathrm{O}}_{2(\mathrm{g})}\Delta H^{\circ }=-393.5\text{kJ}$
• Formation of Water: $2{\mathrm{H}}_{2(\mathrm{g})}+{\mathrm{O}}_{2(\mathrm{g})}\to 2{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\Delta H^{\circ }=-571.6\text{kJ}$
• Formation of Carbon Monoxide: ${\mathrm{C}}_{(\mathrm{s})}+\frac{1}{2}{\mathrm{O}}_{2(\mathrm{g})}\to \mathrm{C}{\mathrm{O}}_{(\mathrm{g})}\Delta H^{\circ }=-110.5\text{kJ}$

(2) Endothermic Reactions

An endothermic reaction is a chemical reaction that proceeds with the absorption of heat ($\Delta H$ is positive). In these reactions, the system absorbs heat from the surroundings.

Examples:

• Formation of Hydrogen Iodide: ${\mathrm{H}}_{2(\mathrm{g})}+{\mathrm{I}}_{2(\mathrm{g})}\to 2\mathrm{H}{\mathrm{I}}_{(\mathrm{g})}\Delta H^{\circ }=+53.8\text{kJ}$
• Steam Reforming of Carbon: ${\mathrm{C}}_{(\mathrm{s})}+{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{g})}\to \mathrm{C}{\mathrm{O}}_{(\mathrm{g})}+{\mathrm{H}}_{2(\mathrm{g})}\Delta H^{\circ }=+131.4\text{kJ}$
• Formation of Nitric Oxide: ${\mathrm{N}}_{2(\mathrm{g})}+{\mathrm{O}}_{2(\mathrm{g})}\to 2\mathrm{N}{\mathrm{O}}_{(\mathrm{g})}\Delta H^{\circ }=+180.5\text{kJ}$

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6.1.3 Heat of Reaction ($\Delta H$)

The heat of reaction is the amount of heat evolved or absorbed in a chemical reaction, corresponding to the molar quantities shown in the balanced chemical equation.

The standard enthalpy change ($\Delta H^{\circ }$) is the heat of reaction measured under standard conditions:

• Temperature: $25^{\circ }\mathrm{C}$ (298 K)
• Pressure: 1 atmosphere (atm)

Important Note: If a reaction is reversed, the magnitude of $\Delta H^{\circ }$ remains the same, but its sign changes.

• Forward Reaction (Exothermic): ${\mathrm{C}}_{(\mathrm{s})}+{\mathrm{O}}_{2(\mathrm{g})}\to \mathrm{C}{\mathrm{O}}_{2(\mathrm{g})}\Delta H^{\circ }=-393.5\text{kJ}$
• Reverse Reaction (Endothermic): $\mathrm{C}{\mathrm{O}}_{2(\mathrm{g})}\to {\mathrm{C}}_{(\mathrm{s})}+{\mathrm{O}}_{2(\mathrm{g})}\Delta H^{\circ }=+393.5\text{kJ}$

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6.1.4 Relation between Enthalpy Change and Heat of Reaction

For reactions occurring at constant pressure, the heat change is equal to the change in enthalpy ($\Delta H$).

Enthalpy change is the difference between the enthalpies of the products and the reactants: $\Delta H=H_{\text{products}}-H_{\text{reactants}}$

• For an exothermic process, heat is released, so $H_{\text{products}}<H_{\text{reactants}}$, and $\Delta H$ is negative.
• For an endothermic process, heat is absorbed, so $H_{\text{products}}>H_{\text{reactants}}$, and $\Delta H$ is positive.

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6.1.5 Standard States and Standard Enthalpy Changes

To ensure consistency, enthalpy changes are reported for reactions where all substances are in their standard state.

Conditions for Standard States:

1. Gas: A pure gas at 1 atm pressure.
2. Element/Compound: The most stable physical state (solid, liquid, or gas) at 1 atm and $25^{\circ }\mathrm{C}$ (298 K).
3. Aqueous Solution: A substance at a 1 M concentration.

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6.1.6 Types of Standard Enthalpies

1. Standard Enthalpy of Reaction ($\Delta H_r^{\circ }$)

The enthalpy change when molar quantities of reactants as shown in the balanced equation react to form products, with all substances in their standard states.

Example: Methane Combustion $\mathrm{C}{\mathrm{H}}_{4(\mathrm{g})}+2{\mathrm{O}}_{2(\mathrm{g})}\to \mathrm{C}{\mathrm{O}}_{2(\mathrm{g})}+2{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\Delta H_r^{\circ }=-890.4\text{kJ}$

2. Standard Enthalpy of Formation ($\Delta H_f^{\circ }$)

The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states.

Convention: The standard enthalpy of formation of an element in its most stable standard state is zero ($\Delta H_f^{\circ }=0$).

Examples: ${\mathrm{C}}_{(\mathrm{s})}+{\mathrm{O}}_{2(\mathrm{g})}\to \mathrm{C}{\mathrm{O}}_{2(\mathrm{g})}\Delta H_f^{\circ }=-393.5\text{kJ/mol}$ ${\mathrm{H}}_{2(\mathrm{g})}+\frac{1}{2}{\mathrm{O}}_{2(\mathrm{g})}\to {\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\Delta H_f^{\circ }=-285.8\text{kJ/mol}$

3. Standard Enthalpy of Combustion ($\Delta H_c^{\circ }$)

The enthalpy change when one mole of a substance is completely burned in excess oxygen under standard conditions.

Example: $\mathrm{C}{\mathrm{H}}_{4(\mathrm{g})}+2{\mathrm{O}}_{2(\mathrm{g})}\to \mathrm{C}{\mathrm{O}}_{2(\mathrm{g})}+2{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\Delta H_c^{\circ }=-890.4\text{kJ/mol}$

4. Standard Enthalpy of Neutralization ($\Delta H_n^{\circ }$)

The enthalpy change when an acid and a base react to form one mole of water under standard conditions.

Example: $\mathrm{HC}{\mathrm{l}}_{(\mathrm{aq})}+\mathrm{NaO}{\mathrm{H}}_{(\mathrm{aq})}\to \mathrm{NaC}{\mathrm{l}}_{(\mathrm{aq})}+{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\Delta H_n^{\circ }=-57.3\text{kJ/mol}$

For strong acid–strong base neutralization, $\Delta H_n^{\circ }\approx -57.3\text{kJ/mol}$ because the net ionic equation is always: ${\mathrm{H}}_{(\mathrm{aq})}^{+}+\mathrm{O}{\mathrm{H}}_{(\mathrm{aq})}^{-}\to {\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}$