Chapter 7: Chemical Reactions—Energy, Rates, and Equilibrium (GOB Chemistry Study Notes)

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Chapter 7: Chemical Reactions—Energy, Rates, and Equilibrium

7.1 Energy and Chemical Bonds

Chemical bonds store energy, which can be classified as potential or kinetic. Understanding these forms of energy is essential for analyzing chemical reactions.

Potential Energy: Stored energy due to position or arrangement. Examples include water behind a dam, a poised car, or a coiled spring.
Kinetic Energy: Energy of motion, such as water falling over a dam or a rolling car.
When bonds form, potential energy is often converted to heat (kinetic energy of particles).
Breaking bonds requires energy input; forming bonds releases energy.
If products have less potential energy than reactants, they are more stable and less likely to undergo further change.

7.2 Heat Changes during Chemical Reactions

Heat changes in chemical reactions are governed by the breaking and forming of bonds, measured by bond dissociation energy.

Bond Dissociation Energy: The energy required to break a bond in a molecule. Higher values indicate more stable, less reactive bonds.
Endothermic Reaction: Absorbs heat (positive sign, energy in). Example: bond breaking.
Exothermic Reaction: Releases heat (negative sign, energy out). Example: bond formation.
The law of conservation of energy states that energy cannot be created or destroyed, only transformed.
Heat of Reaction (ΔH): The difference between heat absorbed in breaking bonds and heat released in forming bonds.

Table of average bond dissociation energies

7.3 Exothermic and Endothermic Reactions

Chemical reactions can be classified based on their heat exchange with the surroundings.

Exothermic: Products have stronger bonds than reactants; energy is released (ΔH negative).
Endothermic: More energy is required to break bonds than is released in forming new ones; energy is absorbed (ΔH positive).
The amount of heat absorbed or released in the reverse reaction is equal in magnitude but opposite in sign to the forward reaction.

Thermite reaction, a strongly exothermic process Exothermic reaction equation: combustion of methane Table of bond energies for reactants and products Endothermic reaction equation: formation of NO from N2 and O2

7.4 Why Do Chemical Reactions Occur? Free Energy

The spontaneity of a reaction depends on both energy changes and molecular disorder (entropy).

Spontaneous Process: Proceeds without external influence (e.g., water flowing downhill).
Entropy (S): Measure of molecular disorder; increases when a system becomes more random (e.g., solid to liquid to gas).
ΔS is positive if disorder increases, negative if disorder decreases.
Spontaneity is favored by heat release (ΔH negative) and increased entropy (ΔS positive).
Free Energy Change (ΔG): Determines spontaneity, combining enthalpy and entropy effects:

Waterfall illustrating spontaneous process Diagram showing entropy increases from solid to liquid to gas and with more moles Equation for free energy change

Exergonic: Spontaneous, ΔG negative (free energy released).
Endergonic: Nonspontaneous, ΔG positive (free energy absorbed).
Temperature can affect spontaneity; some reactions become spontaneous only at higher temperatures.

7.5 How Do Chemical Reactions Occur? Reaction Rates

Reaction rates depend on molecular collisions, orientation, and energy.

Reactant particles must collide with correct orientation and sufficient energy to break bonds and form new ones.
Activation Energy (Eact): Minimum energy required for a reaction to occur; determines reaction rate.
Lower activation energy means a faster reaction; higher activation energy means a slower reaction.
ΔG indicates if a reaction can occur, but not how fast it will occur.

Effective and ineffective molecular collisions Reaction energy diagrams for exergonic and endergonic reactions Energy diagram showing small activation energy and negative ΔG

7.6 Effects of Temperature, Concentration, and Catalysts on Reaction Rates

Several factors influence how quickly a reaction proceeds:

Temperature: Raising temperature increases particle energy, frequency, and force of collisions, thus increasing reaction rate. A 10°C rise typically doubles the rate.
Concentration: Higher concentration increases collision frequency, often increasing reaction rate proportionally.
Catalysts: Lower activation energy by providing an alternative pathway or orienting molecules, increasing reaction rate without being consumed.

Effect of temperature on reaction rate Effect of concentration on reaction rate Energy diagram showing effect of catalyst on activation energy

Change	Effect
Concentration	Increase in reactant concentration increases rate; decrease decreases rate.
Temperature	Increase in temperature increases rate; decrease decreases rate.
Catalyst added	Increases reaction rate.

7.7 Reversible Reactions and Chemical Equilibrium

Many reactions are reversible, proceeding in both directions until equilibrium is reached.

Reversible Reaction: Can proceed from reactants to products and vice versa.
Chemical Equilibrium: State where forward and reverse reaction rates are equal; concentrations of all species remain constant.
Equilibrium is dynamic—reactions continue, but no net change in concentrations occurs.

Reversible reaction: acetic acid and ethyl alcohol Graph showing rates of forward and reverse reactions reaching equilibrium

7.8 Equilibrium Equations and Equilibrium Constants

The position of equilibrium is described quantitatively by the equilibrium constant (K).

For a general reaction: aA + bB ⇌ mM + nN, the equilibrium constant is:

Equilibrium equation and constant

K indicates the extent of reaction: large K means more products, small K means more reactants.
Pure solids and liquids are omitted from K expressions.

Scale showing meaning of K values Equilibrium constant for steam reforming of methane Calculation of K for PCl5 equilibrium

7.9 Le Châtelier’s Principle: The Effect of Changing Conditions on Equilibria

Le Châtelier’s principle predicts how a system at equilibrium responds to disturbances (stresses).

Stress: Any change in concentration, temperature, or pressure that disturbs equilibrium.
Adding reactant or removing product shifts equilibrium to the right (toward products).
Adding product or removing reactant shifts equilibrium to the left (toward reactants).
Increasing temperature favors the endothermic direction; decreasing temperature favors the exothermic direction.
Increasing pressure (by decreasing volume) favors the side with fewer moles of gas.
Catalysts do not affect the position of equilibrium, only the rate at which it is reached.

Adding reactants shifts equilibrium right Adding products shifts equilibrium left Equilibrium as a balance of free energy Removing products shifts equilibrium right Heat as a reactant or product in equilibrium Pressure effect on equilibrium with gases

Change	Effect
Concentration	Increase in reactant or decrease in product favors forward reaction; increase in product or decrease in reactant favors reverse reaction.
Temperature	Increase favors endothermic; decrease favors exothermic.
Pressure	Increase favors side with fewer moles of gas; decrease favors side with more moles of gas.
Catalyst	Equilibrium reached faster; K unchanged.

Concept Map: Chemical Reactions

Concept map summarizing energy, spontaneity, rate, and extent of chemical reactions