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.

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.

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:

• 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.

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.

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.

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:

• 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.

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.

Concept Map: Chemical Reactions