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

7.1 Energy and Chemical Bonds

Chemical bonds involve energy changes, primarily in the form of potential and kinetic energy. Understanding these concepts 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. Examples include water falling, a car rolling downhill, or a spring uncoiling.

• When bonds form (ionic or covalent), potential energy is often converted to heat (kinetic energy of particles).

• In reactions, bond breaking requires energy (energy in), while bond formation releases energy (energy out).

• 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

Energy changes during chemical reactions are governed by the breaking and forming of bonds, described quantitatively by bond dissociation energies and enthalpy changes.

• Bond Dissociation Energy: The energy required to break a bond in a gaseous molecule. Higher values indicate stronger, more stable bonds.

• Endothermic Reaction: Absorbs heat (bond breaking dominates). ΔH is positive.

• Exothermic Reaction: Releases heat (bond formation dominates). ΔH is negative.

• The law of conservation of energy states that energy cannot be created or destroyed, only transformed.

• Heat of Reaction (ΔH): The difference between energy absorbed in breaking bonds and energy released in forming bonds.

• ΔH is measured at constant pressure and is called enthalpy change.

Table: Average Bond Dissociation Energies

7.3 Exothermic and Endothermic Reactions

Reactions are classified based on whether they release or absorb heat. The sign and magnitude of ΔH indicate the direction and amount of heat flow.

• Exothermic: Products have stronger bonds than reactants; energy is released (ΔH < 0).

• Endothermic: More energy is required to break bonds than is released in forming new ones; energy is absorbed (ΔH > 0).

• The reverse of an exothermic reaction is endothermic, and vice versa. The magnitude of heat transfer is the same, but the sign of ΔH is opposite.

Formula for Heat of Reaction:

Worked Example: Calculating ΔH for H2 + O2 → H2O

• Given: 2(H—H) + (O=O) → 4(O—H)

• Calculation:

  • Bond energies: H—H = 103 kcal/mol, O=O = 119 kcal/mol, O—H = 112 kcal/mol

  • kcal/mol

7.4 Why Do Chemical Reactions Occur? Free Energy

Whether a reaction occurs spontaneously depends on both energy changes (enthalpy) and changes in disorder (entropy).

• Spontaneous Process: Proceeds on its own without external influence (e.g., water flowing downhill).

• Nonspontaneous Process: Requires continuous external input.

• Entropy (S): A measure of molecular disorder, in units of cal/(mol·K) or J/(mol·K).

• Entropy increases when a gas forms from a solid or liquid, or when the number of particles increases.

• ΔS > 0: Disorder increases; ΔS < 0: Disorder decreases.

Free Energy Change (ΔG): Determines spontaneity by combining enthalpy and entropy effects.

• Exergonic: ΔG < 0, spontaneous, free energy released.

• Endergonic: ΔG > 0, nonspontaneous, free energy absorbed.

• Temperature can affect spontaneity; some reactions become spontaneous only at higher temperatures.

Table: Factors Affecting Spontaneity

7.5 How Do Chemical Reactions Occur? Reaction Rates

The rate of a chemical reaction depends on the frequency and energy of collisions between reactant molecules.

• Collision Theory: Reactants must collide with proper orientation and sufficient energy to break bonds and form new ones.

• Activation Energy (Ea): The minimum energy required for a reaction to occur.

• Reaction Rate: How quickly reactants are converted to products.

• The lower the activation energy, the faster the reaction; the higher the activation energy, the slower the reaction.

• ΔG (free energy change) does not determine reaction rate.

Energy Diagram:

• Shows reactant and product energy levels, activation energy barrier, and ΔG.

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

Several factors can increase the rate of a chemical reaction by affecting the frequency and energy of collisions.

• Temperature: Raising temperature increases particle energy and collision frequency, often doubling the rate for every 10°C rise.

• Concentration: Higher concentration increases collision frequency, often increasing the reaction rate proportionally.

• Catalysts: Substances that lower the activation energy by providing an alternative pathway or orienting reactants. Catalysts do not affect ΔH or ΔG, only the rate.

Table: Effects of Changes in Reaction Conditions on Reaction Rates

7.7 Reversible Reactions and Chemical Equilibrium

Many reactions are reversible, meaning they can proceed in both forward and reverse directions. Chemical equilibrium is reached when the rates of the forward and reverse reactions are equal.

• Reversible Reaction: Can proceed in both directions.

• Chemical Equilibrium: State where concentrations of reactants and products remain constant because forward and reverse rates are equal.

• Equilibrium is dynamic; reactions continue to occur, but no net change in concentrations.

7.8 Equilibrium Equations and Equilibrium Constants

The position of equilibrium in a reversible reaction is described by the equilibrium constant (K), which quantifies the ratio of product to reactant concentrations at equilibrium.

• Equilibrium Constant (K): For a reaction aA + bB ⇌ cC + dD,

• Pure solids and liquids are omitted from the expression.

• K varies with temperature; units are often omitted.

• A large K (>1) means products are favored; a small K (<1) means reactants are favored.

Worked Example: Writing an Equilibrium Expression

• For CH4(g) + H2O(g) ⇌ CO(g) + 3H2(g):

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) such as changes in concentration, temperature, or pressure.

• Le Châtelier’s Principle: When a stress is applied to a system at equilibrium, the system shifts to relieve the stress.

• Concentration:

  • Increase reactant: shifts right (toward products).

  • Increase product: shifts left (toward reactants).

  • Decrease reactant: shifts left; decrease product: shifts right.

• Temperature:

  • Increase temperature: favors endothermic direction.

  • Decrease temperature: favors exothermic direction.

• Pressure (for gases):

  • Increase pressure (decrease volume): shifts toward side with fewer gas molecules.

  • Decrease pressure (increase volume): shifts toward side with more gas molecules.

• Catalysts: Do not affect equilibrium position; only speed up attainment of equilibrium.

Table: Effects of Changes in Reaction Conditions on Equilibria

Summary Table: Key Terms and Definitions

Additional info: For more detailed calculations, refer to worked examples involving stoichiometry, unit conversions (kcal to kJ), and equilibrium constant calculations as shown in the original material.