Chemical Reactions: Energy, Rates, and Equilibrium – Study Notes

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 over a dam, a rolling car, or an uncoiling spring.

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

• In reactions, bond breaking requires energy input, while bond formation releases energy.

• If products have less potential energy than reactants, they are more stable and less likely to undergo further change.

• The energy change in a reaction depends on the difference in potential energy between reactants and products.

7.2 Heat Changes during Chemical Reactions

Heat changes in chemical reactions are governed by the breaking and forming of bonds, and are subject to the law of conservation of energy.

• 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). ΔH is positive.

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

• The direction of energy flow is indicated by the sign of ΔH: positive for absorbed, negative for released.

• The amount of heat absorbed in one direction equals the amount released in the reverse direction (but with opposite sign).

• Law of Conservation of Energy: 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. Also called enthalpy change.

Table: Average Bond Dissociation Energies

Bond	Average Bond Energy (kcal/mol)
H—H	103
O═O	119
O—H	112
Additional info: Other common bonds (C—H, C—C, etc.) can be found in standard tables.

7.3 Exothermic and Endothermic Reactions

Reactions are classified based on whether they release or absorb heat.

• Exothermic: Bonds formed in products are stronger than those broken in 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 absorbed or released is the same in both directions, but the sign of ΔH changes.

Example Calculation:

• For the reaction: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$

• ΔH calculation: $\Delta H = [2(103) + 119] - [4(112)] = -123$ kcal/mol

7.4 Why Do Chemical Reactions Occur? Free Energy

Spontaneity of reactions depends on energy changes and molecular disorder (entropy).

• Spontaneous Process: Proceeds on its own once started (e.g., water flowing downhill).

• Nonspontaneous Process: Requires continuous external influence.

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

• Gases have higher entropy than liquids, which have higher entropy than solids.

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

• Spontaneity is favored by heat release (ΔH < 0) and increased disorder (ΔS > 0).

• Free Energy Change (ΔG): Combines enthalpy and entropy to predict spontaneity.

Key Equation:

• $\Delta G = \Delta H - T\Delta S$

• ΔG < 0: Spontaneous (exergonic); ΔG > 0: Nonspontaneous (endergonic)

Example: Boiling water (liquid to vapor) increases entropy; smoke dispersing in a room increases entropy.

7.5 How Do Chemical Reactions Occur? Reaction Rates

The rate of a reaction depends on molecular collisions, orientation, and energy (activation energy).

• Collision Theory: Reactant particles must collide with proper orientation and sufficient energy to react.

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

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

• Lower activation energy means a faster reaction; higher activation energy means a slower reaction.

• The size of activation energy is independent of the free-energy change (ΔG).

Energy Diagram:

• Exergonic: Product energy < Reactant energy; ΔG negative.

• Endergonic: Product energy > Reactant energy; ΔG positive.

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

Reaction rates can be influenced by temperature, concentration, and catalysts.

• Temperature: Raising temperature increases particle energy, collision frequency, and reaction rate. A 10°C increase typically doubles the rate.

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

• Catalyst: Lowers activation energy by providing an alternative pathway or orienting molecules. Does not change ΔH or ΔG, only the rate.

Table: Effects of Changes in Reaction Conditions on Reaction Rates

Condition	Effect on Rate
Increase Temperature	Increases rate
Increase Concentration	Increases rate
Add Catalyst	Increases rate (no effect on equilibrium position)
Additional info: Decrease in temperature or concentration decreases rate.

7.7 Reversible Reactions and Chemical Equilibrium

Many reactions are reversible and can reach a state of equilibrium.

• Reversible Reaction: Can proceed in both forward and reverse directions.

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

• Equilibrium is dynamic: reactions continue, but no net change in concentrations occurs.

7.8 Equilibrium Equations and Equilibrium Constants

The equilibrium constant (K) quantifies the position of equilibrium for a reversible reaction.

• For a reaction: $aA + bB \rightleftharpoons cC + dD$

• Equilibrium constant expression: $K = \frac{[C]^c [D]^d}{[A]^a [B]^b}$

• Pure solids and liquids are omitted from the expression.

• K varies with temperature; units are often omitted.

• Large K: products favored; small K: reactants favored.

Example: For $\text{CH}_4 + \text{H}_2\text{O} \rightleftharpoons \text{CO} + 3\text{H}_2$

• $K = \frac{[\text{CO}][\text{H}_2]^3}{[\text{CH}_4][\text{H}_2\text{O}]}$

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).

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

• Stresses include: Changes in concentration, temperature, or pressure. Catalysts do not affect equilibrium position, only the rate at which equilibrium is reached.

• Concentration:

  • Increase reactant: shifts right (toward products).

  • Increase product: shifts left (toward reactants).

  • Decrease product: shifts right.

  • Continuous addition/removal prevents equilibrium.

• Temperature:

  • Increase favors endothermic direction.

  • Decrease favors exothermic direction.

  • Heat can be treated as a reactant or product.

• Pressure (for gases):

  • Increase (by decreasing volume): shifts toward side with fewer gas molecules.

  • Decrease: shifts toward side with more gas molecules.

Table: Effects of Changes in Reaction Conditions on Equilibria

Change	Effect on Equilibrium
Increase reactant concentration	Shifts right (toward products)
Increase product concentration	Shifts left (toward reactants)
Increase temperature (endothermic)	Shifts right
Increase temperature (exothermic)	Shifts left
Increase pressure (gases)	Shifts toward fewer gas molecules
Add catalyst	No effect on equilibrium position
Additional info: Removing product shifts equilibrium toward products.

Summary Table: Key Terms and Concepts

Term	Definition
Potential Energy	Stored energy due to position or arrangement
Kinetic Energy	Energy of motion
Bond Dissociation Energy	Energy required to break a bond
ΔH (Enthalpy Change)	Heat of reaction at constant pressure
ΔS (Entropy Change)	Change in molecular disorder
ΔG (Free Energy Change)	Predicts spontaneity of a process
Activation Energy	Minimum energy required for reaction
Catalyst	Substance that increases reaction rate without being consumed
Chemical Equilibrium	State where forward and reverse reaction rates are equal
Equilibrium Constant (K)	Quantifies position of equilibrium
Le Châtelier’s Principle	System shifts to relieve applied stress

Additional info: For more detailed calculations and worked examples, refer to the main textbook or instructor-provided materials.