Study Guide: Chemical Reactions – Energy, Rate, and Equilibrium

Ch.7 Chemical Reactions: Energy, Rate and Equilibrium

Nature of Energy

Energy is a fundamental concept in chemistry, representing the capacity to do work or produce heat. Thermochemistry studies the energy changes associated with chemical reactions and physical changes.

• Energy: The ability to do work or produce heat.

• SI Unit for Energy: Joule (J).

• Classification of Energy:

  • Potential Energy: Energy due to position or arrangement of atoms.

  • Kinetic Energy: Energy due to motion of atoms.

  • Chemical Energy: Energy stored in chemical bonds.

  • Thermal Energy: Energy related to temperature and motion of atoms.

• Energy Conversion Factors:

  • 1 calorie (cal) = 4.184 J

  • 1 Calorie (Cal) = 4184 J

  • 1 kilowatt-hour (kWh) = 3.60 x 106 J

First Law of Thermodynamics

The First Law states that energy cannot be created or destroyed, only transferred between a system and its surroundings.

• System: The part of the universe being studied (e.g., a chemical reaction).

• Surroundings: Everything outside the system.

• Heat (q): Transfer of thermal energy from a higher temperature object to a lower temperature object.

• Work (w): Energy transfer due to movement against a force.

Endothermic & Exothermic Reactions

Chemical reactions can absorb or release energy, affecting the speed and nature of molecular changes.

• Endothermic Reactions: Absorb heat from surroundings; molecules speed up and may break bonds.

• Exothermic Reactions: Release heat to surroundings; molecules slow down and may form bonds.

Bond Energy and Enthalpy of Reaction

Bond energy is the energy required to break a bond between atoms. It is used to calculate the enthalpy change (ΔH) of a reaction.

• Bond Energy (ΔHB.E.): Energy stored in a bond.

• Enthalpy of Reaction (ΔHRxn): Calculated using bond energies of reactants and products.

• Endothermic Process: Energy is absorbed to break bonds (ΔH > 0).

• Exothermic Process: Energy is released when bonds form (ΔH < 0).

Enthalpy of Reaction Formula:

$\Delta H_{\text{Rxn}} = \text{Sum of bond energies of reactants} - \text{Sum of bond energies of products}$

Thermochemical Equations and Stoichiometry

Thermochemical equations include enthalpy changes and are used to relate quantities of reactants and products.

• Convert given quantity to moles.

• Use mole-to-mole comparison from balanced equation.

• Convert moles of unknown to desired units (grams, molecules, etc.).

• Compare theoretical yields if multiple products are calculated.

Spontaneous vs Nonspontaneous Reactions

Spontaneity determines whether a reaction occurs without continuous external energy.

• Spontaneous: Occurs without constant outside energy; products form at equilibrium.

• Nonspontaneous: Requires constant outside energy; does not form products at equilibrium.

• Spontaneity does not determine reaction speed.

Entropy (S)

Entropy is a measure of disorder or randomness in a system. The Second Law of Thermodynamics states that the entropy of the universe always increases in spontaneous reactions.

• Entropy Change (ΔS): Indicates increase or decrease in disorder.

• Physical changes (solid → liquid → gas) generally increase entropy.

• Chemical changes can increase or decrease entropy depending on bond formation and number of gas molecules.

Gibbs Free Energy (ΔG)

Gibbs Free Energy determines the spontaneity of a reaction. It combines enthalpy (ΔH), entropy (ΔS), and temperature (T).

• ΔG Formula: $\Delta G = \Delta H - T\Delta S$

• ΔG & Spontaneity:

  • ΔG < 0: Spontaneous

  • ΔG > 0: Nonspontaneous

  • ΔG = 0: Equilibrium

• Sign of ΔH and ΔS can predict spontaneity at different temperatures.

Rate of Reaction

Chemical kinetics studies the rate at which reactants are converted to products.

• Factors Affecting Rate:

  • Concentration: Higher concentration increases collision frequency.

  • Surface Area: Greater surface area increases collision frequency.

  • Temperature: Higher temperature increases collision energy and frequency.

  • Catalyst: Lowers activation energy, increasing reaction rate.

Energy Diagrams

Energy diagrams illustrate the energy changes during a reaction, including reactants, products, transition state, and activation energy.

• Reactants: Start of the diagram.

• Products: End of the diagram.

• Transition State: Highest energy point.

• Activation Energy (Ea): Energy barrier for the reaction.

• Overall Energy: Difference between reactants and products.

Rate Law

The rate law expresses the relationship between reaction rate, reactant concentrations, and reaction order.

• Rate Law Expression: $\text{Rate} = k [A]^m [B]^n$

• k: Rate constant

• m, n: Reaction orders (determined experimentally)

• Rate law ignores products.

Chemical Equilibrium

Most reactions reach equilibrium, where the forward and reverse reactions occur at equal rates and concentrations remain constant.

• Equilibrium: Reactant and product concentrations are constant, not necessarily equal.

• Dynamic Process: Both forward and reverse reactions continue.

The Equilibrium Constant (K)

The equilibrium constant (K) is the ratio of product concentrations to reactant concentrations at equilibrium.

• K > 1: Products favored.

• K < 1: Reactants favored.

• K excludes solids and pure liquids.

• Equilibrium Expression: $K = \frac{[\text{products}]}{[\text{reactants}]}$

Le Chatelier’s Principle

If a system at equilibrium is disturbed, it will shift to minimize the disturbance and re-establish equilibrium.

• Disturbances: Changes in concentration, pressure, or temperature.

• Catalysts: Do not shift equilibrium, only change rate.

Additional info: Academic context and explanations have been expanded for clarity and completeness. All included images directly reinforce the adjacent explanations and are strictly relevant to the topic.