Energy Changes, Reaction Rates, and Chemical Equilibrium6

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Energy Changes, Reaction Rates, and Equilibrium

Energy in Chemical Systems

Energy is a fundamental concept in chemistry, describing the capacity to do work. Chemical reactions involve changes in energy, which can be stored (potential energy) or associated with motion (kinetic energy). The law of conservation of energy states that energy cannot be created or destroyed, only transformed.

Potential energy: Stored energy, often in chemical bonds.
Kinetic energy: Energy of motion.
Chemical bonds: Store potential energy; reactions favor products with lower potential energy (greater stability).

Energy is measured in calories (cal) and joules (J):

1 cal = 4.184 J
1 kcal = 1000 cal = 4.184 kJ

Energy Changes in Chemical Reactions

During chemical reactions, bonds in reactants are broken and new bonds are formed in products. The energy required to break a bond is called the bond dissociation energy, and the energy change for the reaction is denoted as ΔH (enthalpy change).

Bond breaking: Always requires energy (endothermic, ΔH > 0).
Bond formation: Always releases energy (exothermic, ΔH < 0).

For example, breaking a Cl–Cl bond requires +58 kcal/mol, while forming it releases –58 kcal/mol.

Bond dissociation energies for hydrogen halides

Bond Dissociation Energy

Stronger bonds have higher bond dissociation energies.
Bond dissociation energies generally decrease down a group in the periodic table.

Endothermic vs. Exothermic Reactions

Endothermic Reaction	Exothermic Reaction
Heat is absorbed. ΔH is positive. Bonds broken in reactants are stronger than bonds formed in products. Products are higher in energy than reactants.	Heat is released. ΔH is negative. Bonds formed in products are stronger than bonds broken in reactants. Products are lower in energy than reactants.

Table comparing endothermic and exothermic reactions

Energy Diagrams and Activation Energy

For a reaction to occur, molecules must collide with sufficient kinetic energy and proper orientation. The activation energy (Ea) is the minimum energy required for a reaction to proceed.

High Ea: Slow reaction (few molecules have enough energy).
Low Ea: Fast reaction (many molecules have enough energy).

Proper and improper molecular orientation in collisions Bond breaking and forming in a reaction Energy diagram showing activation energy and enthalpy change

Exothermic and Endothermic Energy Diagrams

Exothermic: Products are lower in energy than reactants (ΔH < 0).
Endothermic: Products are higher in energy than reactants (ΔH > 0).

Exothermic reaction energy diagram Endothermic reaction energy diagram

Reaction Rates

The rate of a chemical reaction depends on several factors:

Concentration: Higher concentration increases collision frequency and reaction rate.
Temperature: Higher temperature increases kinetic energy and reaction rate.
Catalysts: Substances that speed up reactions by lowering Ea without being consumed.

Catalyst in a reaction Energy diagram comparing catalyzed and uncatalyzed reactions

Biological Catalysts: Enzymes

Enzymes are protein catalysts with specific three-dimensional shapes.
They bind reactants (substrates) at the active site and increase reaction rates.
Example: Lactase converts lactose into glucose and galactose.

Chemical Equilibrium

A reversible reaction can proceed in both forward and reverse directions. At equilibrium, the rates of the forward and reverse reactions are equal, and the concentrations of reactants and products remain constant.

Forward reaction proceeds to the right Reverse reaction proceeds to the left

The Equilibrium Constant (K)

The equilibrium constant, K, expresses the ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their coefficients:

Brackets [ ] indicate concentration in mol/L.
Coefficients become exponents in the expression.

Magnitude of K and Position of Equilibrium

Value of K	Position of Equilibrium
K > 1	Equilibrium favors products (lies to the right).
K < 1	Equilibrium favors reactants (lies to the left).
K ≈ 1	Both reactants and products are present in similar amounts.

Table relating K to equilibrium position

Calculating the Equilibrium Constant

For a reaction: A2 + B2 ⇌ 2 AB, with equilibrium concentrations [A2] = 0.25 M, [B2] = 0.25 M, [AB] = 0.50 M:

Calculation of equilibrium constant

Le Châtelier’s Principle

If a system at equilibrium is disturbed, it will shift to counteract the disturbance. Disturbances include changes in concentration, temperature, or pressure.

Concentration Changes

Adding reactant: Shifts equilibrium to the right (toward products).
Adding product: Shifts equilibrium to the left (toward reactants).
Removing product: Shifts equilibrium to the right.

Adding reactant drives reaction right Adding product drives reaction left Removing product drives reaction right

Temperature Changes

Increasing temperature favors the endothermic direction (absorbs heat).
Decreasing temperature favors the exothermic direction (releases heat).

Increasing temperature favors endothermic reaction

Pressure Changes

Increasing pressure: Shifts equilibrium toward the side with fewer moles of gas.
Decreasing pressure: Shifts equilibrium toward the side with more moles of gas.

Increasing pressure favors side with fewer moles Decreasing pressure favors side with more moles

Summary Table: Effects of Changes on Equilibrium

Change	Effect on Equilibrium
Adding reactant	Favors products
Removing reactant	Favors reactants
Adding product	Favors reactants
Removing product	Favors products
Increasing temperature	Favors products in endothermic, reactants in exothermic
Decreasing temperature	Favors reactants in endothermic, products in exothermic
Increasing pressure	Favors side with fewer moles
Decreasing pressure	Favors side with more moles

Summary table of equilibrium changes