Energy Changes, Reaction Rates, and Equilibrium – Chemistry Chapter 6 Study Notes

Chapter 6: Energy Changes, Reaction Rates, and Equilibrium

6.1 Energy

Energy is a fundamental concept in chemistry and biology, representing the capacity to do work. Understanding energy is essential for analyzing chemical reactions and physiological processes.

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

• Potential energy: Stored energy, often found in chemical bonds.

• Kinetic energy: The energy of motion.

• Law of Conservation of Energy: States that the total energy in a system remains constant; energy cannot be created or destroyed.

Chemical bonds store potential energy. Compounds with lower potential energy are more stable than those with higher potential energy. Reactions that produce products with lower potential energy than the reactants are generally favored.

A. The Units of Energy

• Calorie (cal): The amount of energy needed to raise the temperature of 1 g of water by 1°C.

• Joule (J): Another unit of energy. $1\,\text{cal} = 4.184\,\text{J}$

• Kilojoules (kJ) and kilocalories (kcal) are larger units. $1,000\,\text{J} = 1\,\text{kJ}$ $1,000\,\text{cal} = 1\,\text{kcal}$ $1\,\text{kcal} = 4.184\,\text{kJ}$

6.2 Energy Changes in Reactions

Chemical reactions involve the breaking and forming of bonds, which are associated with energy changes.

• Bond breaking requires an input of energy (endothermic).

• Bond formation releases energy (exothermic).

Example: To break a Cl–Cl bond, 58 kcal/mol must be added; to form it, 58 kcal/mol is released.

A. Bond Dissociation Energy

• ΔH (Enthalpy change): The energy absorbed or released in a reaction; also called the heat of reaction.

• Endothermic reaction: Energy is absorbed, ΔH is positive ($\Delta H > 0$).

• Exothermic reaction: Energy is released, ΔH is negative ($\Delta H < 0$).

• Bond dissociation energy: The ΔH for breaking a covalent bond by equally dividing the electrons between the two atoms. Always positive because bond breaking is endothermic.

Example equations: Bond breaking: $\text{H–H} \rightarrow \text{H}^\cdot + \text{H}^\cdot$   $\Delta H = +104\,\text{kcal/mol}$ Bond formation: $\text{H}^\cdot + \text{H}^\cdot \rightarrow \text{H–H}$   $\Delta H = -104\,\text{kcal/mol}$

• The stronger the bond, the higher its bond dissociation energy.

• Bond dissociation energies generally decrease down a group in the periodic table.

Bond	ΔH (kcal/mol)
H–F	+136
H–Cl	+103
H–Br	+88
H–I	+71

B. Calculations Involving ΔH Values

• ΔH indicates the relative strength of bonds broken and formed.

• If ΔH is negative: More energy is released in forming bonds than is needed to break bonds; products have stronger bonds than reactants.

• If ΔH is positive: More energy is needed to break bonds than is released in forming new bonds; reactants have stronger bonds than products.

Example (Exothermic): $\text{CH}_4 (g) + 2\text{O}_2 (g) \rightarrow \text{CO}_2 (g) + 2\text{H}_2\text{O} (l)$ $\Delta H = -213\,\text{kcal/mol}$ (Heat is released)

Example (Endothermic): $6\text{CO}_2 (g) + 6\text{H}_2\text{O} (l) \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 (aq) + 6\text{O}_2 (g)$ $\Delta H = +678\,\text{kcal/mol}$ (Heat is absorbed)

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

6.3 Energy Diagrams

Energy diagrams illustrate the energy changes during a chemical reaction, including the activation energy and enthalpy change.

• For a reaction to occur, molecules must collide with sufficient kinetic energy and proper orientation.

• Activation energy (Ea): The minimum energy required for reactants to reach the transition state and react.

• The height of the energy barrier (Ea) determines the reaction rate.

• If Ea is high, the reaction is slow; if Ea is low, the reaction is fast.

• The difference in energy between reactants and products is ΔH.

• If ΔH is negative, the reaction is exothermic; if positive, endothermic.

6.4 Reaction Rates

The rate of a chemical reaction depends on several factors, including concentration, temperature, and the presence of catalysts.

A. How Concentration and Temperature Affect Reaction Rate

• Increasing reactant concentration increases the number of collisions and the reaction rate.

• Increasing temperature increases the average kinetic energy of molecules and the reaction rate.

B. Catalysts

• Catalyst: A substance that speeds up the rate of a reaction without being consumed.

• Catalysts lower the activation energy (Ea) but do not affect ΔH.

• Uncatalyzed reactions (higher Ea) are slower; catalyzed reactions (lower Ea) are faster.

C. Focus on the Human Body: Biological Catalysts

• Enzymes: Biological catalysts, usually proteins, with specific three-dimensional shapes.

• The active site of an enzyme binds a reactant (substrate), facilitating a specific reaction at an enhanced rate.

• Lactase: An enzyme that converts lactose into glucose and galactose. Deficiency leads to lactose intolerance.

6.5 Equilibrium of Reaction

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

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

• At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction.

• The concentrations of reactants and products remain constant at equilibrium.

A. The Equilibrium Constant (K)

• K expresses the ratio of product concentrations to reactant concentrations at equilibrium.

• For a general reaction: $aA + bB \rightarrow cC + dD$ $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

• Brackets [ ] indicate concentration in moles per liter (mol/L).

• Coefficients in the balanced equation become exponents in the K expression.

Example: $\text{N}_2 (g) + \text{O}_2 (g) \rightarrow 2\text{NO}$ $K = \frac{[NO]^2}{[N_2][O_2]}$

B. The Magnitude of the Equilibrium Constant

• If K > 1: Products are favored; equilibrium lies to the right.

• If K < 1: Reactants are favored; equilibrium lies to the left.

• If K ≈ 1 (0.01 < K < 100): Both reactants and products are present in similar amounts.

Value of K	Position of Equilibrium
K > 1	Equilibrium favors the products (right).
K < 1	Equilibrium favors the reactants (left).
K = 1	Both reactants and products are present at equilibrium.

Example calculation: Given: [A2] = 0.25 M, [B2] = 0.25 M, [AB] = 0.50 M Reaction: $A_2 (g) + B_2 (g) \rightarrow 2AB$ $K = \frac{[AB]^2}{[A_2][B_2]} = \frac{(0.50)^2}{0.25 \times 0.25} = \frac{0.25}{0.0625} = 4.0$

6.6 Le Châtelier’s Principle

Le Châtelier’s Principle describes how a system at equilibrium responds to disturbances.

• If a system at equilibrium is disturbed (by changes in concentration, temperature, or pressure), it will shift to counteract the disturbance.

A. Concentration Changes

• Increasing the concentration of a reactant drives the reaction to the right (toward products).

• Increasing the concentration of a product drives the reaction to the left (toward reactants).

• Removing a product drives the reaction to the right.

B. Temperature Changes

• Increasing temperature favors the endothermic direction (absorbs heat).

• Increasing temperature favors the reverse reaction if the forward reaction is exothermic (releases heat).

• Decreasing temperature favors the exothermic direction.

C. 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.

Summary Table: Effects of Changes on Equilibrium

Change	Effect on Equilibrium
Increase reactant concentration	Shifts right (toward products)
Increase product concentration	Shifts left (toward reactants)
Remove product	Shifts right
Increase temperature (endothermic)	Shifts right
Increase temperature (exothermic)	Shifts left
Increase pressure	Shifts toward fewer moles of gas
Decrease pressure	Shifts toward more moles of gas

6.7 Focus on the Human Body

Body temperature regulation is a key physiological process that depends on chemical reactions and energy changes. Enzymes play a crucial role in maintaining metabolic rates and homeostasis.

Additional info: These chemistry principles are foundational for understanding metabolic reactions, enzyme function, and physiological equilibrium in Anatomy & Physiology.