Chapter 5: Thermochemistry

5.1 The Nature of Energy

Thermochemistry is the study of the energy changes that accompany chemical reactions, focusing primarily on heat transfer. Energy is the capacity to do work or produce heat.

• Energy: Exists in two main forms: kinetic energy (energy of motion) and potential energy (energy due to position or composition).

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed from one form to another.

• System and Surroundings: The system is the part of the universe being studied, while the surroundings are everything else.

• Work (w): Energy used to move an object against a force.

• Heat (q): Energy transferred between objects due to a temperature difference.

Example: When gasoline burns in a car engine, chemical potential energy is converted to kinetic energy and heat.

5.2 The First Law of Thermodynamics

The first law states that the total energy of the universe is constant. For a system, the change in internal energy (ΔE) is the sum of heat and work exchanged with the surroundings.

• Internal Energy (E): The sum of all kinetic and potential energies of the system.

• First Law Equation:

• Sign Conventions: q is positive if heat is absorbed by the system; w is positive if work is done on the system.

Example: Compressing a gas (work done on the system) increases its internal energy.

5.3 Enthalpy

Enthalpy (H) is a thermodynamic quantity equivalent to the total heat content of a system at constant pressure.

• Definition:

• Change in Enthalpy (ΔH): At constant pressure, the heat absorbed or released is equal to the change in enthalpy.

• Endothermic Process: ΔH > 0 (system absorbs heat)

• Exothermic Process: ΔH < 0 (system releases heat)

Example: Melting ice is endothermic; burning methane is exothermic.

5.4 Enthalpies of Reaction

The enthalpy change for a chemical reaction is called the enthalpy of reaction (ΔHrxn).

• Standard Enthalpy Change: Measured under standard conditions (1 atm, 25°C, 1 M concentration).

• Thermochemical Equations: Chemical equations that include enthalpy changes.

Example:

5.5 Calorimetry

Calorimetry is the measurement of heat flow in a chemical reaction or physical process.

• Calorimeter: Device used to measure heat changes.

• Specific Heat (c): Amount of heat required to raise the temperature of 1 g of a substance by 1°C.

• Heat Equation:

• Constant-Pressure Calorimetry: Used for reactions in solution (coffee-cup calorimeter).

• Bomb Calorimetry: Used for combustion reactions at constant volume.

Example: Measuring the heat released when dissolving NaOH in water.

5.6 Hess's Law

Hess's Law states that the total enthalpy change for a reaction is the same, no matter how many steps the reaction is carried out in.

• Application: Allows calculation of ΔH for reactions by combining known enthalpy changes of other reactions.

Example: If A → B (ΔH1) and B → C (ΔH2), then A → C (ΔH = ΔH1 + ΔH2).

5.7 Enthalpies of Formation

The standard enthalpy of formation (ΔHf°) is the enthalpy change when one mole of a compound is formed from its elements in their standard states.

• Standard State: Most stable form of an element at 1 atm and 25°C.

• Calculation of Reaction Enthalpy:

• Table of Standard Enthalpies of Formation: Used to calculate ΔH for various reactions.

Example: Calculating ΔH for the formation of water from H2 and O2.

Chapter 14: Chemical Kinetics

14.1 Factors that Affect Reaction Rates

Chemical kinetics is the study of the speed (rate) of chemical reactions and the factors that affect them.

• Concentration: Higher concentration usually increases reaction rate.

• Temperature: Higher temperature increases kinetic energy, leading to more effective collisions.

• Catalysts: Substances that increase reaction rate without being consumed.

• Surface Area: Greater surface area increases rate for heterogeneous reactions.

• Nature of Reactants: Some substances react faster than others due to bond strengths and molecular structure.

Example: Powdered zinc reacts faster with acid than a zinc strip due to increased surface area.

14.2 Reaction Rates

The reaction rate is the change in concentration of a reactant or product per unit time.

• Average Rate:

• Instantaneous Rate: The rate at a specific moment, found by the slope of the concentration vs. time curve.

• Rate Law: Expresses the rate as a function of reactant concentrations.

Example: For , rate =

14.3 Concentration and Rate

The rate law relates the rate of a reaction to the concentration of reactants, often in the form:

• k: Rate constant (depends on temperature and reaction).

• m, n: Reaction orders (determined experimentally).

• Overall Order: Sum of exponents in the rate law.

Example: If rate = k[A][B]^2, the reaction is first order in A, second order in B, and third order overall.

14.4 The Change of Concentration with Time

Integrated rate laws describe how concentrations change over time for different reaction orders.

• First-Order:

• Second-Order:

• Zero-Order:

• Half-Life (t1/2): Time for concentration to decrease by half; for first-order,

Example: Radioactive decay follows first-order kinetics.

14.5 Temperature & Rate

Reaction rates generally increase with temperature due to higher kinetic energy and more frequent effective collisions.

• Arrhenius Equation:

• Ea: Activation energy (minimum energy required for reaction).

• A: Frequency factor (related to collision frequency and orientation).

• R: Gas constant; T: Temperature in Kelvin.

Example: Increasing temperature from 25°C to 35°C can double the rate of some reactions.

14.6 Reaction Mechanisms

A reaction mechanism is a sequence of elementary steps that describes the pathway from reactants to products.

• Elementary Step: A single step in a mechanism, with its own rate law.

• Intermediate: Species produced and consumed during the reaction, not present in the overall equation.

• Rate-Determining Step: The slowest step, which controls the overall rate.

Example: The decomposition of hydrogen peroxide proceeds via a two-step mechanism with an intermediate.

14.7 Catalysis

Catalysts increase reaction rates by providing an alternative pathway with lower activation energy.

• Homogeneous Catalysis: Catalyst is in the same phase as reactants.

• Heterogeneous Catalysis: Catalyst is in a different phase (e.g., solid catalyst with gaseous reactants).

• Enzymes: Biological catalysts that are highly specific and efficient.

Example: Platinum in catalytic converters speeds up the breakdown of exhaust gases.