Chemical Thermodynamics: Concepts

• Chemistry is the study of matter, its properties, and how matter changes.
• Physical chemistry studies macroscopic and particulate phenomena in chemical systems using principles and concepts from physics.
  • Includes motion, energy, force, time, thermodynamics, quantum chemistry, statistical mechanics, analytical dynamics, and chemical equilibria.

Chemical Thermodynamics

• The study of the interrelation of heat and work with chemical reactions or physical changes of state.
• Involves applying thermodynamic principles to chemical systems and reactions.
• Predicts the feasibility and equilibrium composition of chemical reactions.
• Key concepts: energy, entropy, Gibbs free energy, and chemical potential.
• Focuses on the bulk properties of matter, not microscopic details.

Basic Concepts in Thermodynamics

• System: The part of the universe under investigation.
• Surroundings: Everything outside the system.
• Boundary: Separates the system from the surroundings.
• Types of systems:
  • Open system: Exchanges both matter and energy with the surroundings.
  • Closed system: Exchanges energy but not matter with the surroundings.
  • Isolated system: Exchanges neither matter nor energy with the surroundings.
• State variables: Properties defining the state of a system like pressure, volume, temperature, and composition.
• State functions: Properties depending only on the initial and final states of the system, such as internal energy, enthalpy, entropy, and Gibbs free energy.
• Intensive properties: Independent of the amount of substance, examples include temperature, pressure, and density.
• Extensive properties: Dependent on the amount of substance, examples include volume, mass, and energy.

Laws of Thermodynamics

• Zeroth Law: If two systems are each in thermal equilibrium with a third system, then they are in thermal equilibrium with each other.
• First Law: The total energy of an isolated system is constant; energy can be transformed, but not created or destroyed.
  • ΔU = Q - W, where ΔU is the change in internal energy, Q is the heat added to the system, and W is the work done by the system.
• Second Law: The entropy of an isolated system tends to increase.
  • Entropy (S) measures the disorder or randomness of a system.
  • ΔS ≥ 0 for an isolated system.
• Third Law: The entropy of a perfect crystal at absolute zero (0 K) is zero.
  • Provides a reference point for determining the absolute entropy of a substance at any other temperature.

Enthalpy (H)

• A thermodynamic property of a system, defined as the sum of the system's internal energy (U) and the product of its pressure (P) and volume (V).
• H = U + PV
• A state function, so the change in enthalpy (ΔH) depends only on the initial and final states.
• ΔH = ΔU + PΔV (at constant pressure)
• In chemical reactions, enthalpy change (ΔH) represents the heat absorbed or released during the reaction at constant pressure.
• Exothermic reaction: ΔH < 0 (heat is released).
• Endothermic reaction: ΔH > 0 (heat is absorbed).
• Standard enthalpy of formation (ΔHf°): The enthalpy change when one mole of a compound is formed from its elements in their standard states (usually 298 K and 1 atm).

Entropy (S)

• Measures the disorder or randomness of a system.
• A state function.
• The change in entropy (ΔS) is related to the heat absorbed or released (Q) and the temperature (T) by the equation: ΔS = Qrev/T, where Qrev is the heat exchanged in a reversible process.
• Entropy increases with increasing temperature, volume, and number of particles.
• A key factor in determining the spontaneity of a process.
• Standard entropy (S°): The absolute entropy of a substance at 298 K and 1 atm.

Gibbs Free Energy (G)

• A thermodynamic potential that measures the amount of energy available in a thermodynamic system to do useful work at constant temperature and pressure.
• G = H - TS, where H is enthalpy, T is temperature, and S is entropy.
• A state function.
• The change in Gibbs free energy (ΔG) is a criterion for spontaneity at constant temperature and pressure.
• ΔG < 0: The process is spontaneous ( Gibbs free energy decreases).
• ΔG > 0: The process is non-spontaneous ( Gibbs free energy increases).
• ΔG = 0: The process is at equilibrium ( Gibbs free energy is at a minimum).
• ΔG = ΔH - TΔS
• Standard Gibbs free energy change (ΔG°): The change in Gibbs free energy when a reaction is carried out under standard conditions (usually 298 K and 1 atm).

Chemical Potential (μ)

• The change in the Gibbs free energy of a system when one mole of a substance is added to the system at constant temperature and pressure.
• Indicates how the Gibbs free energy of a system changes with variations in the number of moles of a particular component.
• An intensive property.
• For a pure substance, the chemical potential is equal to the molar Gibbs free energy.
• Crucial in determining phase equilibria and chemical equilibria.

Phase Equilibria

• Deals with the conditions under which two or more phases of a substance can coexist in equilibrium.
• Phase: A homogeneous part of a system that is uniform in chemical composition and physical properties.
• Phase transitions: Processes such as melting, boiling, sublimation, and phase transformations between solid forms.
• Gibbs Phase Rule: F = C - P + 2
  • F is the number of degrees of freedom.
    • The number of independent variables that can be changed without altering the number of phases in equilibrium.
  • C is the number of components.
    • The minimum number of independent chemical species necessary to define the composition of all phases in the system.
  • P is the number of phases in equilibrium.

Chemical Equilibria

• The state in which the rates of the forward and reverse reactions are equal.
• The net change in concentrations of reactants and products is zero.
• Equilibrium constant (K): A quantitative measure of the extent to which a reaction will proceed to completion.
• For a reversible reaction: aA + bB ⇌ cC + dD
• The equilibrium constant is expressed as: K = ([C]^c [D]^d) / ([A]^a [B]^b), where [A], [B], [C], and [D] are the equilibrium concentrations of the reactants and products.
• ΔG° = -RTlnK, where R is the gas constant and T is the temperature in Kelvin.
• Le Chatelier's Principle: If a change of condition is applied to a system in equilibrium, the system will shift in a direction that relieves the stress.
  • Changes in conditions include changes in concentration, pressure, temperature, or the addition of an inert gas.

Applications of Chemical Thermodynamics

• Predicting the spontaneity of chemical reactions.
• Determining the equilibrium composition of reaction mixtures.
• Calculating the heat and work associated with chemical or physical transformations.
• Understanding phase equilibria and phase transitions.
• Optimizing industrial chemical processes to maximize yield and minimize energy consumption.
• Studying the thermodynamics of solutions and mixtures.
• Analyzing electrochemical cells and batteries.