Fundamental Concepts of Thermodynamics and Energy Changes

Posted on Apr 3, 2025 in Chemistry

What is Thermodynamics?

Thermodynamics is the part of physics that studies the heat and work exchanges accompanying physicochemical processes. When these processes are chemical reactions, the specific branch studying them is called thermochemistry. The energy exchanged in a chemical reaction can manifest as heat, light, electrical energy, etc.

Thermodynamic Systems

System, Environment, and Universe

A thermodynamic system is the specific part of the universe isolated for study of its physicochemical properties. Everything surrounding the system is called the environment, which constitutes the part of the universe outside the system. The combination of the system and its environment makes up the universe.

Types of Systems

The relationship between a system and its environment defines its type:

  • Open System: Exchanges both matter and energy with the environment.
  • Closed System: Exchanges energy with the environment, but not matter.
  • Isolated System: Exchanges neither matter nor energy with the environment.

Thermodynamic Processes

Reversible vs. Irreversible Processes

  • Reversible Process: Occurs through a series of equilibrium states, where a tiny change in conditions can reverse the process direction.
  • Irreversible Process: The system evolves spontaneously in only one direction.

Common Process Types

  • Isothermal Process: Temperature remains constant (ΔT = 0).
  • Isobaric Process: Pressure remains constant (ΔP = 0).
  • Isochoric Process (or Isovolumetric): Volume remains constant (ΔV = 0).

Internal Energy (U)

The internal energy (U) is the sum of all kinetic and potential energies of the particles that constitute the system. It is measured in Joules (J) and is a state function, meaning its value depends only on the current state of the system, not how it got there.

First Law of Thermodynamics

The first law of thermodynamics (law of conservation of energy) states that when a system evolves from an initial state to a final state, the change experienced by its internal energy (ΔU) equals the sum of the heat (Q) exchanged with the environment and the work (W) done on or by the system during this transformation:

ΔU = Q + W

Note: Sign conventions for Q and W may vary. Common convention: Q > 0 for heat added to the system, W > 0 for work done on the system.

Enthalpy (H)

Enthalpy (H) is a thermodynamic property representing the total heat content of a system. The change in enthalpy (ΔH) represents the heat exchanged with the environment during a process occurring at constant pressure. Enthalpy is also a state function and is measured in Joules (J).

Thermochemical Equations

A thermochemical equation is a balanced chemical equation that includes:

  • The substances involved (reactants and products).
  • The appropriate stoichiometric coefficients.
  • The physical state (solid (s), liquid (l), gas (g), aqueous (aq)) of each substance.
  • The enthalpy change (ΔH) associated with the reaction, indicating the amount of heat absorbed or released.

It often specifies the pressure and temperature at which the reaction occurs. The sign of ΔH indicates the nature of the process:

  • ΔH < 0: Exothermic (heat is released)
  • ΔH > 0: Endothermic (heat is absorbed)

Hess’s Law

Hess’s Law states that if a chemical reaction can be expressed as the sum of a series of steps (either hypothetical or actual), the total enthalpy change for the overall reaction is equal to the sum of the enthalpy changes for the individual steps. This holds true regardless of the path taken.

Standard Enthalpy Changes

Standard Enthalpy of Formation (ΔH°f)

The standard enthalpy of formation is the enthalpy change when one mole of a compound is formed from its constituent elements in their most stable states under standard conditions (usually 1 atm pressure and 25°C or 298 K). By convention, the ΔH°f of any element in its most stable form under standard conditions is zero.

Standard Enthalpy of Combustion (ΔH°c)

The standard enthalpy of combustion is the enthalpy change when one mole of a substance reacts completely with excess oxygen (O₂) under standard conditions to produce specified combustion products (commonly CO₂(g) for carbon compounds and H₂O(l) for hydrogen compounds).

Standard Enthalpy of Hydrogenation

The standard enthalpy of hydrogenation is the enthalpy change when one mole of an unsaturated compound reacts with hydrogen (H₂) to become saturated, under standard conditions.

Bond Enthalpy (Bond Energy)

Bond enthalpy is the average energy required to break one mole of a specific type of bond between two atoms in the gaseous state.

Entropy (S)

Entropy (S) is a thermodynamic state function that measures the degree of disorder, randomness, or energy dispersal in a system. The greater the disorder, the higher the entropy. It is typically measured in Joules per Kelvin (J/K).

Second Law of Thermodynamics

The second law of thermodynamics states that for any spontaneous process, the total entropy of the universe (system plus environment) must increase (ΔS_universe > 0). Key points include:

  1. When a system releases heat to the environment, the entropy of the environment increases.
  2. If a system absorbs heat from the environment, the system’s entropy increases while the environment’s entropy decreases. The process will be spontaneous only if the increase in the system’s entropy is greater than the decrease in the environment’s entropy, resulting in a net positive ΔS_universe.

Third Law of Thermodynamics

The third law of thermodynamics states that the entropy of a perfect crystalline substance approaches zero as the temperature approaches absolute zero (0 Kelvin).

Properties of Entropy Change (ΔS)

  1. Entropy is an extensive property (it depends on the amount of substance).
  2. The entropy change (ΔS) for a process is equal in magnitude but opposite in sign to the ΔS for the reverse process.
  3. When a reaction results in an increase in the number of moles of gas or breaks down larger molecules into smaller ones, the entropy change (ΔS) is generally positive (ΔS > 0).

Gibbs Free Energy (G) and Spontaneity

Gibbs Free Energy (G) is a thermodynamic potential that combines enthalpy (H) and entropy (S) to determine the spontaneity of a process at constant temperature and pressure (G = H – TS). The change in Gibbs Free Energy (ΔG) is the ultimate criterion for spontaneity under these conditions.

Standard Free Energy of Formation (ΔG°f)

The standard free energy of formation (ΔG°f) is the change in free energy associated with the formation of one mole of a compound from its constituent elements in their standard states (most stable form, standard conditions). For elements in their standard states, ΔG°f = 0.

Spontaneity Criteria (ΔG)

At constant temperature and pressure:

  • If ΔG < 0: The process is spontaneous in the forward direction (as written).
  • If ΔG > 0: The process is non-spontaneous in the forward direction but spontaneous in the reverse direction.
  • If ΔG = 0: The system is at equilibrium.