Thermodynamics is a Greek word derived from “Thermo” (heat) and “Dynamics” (movement or transformation).
It is an essential branch of science that studies the interconversion of heat, work, and other forms of energy.
A system is a fixed quantity of matter with a defined boundary.
Surroundings are everything outside the system’s boundary that interacts with it.
A boundary is a real or imaginary surface that separates the system from its surroundings.
Thermodynamic systems are classified according to their ability to exchange mass and energy.
An open system allows the transfer of both mass and energy with its surroundings (e.g., a boiling tea kettle).
A closed system allows the transfer of energy but not mass (e.g., a pressure cooker).
An isolated system does not exchange matter or energy with its surroundings (e.g., a thermos flask).
Macroscopic properties describe thermodynamic systems using measurable bulk quantities.
Entropy (S) is a measure of the disorder of a system and is higher in liquids and gases than in solids.
The state of a system is defined by macroscopic properties such as temperature, pressure, volume, and the number of moles.
A state function depends only on the initial and final states, not on the path taken.
Enthalpy (H) equals internal energy (E) plus the product of pressure and volume (PV).
Exothermic reactions release heat, and ΔH is negative.
Endothermic reactions absorb heat, and ΔH is positive.
Enthalpy change (ΔH) represents the energy change at constant pressure.
The First Law of Thermodynamics states that energy cannot be created or destroyed; it can only be transferred or transformed.
ΔE = q + W, where ΔE is the change in internal energy, q is heat, and W is work.
Internal energy is the sum of the microscopic kinetic and potential energies of the particles in a system.
At constant volume, ΔE = qᵥ because no work is done.
At constant pressure, ΔH = qₚ, meaning that the enthalpy change equals the heat transferred.
For gaseous reactions, ΔH = ΔE + PΔV.
Hess’s Law states that the net enthalpy change is independent of the reaction pathway.
The standard enthalpy of formation (ΔH°f) is the enthalpy change when one mole of a compound is formed from its constituent elements in their standard states.
The standard enthalpy of formation (ΔH°f) of an element in its most stable form is zero.
ΔH can be determined using calorimetry or Hess’s Law.
The Born–Haber cycle relates the lattice energy of an ionic solid to other enthalpy changes.
Lattice energy (ΔHLE) is the energy released when one mole of an ionic solid lattice is formed from its gaseous ions.
