Thermodynamics - Concepts, Laws, and Enthalpy Changes

Comprehensive vocabulary flashcards covering fundamental terms, state functions, laws of thermodynamics, and enthalpy types from Unit 5 Chemistry.

System

The specific part of the universe selected for thermodynamic study in which observations are made.

Surroundings

The entire portion of the universe outside the system that can interact with it, usually the region of space in its immediate neighbourhood.

Boundary

The real or imaginary wall separating a thermodynamic system from its surroundings, which regulates the flow of matter and energy.

Open System

A system that can exchange both matter and energy with its surroundings, such as reactants held in an open beaker.

Closed System

A system that allows the exchange of energy with its surroundings but prevents the exchange of matter, such as reactants sealed inside a conducting vessel.

Isolated System

A system in which there is no exchange of either matter or energy with the surroundings, such as reactants contained in a thermos flask.

State Functions (State Variables)

Macroscopic properties (such as pressure $p$, volume $V$, and temperature $T$) whose values depend only on the current state of the system and not on the pathway taken to reach that state.

Internal Energy ($U$)

The total energy contained within a system, representing the sum of all microscopic and macroscopic energy forms (chemical, electrical, mechanical, and thermal).

Adiabatic Process

A thermodynamic process in which no heat enters or leaves the system across its boundary ($q = 0$).

First Law of Thermodynamics

The statement of energy conservation asserting that the energy of an isolated system is constant, expressed mathematically as $\Delta U = q + w$.

Reversible Process

A change carried out infinitely slowly through a continuous series of equilibrium states such that the direction can be reversed at any stage by an infinitesimal adjustment.

Irreversible Process

A process that does not occur quasi-statically through equilibrium states and cannot be reversed by an infinitesimal change, representing all natural spontaneous processes.

Free Expansion

The expansion of a gas into an evacuated space or vacuum where external pressure is zero ($p_{\text{ex}} = 0$), resulting in zero work done ($w = 0$).

Enthalpy ($H$)

A thermodynamic state function defined as the sum of internal energy and the pressure-volume product: $H = U + pV$.

Extensive Property

A macroscopic property whose numerical value depends directly on the size or quantity of matter present in the system, such as mass, volume, internal energy, and enthalpy.

Intensive Property

A macroscopic property that is independent of the quantity or amount of matter present in the system, such as temperature, density, and pressure.

Heat Capacity ($C$)

The coefficient of proportionality between the heat transferred to a system and the resulting rise in temperature, defined by $q = C\Delta T$.

Molar Heat Capacity ($C_m$)

The amount of heat required to raise the temperature of one mole of a substance by $1\,^{\circ}\text{C}$ (or $1\,\text{K}$), given by $C_m = \frac{C}{n}$.

Specific Heat Capacity ($c$)

The quantity of heat required to raise the temperature of one unit mass of a substance by $1\,^{\circ}\text{C}$ (or $1\,\text{K}$).

$$C_p - C_V = R$$

The relation expressing the difference between the molar heat capacity at constant pressure ($C_p$) and constant volume ($C_V$) for one mole of an ideal gas.

Calorimetry

An experimental technique used to measure the heat changes and energy transferred during a physical or chemical process by monitoring temperature changes in a calibrated device called a calorimeter.

Bomb Calorimeter

A strong, sealed steel vessel immersed in a water bath designed to measure the heat of combustion at constant volume ($q_V = \Delta U$).

Reaction Enthalpy ($\Delta_r H$)

The net enthalpy change accompanying a chemical reaction, calculated as the sum of enthalpies of the products minus the sum of enthalpies of the reactants.

Standard State

The pure form of a substance in its specified state of aggregation under a pressure of exactly $1\,\text{bar}$ at a specified temperature (usually $298\,\text{K}$).

Standard Enthalpy of Fusion ($\Delta_{\text{fus}}H^{\ominus}$)

The enthalpy change accompanying the complete melting of one mole of a solid substance into liquid under standard conditions.

Standard Enthalpy of Vaporization ($\Delta_{\text{vap}}H^{\ominus}$)

The amount of heat required to convert one mole of a liquid into vapour at its boiling point under standard pressure ($1\,\text{bar}$).

Standard Enthalpy of Sublimation ($\Delta_{\text{sub}}H^{\ominus}$)

The enthalpy change when one mole of a solid substance transforms directly into its vapour phase under standard pressure at constant temperature.

Standard Molar Enthalpy of Formation ($\Delta_f H^{\ominus}$)

The enthalpy change accompanying the formation of one mole of a compound in its standard state directly from its constituent elements in their most stable reference states.

Thermochemical Equation

A balanced chemical equation that explicitly indicates the physical states of all participants alongside the corresponding reaction enthalpy ($\Delta_r H$).

Hess's Law of Constant Heat Summation

A principle stating that if a reaction is carried out in a series of steps, its overall standard reaction enthalpy is the algebraic sum of the standard enthalpies of the individual intermediate steps.

Standard Enthalpy of Combustion ($\Delta_c H^{\ominus}$)

The enthalpy change per mole of a substance when it undergoes complete combustion in excess oxygen, with all reactants and products in their standard states.

Enthalpy of Atomization ($\Delta_a H^{\ominus}$)

The enthalpy change that occurs when one mole of a substance has all its chemical bonds broken completely to yield gaseous atoms.

Bond Dissociation Enthalpy

The enthalpy change required to break one mole of a specific covalent bond in a gaseous diatomic or polyatomic compound into gaseous fragments.

Mean Bond Enthalpy

The average enthalpy value per mole required to break a particular type of chemical bond across the successive steps of dissociation in a polyatomic molecule.

Lattice Enthalpy ($\Delta_{\text{lattice}}H^{\ominus}$)

The enthalpy change that occurs when one mole of a solid ionic crystalline compound completely dissociates into its constituent gaseous ions.

Born-Haber Cycle

A closed thermochemical cycle based on Hess's law used to indirectly deduce the lattice enthalpy of an ionic crystal from measurable steps including sublimation, ionization, dissociation, and electron gain.

Enthalpy of Solution ($\Delta_{\text{sol}}H^{\ominus}$)

The enthalpy change observed when one mole of a substance dissolves in a specified quantity of solvent, equal to the sum of its lattice enthalpy and hydration enthalpy: $\Delta_{\text{sol}}H^{\ominus} = \Delta_{\text{lattice}}H^{\ominus} + \Delta_{\text{hyd}}H^{\ominus}$.

Enthalpy of Dilution

The heat change accompanying the addition of a specified amount of solvent to a solution, depending on the initial solution concentration and quantity of added solvent.

Spontaneous Process

A natural, irreversible process that possesses an inherent tendency to proceed in a forward direction without the continuous assistance of an external driving agency.

Entropy ($S$)

A thermodynamic state function that provides a quantitative measure of the degree of disorder, randomness, or chaotic energy dispersal in a system.

Second Law of Thermodynamics

The fundamental law stating that the total entropy change of the universe (system plus surroundings) must be positive ($\Delta S_{\text{total}} > 0$) for any spontaneous process to occur.

Third Law of Thermodynamics

The law stating that the entropy of a perfectly ordered, pure crystalline solid approaches zero as the absolute temperature approaches zero kelvin ($0\,\text{K}$).

Gibbs Energy ($G$)

A thermodynamic state function defined as $G = H - TS$, which represents the maximum portion of energy in a system available to perform useful non-mechanical work at constant temperature and pressure.

Gibbs Equation

The fundamental relation $\Delta G = \Delta H - T\Delta S$ used to evaluate the spontaneity of a process at constant temperature and pressure, where $\Delta G < 0$ indicates spontaneity.