Engineering Chemistry: Thermodynamics and Chemical Kinetics

Comprehensive practice flashcards covering Module 1: Thermodynamics and Chemical Kinetics, including terminology, laws of thermodynamics, entropy, Carnot cycle, and reaction mechanics as per the Fall Semester 2025-26 syllabus.

Thermodynamic System

The quantity of matter or a region in space upon which attention is concentrated in the analysis of a problem.

Intensive Properties

Properties of a system that are independent of the amount of matter present.

State Function

A property that depends only on the initial and final states of the system and is independent of the path taken, such as Temperature ($T$), Pressure ($P$), Internal energy ($U$), and Enthalpy ($H$).

Path Function

A property that depends on the path taken between the initial and final states, such as work done ($W$) and heat transferred ($q$).

Internal Energy ($U$)

The sum of kinetic energy and potential energy of a system; it is a state function and an extensive property.

Adiabatic Process

A thermodynamic process in which no heat is transferred into or out of the system.

Isothermal Process

A thermodynamic process that occurs at a constant temperature.

Isobaric Process

A thermodynamic process that occurs at a constant pressure.

Isochoric Process

A thermodynamic process that occurs at a constant volume, resulting in zero work done.

Reversible Process

A process that can be turned back such that both the system and surroundings return to their original states with no other change in the universe; it is reversed by an infinitesimal modification of a variable.

Zeroth Law of Thermodynamics

States that if two systems are in thermal equilibrium with a third system, then those two systems are in thermal equilibrium with each other; it forms the basis for temperature measurement.

First Law of Thermodynamics

The law of conservation of energy stating that the energy of an isolated system remains constant ($\Delta U = q + w$).

Enthalpy ($H$)

A thermodynamic quantity defined as $(U + PV)$; the change in enthalpy is given by $\Delta H = \Delta U + P\Delta V$ at constant pressure.

Heat Capacity ($C$)

The quantity of heat ($q$) required to raise the temperature of a system from $T_1$ to $T_2$ divided by the temperature difference.

Second Law of Thermodynamics

States that heat does not flow spontaneously from a cool body to a hotter body and that the entropy ($S$) of an isolated system increases during a spontaneous change.

Entropy ($S$)

A state function that measures the disorder of a system; its thermodynamic definition is $dS = \frac{dq_{rev}}{T}$.

Carnot Cycle

A reversible cyclic process consisting of four stages (isothermal expansion, adiabatic expansion, isothermal compression, and adiabatic compression) used to demonstrate maximum heat-to-work conversion.

Thermal Efficiency ($\eta$)

Calculated as $\eta = \frac{\text{work performed}}{\text{heat absorbed from the hot source}}$, given by the formula $1 - \frac{T_c}{T_h}$.

Gibbs Free Energy ($G$)

Defined as $G = H - TS$; a process is spontaneous at constant temperature and pressure if $\Delta G < 0$.

Gibbs-Helmholtz Equation

A fundamental thermodynamic relation connecting Gibbs free energy with enthalpy and entropy: $\Delta G = \Delta H - T\Delta S$.

Third Law of Thermodynamics

States that the entropy of all perfect crystalline substances is zero at $T = 0\,K$, where the number of microstates ($W$) equals $1$.

Chemical Kinetics

The study of reaction rates, the factors affecting them, and the step-by-step molecular-level mechanisms of reactions.

Rate Law

The relationship between the reaction rate and reactant concentrations, expressed as $\text{rate} = k[A]^m [B]^n$, where $k$ is the rate constant.

Order of Reaction

The sum of the exponents of the concentrations in the rate law equation, determined only through experimental observation.

Zero-Order Reaction

A reaction where the rate is independent of the concentration of the reactant ($\text{Rate} = k[A]^0$). Its half-life is $t_{1/2} = \frac{[A]_0}{2k}$.

First-Order Reaction

A reaction whose rate depends on the concentration of a single reactant raised to the first power. Its half-life is constant at $t_{1/2} = \frac{0.693}{k}$.

Pseudo-First-Order Reaction

A bimolecular reaction that behaves like a first-order reaction because one reactant is present in great excess, such as the hydrolysis of ester in water.

Activation Energy ($E_a$)

The minimum amount of energy required to initiate a chemical reaction by overcoming the energy barrier between reactants and the transition state.

Arrhenius Equation

Expresses the temperature dependence of the rate constant: $k = A e^{-\frac{E_a}{RT}}$, where $A$ is the pre-exponential factor.

Catalysis

The alteration of a reaction's rate by a substance (catalyst) that provides an alternative mechanism with a lower activation energy without being permanently changed itself.

Homogeneous Catalysis

A catalytic process where the catalyst and reactants are present in the same phase, such as the oxidation of $SO_2$ with $NO$ gas.

Heterogeneous Catalysis

A catalytic process where the catalyst and reactants are in different phases, such as the manufacture of ammonia by the Haber process using an iron ($Fe$) catalyst.

Enzyme Catalysis

Biological catalysis using protein-based molecules called enzymes that operate via a lock-and-key mechanism at an active site.