Class 11 NCERT Chemistry Foundations Review

A comprehensive set of vocabulary flashcards covering the core foundations of Class 11 NCERT Chemistry, including stoichiometry, atomic structure, periodic trends, bonding, thermodynamics, equilibrium, redox, and organic basics.

Mole

The SI unit for amount of substance; one mole contains exactly $6.022 \times 10^{23}$ elementary entities (Avogadro's number).

Avogadro's Constant ($N_A$)

The number of entities per mole of substance, defined as $6.022 \times 10^{23}\,mol^{-1}$.

Empirical Formula

The simplest whole-number ratio of atoms of each element present in a compound.

Molecular Formula

The formula representing the actual number of atoms of each element in a molecule; it is a whole-number multiple of the empirical formula.

Limiting Reagent

The reactant that is completely consumed first in a chemical reaction and determines the maximum amount of product that can form.

Law of Conservation of Mass

A law stating that matter is neither created nor destroyed in a chemical reaction; total mass of reactants equals total mass of products.

Law of Definite Proportions

A law stating that a given compound always contains exactly the same proportion of elements by mass, regardless of its source.

Law of Multiple Proportions

States that when two elements form more than one compound, the different masses of one element reacting with a fixed mass of the other are in a ratio of small whole numbers.

Atomic Mass Unit ($u$)

A mass unit defined as exactly $\frac{1}{12}$ the mass of a carbon-12 atom.

Isotopes

Atoms of the same element that share the same atomic number ($Z$) but differ in their neutron count and mass number ($A$).

Heisenberg's Uncertainty Principle

The principle stating that it is impossible to simultaneously know the exact position ($\Delta x$) and momentum ($\Delta p$) of an electron.

Quantum Numbers

A set of four numbers (principal $n$, azimuthal $l$, magnetic $m_l$, and spin $m_s$) that uniquely describe the energy, shape, and orientation of an orbital and electron spin.

Aufbau Principle

The rule that electrons fill atomic orbitals in order of increasing energy, following the $(n + l)$ rule.

Pauli Exclusion Principle

States that no two electrons in the same atom can have an identical set of all four quantum numbers; an orbital holds a maximum of $2$ electrons with opposite spins.

Hund's Rule of Maximum Multiplicity

States that for degenerate orbitals, electrons fill them singly with parallel spins before any pairing occurs.

Effective Nuclear Charge ($Z_{eff}$)

The net positive charge experienced by a valence electron after accounting for the shielding effect of inner-shell electrons.

Ionization Enthalpy ($\Delta_i H$)

The energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.

Electron Gain Enthalpy ($\Delta_{eg} H$)

The enthalpy change that occurs when an electron is added to a neutral gaseous atom to form a negative ion.

Electronegativity

A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself.

Fajans' Rules

Rules used to predict the covalent character of ionic bonds based on the size and charge of the cation and anion.

Formal Charge ($FC$)

The hypothetical charge assigned to an atom in a Lewis structure, calculated as $V - (L + \frac{1}{2}B)$, where $V$ is valence electrons, $L$ is lone-pair electrons, and $B$ is bonding electrons.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; a model used to predict molecular geometry by minimizing repulsion between electron pairs around a central atom.

Hybridization

The process of mixing atomic orbitals of slightly different energies to produce a new set of equivalent orbitals (such as $sp$, $sp^2$, or $sp^3$).

Sigma ($\sigma$) Bond

A covalent bond formed by the head-on overlap of atomic orbitals along the internuclear axis.

Pi ($\pi$) Bond

A covalent bond formed by the sideways overlap of atomic orbitals, occurring only after a sigma bond is established.

Resonance

A condition where a single Lewis structure cannot represent a molecule; the actual structure is a hybrid of multiple contributing structures.

Hydrogen Bond

A strong intermolecular attraction occurring when hydrogen is covalently bonded to a highly electronegative atom ($N$, $O$, or $F$).

Bond Order

In Molecular Orbital theory, the value calculated as $\frac{1}{2}(\text{bonding electrons } - \text{ antibonding electrons})$; higher values indicate stronger, shorter bonds.

State Function

A property whose value depends only on the current state of the system and not on the path taken (e.g., $U$, $H$, $S$, $G$).

First Law of Thermodynamics

The law of conservation of energy applied to a system, expressed as $\Delta U = q + w$.

Enthalpy ($H$)

A thermodynamic state function defined as $H = U + pV$; its change ($\Delta H$) equals the heat exchanged at constant pressure.

Hess's Law

States that if a reaction occurs in steps, the total enthalpy change is the sum of the enthalpy changes of the individual steps.

Entropy ($S$)

A measure of molecular disorder or the number of accessible microstates in a system.

Gibbs Free Energy ($G$)

A state function defined as $G = H - TS$; a negative change ($\Delta G < 0$) indicates a spontaneous process.

Dynamic Equilibrium

A state in a reversible reaction where the forward and reverse reaction rates are equal, resulting in constant macroscopic concentrations.

Le Chatelier's Principle

States that if a system at equilibrium is disturbed by a change in concentration, temperature, or pressure, the system shifts to counteract the change.

Br nsted-Lowry Acid/Base

An acid is a proton ($H^+$) donor and a base is a proton acceptor.

Lewis Acid/Base

An acid is an electron-pair acceptor and a base is an electron-pair donor.

pH Scale

A logarithmic scale used to specify the acidity or basicity of an aqueous solution, defined as $pH = -\log_{10}[H^+]$.

Buffer Solution

A solution that resists significant changes in $pH$ upon the addition of small amounts of acid or base, typically consisting of a weak acid and its conjugate base.

Solubility Product ($K_{sp}$)

The equilibrium constant for a sparingly soluble salt dissolving in water.

Common Ion Effect

The suppression of the ionization of a weak electrolyte or the solubility of a salt by the addition of a compound providing a shared ion.

Oxidation

The loss of electrons or an increase in the oxidation number of an element.

Reduction

The gain of electrons or a decrease in the oxidation number of an element.

Oxidation Number

The formal charge an atom would carry if all bonds were assumed to be purely ionic; used to track electron transfer in redox reactions.

Catenation

The unique ability of carbon atoms to form stable covalent bonds with other carbon atoms, resulting in long chains or rings.

Structural Isomers

Molecules with the same molecular formula but different connectivity or arrangements of atoms (e.g., chain, position, or functional group isomers).

Inductive Effect

The permanent polarization of a sigma bond due to the electronegativity difference between atoms, shifting electron density along a chain.

Hyperconjugation

The delocalization of $\sigma$-electrons from a $C-H$ or $C-C$ bond into an adjacent empty or partially filled $p$-orbital or $\pi$-orbital.

Markovnikov's Rule

A rule stating that in the addition of $HX$ to an unsymmetrical alkene, the $H$ atom attaches to the carbon with more hydrogen atoms already bonded to it.