Std. 11
Chemistry

Chemistry studies matter, its composition, properties and transformations. The laws of conservation of mass, definite proportions, multiple proportions, reciprocal proportions and gaseous volumes provide the quantitative foundation of chemical combination.

One mole contains 6.022×10²³ entities. Molar mass is the mass of one mole. The number of moles is n = mass/molar mass.

Molarity is moles of solute per litre of solution: M = moles/L solution. Molality is moles of solute per kilogram of solvent: m = moles/kg solvent.

For component A, χA = nA/Σn. Mole fraction is dimensionless and all mole fractions in a mixture add to 1.

The empirical formula gives the simplest whole-number ratio of atoms. The molecular formula gives the actual number of atoms and is an integral multiple of the empirical formula.

A balanced chemical equation gives mole relationships. Convert the given quantity into moles, use the stoichiometric ratio, and convert to the requested quantity. Limiting reagent is the reactant consumed first and therefore limits product formation.

Significant figures communicate measurement precision. Dimensional analysis checks whether units are consistent and can help convert quantities safely.

Equivalent weight expresses combining capacity relative to a standard reaction. Normality relates gram equivalents of solute to volume of solution; use the prescribed formulae and reaction context.

Electrons, protons and neutrons form the basic atomic structure. Atomic number is the number of protons; mass number is protons plus neutrons.

Thomson proposed a positive sphere containing electrons; Rutherford's scattering experiment established a small dense nucleus; Bohr introduced quantised orbits for hydrogen-like atoms.

Light and matter show wave-particle behaviour. de Broglie's relation connects wavelength with momentum: λ=h/p.

Heisenberg's principle states that position and momentum cannot both be known with unlimited precision: Δx·Δp ≥ h/4π.

An orbital is a region of high probability of finding an electron, not a fixed path. Quantum numbers specify energy level, subshell, orbital orientation and spin.

Electrons fill orbitals according to the Aufbau principle, Pauli exclusion principle and Hund's rule. Half-filled and completely filled subshells can have extra stability.

s orbitals are spherical; p orbitals have dumbbell-like shapes in three orientations. d orbitals have more complex directional shapes.

The modern periodic law states that physical and chemical properties of elements are periodic functions of their atomic numbers.

Atomic radius generally decreases across a period because effective nuclear attraction increases, and increases down a group because additional shells are occupied.

Ionisation energy is the energy required to remove an electron from a gaseous atom. It generally increases across a period and decreases down a group, with important exceptions.

Electron affinity/electron-gain enthalpy describes the energy change associated with gaining an electron. Trends depend on atomic size, nuclear attraction and electron-electron repulsion.

Electronegativity is the tendency of an atom in a molecule to attract the shared electron pair. It generally increases across a period and decreases down a group.

Valency and common oxidation states reflect the electron arrangement and bonding behaviour of elements. Transition from one oxidation state to another can produce characteristic compounds.

Periodic trends help predict reactivity, bond character, acid-base behaviour of oxides and many other chemical properties.

An ionic bond results from electron transfer followed by electrostatic attraction between oppositely charged ions. Ionic solids have lattice structures and high melting points in many cases.

A covalent bond forms when atoms share electron pairs. Bond length, bond energy and bond order describe important features of the bond.

Lewis structures show valence electrons and shared/lone pairs. They help predict bonding and formal charge but are a model rather than a complete description of electron distribution.

Electron pairs around a central atom repel one another and arrange themselves to minimise repulsion. Lone pairs generally repel more strongly than bonding pairs, affecting molecular shape.

Hybridisation is a model that combines atomic orbitals to form equivalent hybrid orbitals. Common types include sp, sp² and sp³, associated with linear, trigonal planar and tetrahedral arrangements.

Atomic orbitals combine to form molecular orbitals. Electrons occupy bonding and antibonding orbitals; bond order helps indicate bond strength and stability.

Hydrogen bonding is an attractive interaction involving hydrogen bonded to a highly electronegative atom such as N, O or F. It influences boiling point, solubility and structure.

Dipole moment measures molecular polarity. For a simple dipole, μ=q×r. Molecular polarity depends on both bond polarities and molecular geometry.

The system is the portion of the universe being studied; everything else is the surroundings. A boundary separates them. Systems may be open, closed or isolated.

Internal energy, enthalpy, entropy and Gibbs energy are state functions: their changes depend only on initial and final states, not the path.

Energy is conserved. In chemistry, ΔU=q+w under the usual sign convention where work done on the system is positive.

At constant pressure, heat exchanged is related to enthalpy change: qₚ=ΔH. For a reaction, Hess's law allows enthalpy changes to be combined.

Entropy is associated with the dispersal of energy and the number of accessible microscopic arrangements. A process may be spontaneous when the overall entropy change is favourable.

At constant temperature and pressure, ΔG=ΔH−TΔS. A negative ΔG indicates thermodynamic spontaneity under those conditions; it does not mean the reaction is necessarily fast.

Calorimetry measures heat transfer. Specific heat and molar heat capacity relate heat to temperature change and amount of substance.

Formation, combustion, neutralisation, bond dissociation and solution enthalpies describe specific heat changes. Sign and conditions must always be stated.

In a reversible reaction at equilibrium, forward and reverse processes continue but occur at equal rates, so macroscopic concentrations remain constant.

For a general reaction, the equilibrium constant is written using equilibrium concentrations or activities: Kc = products/reactants, with powers equal to stoichiometric coefficients.

When an equilibrium system is disturbed, it shifts in a direction that tends to oppose the disturbance. Changes in concentration, pressure and temperature have different effects.

Weak electrolytes ionise only partially. Acid and base equilibria explain pH, ionisation constants and buffer action.

pH is related to hydrogen-ion concentration: pH=−log[H⁺]. Similarly, pOH relates to hydroxide concentration. At 25°C, pH+pOH=14 for dilute aqueous solutions.

A buffer resists large changes in pH when small amounts of acid or base are added. It commonly contains a weak acid and its conjugate base, or a weak base and its conjugate acid.

For a sparingly soluble salt, Ksp is the product of equilibrium ionic concentrations raised to their stoichiometric powers. Comparing ionic product with Ksp helps predict precipitation.

Adding an ion already present in an equilibrium suppresses ionisation of a weak electrolyte or decreases the solubility of a sparingly soluble electrolyte, depending on the system.

Oxidation involves loss of electrons/increase in oxidation number; reduction involves gain of electrons/decrease in oxidation number. They always occur together.

An oxidising agent accepts electrons and is itself reduced. A reducing agent donates electrons and is itself oxidised.

Oxidation number is a bookkeeping value assigned according to rules. It helps identify which species undergo oxidation or reduction.

Redox equations can be balanced using oxidation-number or ion-electron methods. The final equation must conserve both atoms and charge.

In disproportionation, the same species is simultaneously oxidised and reduced, producing two different oxidation states.

Redox principles explain combustion, corrosion, batteries, metallurgy and many biological oxidation processes.

Carbon forms four covalent bonds and bonds strongly with other carbon atoms, producing chains, branches and rings. This is the foundation of the enormous variety of organic compounds.

A functional group is an atom or group of atoms responsible for characteristic chemical behaviour. It provides a systematic way to classify organic compounds.

Organic compounds are named by selecting the parent chain/ring, identifying the principal functional group, numbering to give appropriate locants and adding substituent prefixes in the prescribed order.

Structural isomers have the same molecular formula but different connectivity. Stereoisomers have the same connectivity but differ in spatial arrangement.

Crystallisation, sublimation, distillation, differential extraction and chromatography separate or purify organic substances using differences in physical properties.

Organic compounds can be tested for elements such as nitrogen, sulphur and halogens by prescribed methods. Observations must be connected to the correct inference.

Inductive effect, resonance/mesomeric effect, hyperconjugation and electromeric effect help explain stability, polarity and reactivity of organic intermediates and compounds.

Carbocations, carbanions and free radicals are reactive species formed during organic reactions. Their stability depends on structure and electron distribution.

Alkanes are saturated hydrocarbons containing only single C–C bonds. For open-chain alkanes, the general formula is CₙH₂ₙ₊₂.

Alkenes contain a carbon-carbon double bond and are unsaturated. For simple open-chain monoalkenes, the general formula is CₙH₂ₙ.

Alkynes contain a carbon-carbon triple bond. For simple open-chain monoalkynes, the general formula is CₙH₂ₙ₋₂.

Benzene is a planar aromatic ring with delocalised π electrons. Its unusual stability explains why it commonly undergoes substitution rather than simple addition reactions.

Combustion produces carbon dioxide and water when complete. Substitution reactions, especially halogenation, occur under suitable light/heat conditions.

The π bond makes alkenes reactive toward addition. Hydrogenation, halogenation, hydrohalogenation, hydration and oxidation are important prescribed reaction patterns.

For addition of an unsymmetrical reagent such as HX to an unsymmetrical alkene under ordinary ionic conditions, the hydrogen generally adds to the carbon already bearing more hydrogen, and the other part to the more substituted carbon.

Alkynes undergo addition reactions at the triple bond and can be converted into other functional groups under suitable conditions.

Electrophilic substitution reactions such as nitration, sulphonation, halogenation and Friedel–Crafts reactions illustrate preservation of aromatic stability while replacing a ring hydrogen.