Chemistry Part I

NCERT Class 12 Chemistry Chapter 1 covers Solutions — homogeneous mixtures of two or more substances — including types of solutions, concentration units, Henry's law, Raoult's law, colligative properties (vapour pressure lowering, boiling point elevation, freezing point depression, osmotic pressure), and van't Hoff factor for solutes that associate or dissociate.

• 1Solutions are homogeneous mixtures; the component in largest quantity is the solvent, and concentration units include molarity (mol/L), molality (mol/kg), mole fraction, mass percentage, and ppm
• 2Henry's law states that at constant temperature the solubility of a gas in a liquid is directly proportional to its partial pressure above the liquid (p = KH x)
• 3Raoult's law states that the partial vapour pressure of each volatile component equals its mole fraction multiplied by its vapour pressure in the pure state (p1 = x1 p1°)
• 4Ideal solutions obey Raoult's law over the entire concentration range with zero enthalpy and volume of mixing; non-ideal solutions show positive or negative deviations and can form azeotropes
• 5Colligative properties — relative lowering of vapour pressure, elevation of boiling point (DTb = Kb m), depression of freezing point (DTf = Kf m), and osmotic pressure (P = CRT) — depend only on the number of solute particles, not their identity

Class 12 Chemistry Chapter 2, Electrochemistry, covers the interconversion of chemical and electrical energy, including galvanic cells, the Nernst equation, electrolytic conductance, Faraday's laws of electrolysis, batteries, fuel cells, and corrosion as an electrochemical process.

• 1A galvanic cell converts Gibbs energy of a spontaneous redox reaction into electrical work; the Daniell cell (Zn/Cu²⁺) has a standard EMF of 1.1 V.
• 2The Nernst equation E(cell) = E°(cell) – (RT/nF) ln Q relates cell potential to the reaction quotient; at 298 K it simplifies to E(cell) = E°(cell) – (0.059/n) log Q.
• 3Standard cell potential, standard Gibbs energy, and equilibrium constant are interrelated: DrG° = –nFE°(cell) and E°(cell) = (2.303RT/nF) log Kc.
• 4Molar conductivity (Λm = κ/c) increases on dilution for both strong and weak electrolytes; for strong electrolytes Λm = Λ°m – A√c (Debye–Hückel–Onsager equation).
• 5Kohlrausch's law states that limiting molar conductivity equals the sum of individual ionic contributions: Λ°m = ν₊λ°₊ + ν₋λ°₋; it is used to find Λ°m of weak electrolytes.

Chemical Kinetics is the branch of chemistry that studies the rates of chemical reactions and the factors — concentration, temperature, pressure, and catalysts — that control how fast reactions proceed.

• 1Rate of reaction is measured as the change in concentration of a reactant or product per unit time, expressed in mol L⁻¹ s⁻¹
• 2Rate law (Rate = k[A]^x[B]^y) must be determined experimentally; exponents need not equal stoichiometric coefficients
• 3Order of a reaction is the sum of powers of concentration terms in the rate law; it can be zero, fractional, or a whole number
• 4Half-life of a first-order reaction is t½ = 0.693/k, independent of initial concentration; for zero-order, t½ = [R]0/2k
• 5The Arrhenius equation k = Ae^(−Ea/RT) shows that increasing temperature or lowering activation energy increases the rate constant

Class 12 Chemistry Chapter 4 covers the d- and f-block elements — the transition metals (Groups 3–12) and the inner transition metals (lanthanoids and actinoids) — explaining their electronic configurations, variable oxidation states, magnetic behaviour, coloured ions, catalytic properties, and key compounds such as K₂Cr₂O₇ and KMnO₄.

• 1Transition metals are in Groups 3–12; their d orbitals are progressively filled across four series (3d, 4d, 5d, 6d). Zn, Cd and Hg are not true transition metals because their d orbitals are completely filled in ground state and common oxidation states.
• 2Variable oxidation states arise from incomplete d-orbital filling; manganese shows the widest range (+2 to +7). Oxidation states in d-block elements differ by one unit, unlike p-block elements where they differ by two.
• 3Transition metals are paramagnetic due to unpaired d electrons; magnetic moment is calculated using the spin-only formula µ = √(n(n+2)) BM, where n is the number of unpaired electrons.
• 4Coloured ions result from d–d electronic transitions when energy absorbed falls in the visible region; ions with completely filled or empty d orbitals (e.g., Sc³⁺, Zn²⁺) are colourless.
• 5Lanthanoid contraction — the steady decrease in atomic and ionic radii from La to Lu — causes the 5d elements (e.g., Hf, 159 pm) to have almost identical radii to their 4d counterparts (e.g., Zr, 160 pm), making their separation in nature difficult.

NCERT Class 12 Chemistry Chapter 5, Coordination Compounds, covers Werner's theory, nomenclature, isomerism, and bonding theories (VBT and CFT) for complexes where a central metal ion is bound to ligands in a definite geometric arrangement.

• 1Werner's theory distinguishes primary (ionisable) valence and secondary (non-ionisable) valence; the secondary valence equals the coordination number and determines geometry (octahedral, tetrahedral, square planar).
• 2Ligands are classified by denticity: unidentate (Cl⁻, NH₃), didentate (ethane-1,2-diamine, C₂O₄²⁻), polydentate (EDTA⁴⁻ is hexadentate), and ambidentate (NO₂⁻, SCN⁻ can bind through either of two atoms).
• 3Isomerism types include geometric (cis/trans, fac/mer), optical (non-superimposable mirror images in octahedral complexes with didentate ligands), linkage, coordination, ionisation, and solvate isomerism.
• 4Crystal Field Theory explains colour via d-d electronic transitions and magnetic behaviour via d-orbital splitting (Δo for octahedral, Δt = 4/9 Δo for tetrahedral); strong-field ligands cause low-spin and weak-field ligands cause high-spin configurations.
• 5Metal carbonyls exhibit synergic bonding: ligand-to-metal σ donation plus metal-to-ligand π back-donation into the CO antibonding π* orbital, stabilising the M–C bond.