NEB Class 12 Chemistry cheat sheet

Rate law: rate = k[A]^m[B]^n; m and n are experimental orders, not stoichiometric coefficients. For a first-order reaction, k = (2.303/t)·log(a/(a−x)) and t½ = 0.693/k, independent of initial concentration.

Arrhenius equation: k = A·e^(−Ea/RT); log(k₂/k₁) = (Ea/2.303R)·(1/T₁ − 1/T₂). A catalyst lowers Ea without changing ΔH or the equilibrium constant.

Collision theory requires both sufficient energy (≥ Ea) and correct orientation; the steric factor P appears in rate = P·Z·e^(−Ea/RT).

First law: ΔU = q + w; work done on the system is positive. For an isothermal reversible expansion of an ideal gas, w = −nRT·ln(V₂/V₁) and ΔU = 0.

Enthalpy relations: ΔH = ΔU + Δn_g·RT, and Hess's law allows ΔH to be summed over any route. Bond breaking is endothermic; bond formation is exothermic.

Entropy and spontaneity: ΔG = ΔH − TΔS; a reaction is spontaneous when ΔG < 0, and ΔG° = −RT·ln K = −nFE°cell.

Equilibrium: Kp = Kc(RT)^Δn. Le Chatelier shifts oppose a change in concentration, pressure, or temperature; only temperature alters the value of K.

Colligative properties: ΔTb = i·Kb·m and ΔTf = i·Kf·m; osmotic pressure π = i·CRT. The van't Hoff factor i reflects dissociation or association in solution.

Electrochemistry: Nernst equation E = E° − (0.059/n)·log([products]/[reactants]) at 298 K; oxidation occurs at the anode in every cell type.

Periodic trends across a period: atomic radius and metallic character decrease, while ionisation energy, electron affinity, and electronegativity increase. Down a group the reverse holds.

Effective nuclear charge (Slater's rules) explains why ionisation energy rises across a period despite increasing shielding; noble gases have the highest IE in each period.

Diagonal relationship: Li–Mg, Be–Al, and B–Si show similar chemistry (e.g., both Be and Al form amphoteric oxides and covalent chlorides).

Hydrogen: ortho and para forms differ in nuclear spin; H₂O₂ acts as both oxidant and reductant, and its strength as an oxidant is greater in acidic medium.

Group 1 and 2: alkali metals show +1 only; alkaline earths show +2. Solubility of hydroxides increases down Group 2, while thermal stability of carbonates increases down the group.

Group 17: oxidising power decreases down the group (F₂ > Cl₂ > Br₂ > I₂); HF is a weak acid due to strong H–F hydrogen bonding, while HCl is strong.

Group 15–16: NH₃ is basic (lone pair on N), HNO₃ is a strong oxidising acid, and H₂SO₄ is a dehydrating and oxidising agent; SO₂ bleaches by reduction.

Transition metals: variable oxidation states, paramagnetism from unpaired d electrons, catalytic activity, and coloured ions from d–d transitions. Mn shows the widest range (+2 to +7).

Coordination compounds: IUPAC naming lists ligands alphabetically before the metal; EAN and VBT explain geometry, while CFSE explains colour and high/low spin splitting.

IUPAC priority for the principal functional group: carboxylic acid > ester > amide > nitrile > aldehyde > ketone > alcohol > amine > alkene/alkyne > halo/ether.

Nucleophilic substitution: SN1 proceeds via a carbocation (rate depends only on substrate, favoured by tertiary and polar protic solvents); SN2 is concerted and backside with inversion, favoured by primary and polar aprotic solvents.

Markovnikov's rule: in HX addition to an alkene, H adds to the carbon with more hydrogens. Peroxide effect (anti-Markovnikov) operates only with HBr, not HCl or HI.

Elimination (E1/E2) competes with substitution; E2 needs strong base and anti-periplanar geometry, and Saytzeff's rule favours the more substituted alkene.

Aromaticity (Hückel): cyclic, planar, fully conjugated, with 4n+2 π electrons. Benzene undergoes electrophilic substitution; substituents direct ortho/para if electron-donating, meta if electron-withdrawing.

Aldehydes are more reactive than ketones toward nucleophilic addition (less steric hindrance, less electron donation by alkyl groups). Tollens' and Fehling's tests distinguish aldehydes from ketones.

Carboxylic acid derivatives' reactivity: acid chloride > anhydride > ester > amide. Hydrolysis, alcoholysis, and ammonolysis all proceed by nucleophilic acyl substitution.

Alcohols: Lucas test distinguishes 1°, 2°, and 3° alcohols by turbidity time; oxidation gives aldehyde (1°) or ketone (2°), while 3° alcohols resist oxidation.

Amines are basic due to the lone pair on N; in aqueous solution, secondary amines are often most basic because of combined inductive and solvation effects. Hinsberg test distinguishes 1°, 2°, 3° amines.

Titration: at equivalence, moles of acid = moles of base (n₁V₁/n₂V₂ relation). Phenolphthalein suits weak acid–strong base; methyl orange suits strong acid–weak base.

Buffer: Henderson–Hasselbalch gives pH = pKa + log([salt]/[acid]); maximum buffer capacity occurs when pH = pKa.

Solubility product: Ksp = [A]^a[B]^b; precipitation begins when the ionic product exceeds Ksp. Common ions lower solubility.

Qualitative analysis: flame tests (Na golden yellow, K lilac, Ca brick red, Ba apple green, Cu blue-green); Group I cations precipitate as chlorides, Group II as sulphides in acidic medium.

Colorimetry uses the Beer–Lambert law A = εbc; absorbance is proportional to concentration, and a calibration curve is used for unknown samples.

Chromatography: separation depends on differential partition or adsorption between stationary and mobile phases; Rf = distance travelled by solute / distance travelled by solvent front.

Gravimetric analysis requires a pure, filterable, and constant-weight precipitate; the analyte mass is calculated from the gravimetric factor (molar mass ratio).

Alloys and materials: steel is Fe with controlled C; stainless steel adds Cr and Ni for corrosion resistance. Cement composition is expressed as oxides (CaO, SiO₂, Al₂O₃, Fe₂O₃).

Environmental chemistry: BOD measures biodegradable organic load; COD measures total oxidisable matter. Ozone depletion is driven by CFCs releasing Cl radicals; acid rain arises from SO₂ and NOx.