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Complete NCERT Chemistry — Physical, Organic, Inorganic — Atomic Structure, Bonding, Thermodynamics, Reactions, Periodic Table.
| Concept | Formula / Value | Notes |
|---|---|---|
| Mole | 1 mole = 6.022 × 10²³ particles | Avogadro's number (Nₐ) |
| Molar Mass | Mass of 1 mole (g/mol) | Numerically = atomic/molecular mass |
| No. of moles (n) | n = W / M = N / Nₐ = V / 22.4 | W=mass, M=molar mass, N=particles, V=vol at STP |
| Molarity (M) | M = moles of solute / vol(L) | Changes with temperature |
| Molality (m) | m = moles of solute / mass(kg) of solvent | Independent of temperature |
| Mole Fraction (x) | x₁ = n₁ / (n₁ + n₂ + …) | Sum of all mole fractions = 1 |
| Normality (N) | N = equivalents / vol(L) | N = M × n-factor |
| % by Mass | (mass solute / mass solution) × 100 | Used for solid solutions |
| % by Volume | (vol solute / vol solution) × 100 | Used for liquid solutions |
─── Stoichiometry Example ───
2H₂ + O₂ → 2H₂O
If 4g H₂ reacts with 32g O₂:
moles of H₂ = 4/2 = 2
moles of O₂ = 32/32 = 1
Required ratio: H₂:O₂ = 2:1
Available: 2:1 → exact stoichiometric ratio
Moles of H₂O = 2 × 1 = 2 (from limiting O₂)
Mass of H₂O = 2 × 18 = 36g| Q.No. | Name | Symbol | Values | Significance |
|---|---|---|---|---|
| 1 | Principal | n | 1, 2, 3, … | Size & energy of orbital (shell) |
| 2 | Azimuthal | l | 0 to n−1 | Shape of orbital (s,p,d,f) |
| 3 | Magnetic | mₗ | −l to +l | Orientation in space |
| 4 | Spin | mₛ | +½ or −½ | Direction of electron spin |
─── Shapes of Orbitals ───
s-orbital: Spherical (l = 0)
1s (n=1): smallest, closest to nucleus
2s (n=2): larger, contains 1 radial node
Nodes = n − l − 1
p-orbital: Dumbbell shape (l = 1)
pₓ, pᵧ, p_z — 3 orientations (mₗ = −1, 0, +1)
Each has a nodal plane through the nucleus
d-orbital: Double dumbbell / clover leaf (l = 2)
d_xy, d_yz, d_xz, d_x²−y², d_z² — 5 orientations
f-orbital: Complex shapes (l = 3) — 7 orientations
─── Electronic Configuration Rules ───
Aufbau Principle: Fill lower energy orbitals first
Order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d …
Pauli Exclusion Principle: No two electrons can have the
same set of 4 quantum numbers (max 2e⁻ per orbital with
opposite spins: ↑↓)
Hund's Rule of Maximum Multiplicity: Pairing of electrons
in degenerate orbitals occurs only after each orbital has
one electron (with parallel spins)
Example: Nitrogen (Z=7): 1s² 2s² 2p³ → [↑] [↑] [↑]
Oxygen (Z=8): 1s² 2s² 2p⁴ → [↑↓] [↑] [↑]
Exceptions (Stability of half/fully filled):
Cr (Z=24): [Ar] 3d⁵ 4s¹ (not 3d⁴ 4s²)
Cu (Z=29): [Ar] 3d¹⁰ 4s¹ (not 3d⁹ 4s²)| Law | Relationship | Formula | Constant |
|---|---|---|---|
| Boyle's | P ∝ 1/V (T const) | PV = k | Temperature |
| Charles's | V ∝ T (P const) | V/T = k | Pressure |
| Gay-Lussac's | P ∝ T (V const) | P/T = k | Volume |
| Avogadro's | V ∝ n (T,P const) | V/n = k | T, P |
| Combined | PV ∝ nT | PV = nRT | R = gas constant |
─── Ideal Gas Equation ───
PV = nRT
R = 0.0821 L·atm/(mol·K) = 8.314 J/(mol·K)
= 0.0831 L·bar/(mol·K) = 1.987 cal/(mol·K)
= 62.36 L·mmHg/(mol·K)
─── Van der Waals Equation (Real Gas) ───
n²a
(P + ───)(V − nb) = nRT
V²
a → accounts for intermolecular forces
b → accounts for molecular volume
Larger a = stronger attraction; Larger b = bigger molecules
─── Kinetic Molecular Theory Postulates ───
1. Gases consist of tiny particles in constant random motion
2. Particle volume is negligible vs container volume
3. No intermolecular forces (ideal gas)
4. Collisions are perfectly elastic
5. Average KE ∝ absolute temperature
KE = 3/2 × RT/Nₐ per molecule = 3/2 kT
RMS speed: u_rms = √(3RT/M) where M = molar mass
Most probable: u_mp = √(2RT/M)
Average speed: u_avg = √(8RT/πM)
─── Liquefaction of Gases ───
Critical Temperature (Tc): Max temp above which gas
cannot be liquefied by pressure alone
Critical Pressure (Pc): Min pressure required to liquefy
gas at Tc
Tc > 25°C: Easily liquefiable (NH₃, CO₂, SO₂)
Tc < 25°C: Difficult to liquefy (H₂, O₂, N₂)| Term | Definition | Example |
|---|---|---|
| System | Part of universe under study | Reaction mixture in a beaker |
| Surroundings | Rest of universe | Everything outside the beaker |
| Open System | Exchanges both matter & energy | Open beaker |
| Closed System | Exchanges only energy | Sealed container |
| Isolated System | No exchange of matter or energy | Thermos flask |
| State Function | Depends only on initial & final state | ΔU, ΔH, ΔS, ΔG, P, V, T |
| Path Function | Depends on path taken | q (heat), w (work) |
| Extensive Property | Depends on amount of matter | Mass, volume, enthalpy |
| Intensive Property | Independent of amount | Temperature, pressure, density |
─── First Law of Thermodynamics ───
ΔU = q + w
(Energy can neither be created nor destroyed)
Work done by gas: w = −PΔV (at constant pressure)
For ideal gas: w = −nRT ln(V₂/V₁) (isothermal)
─── Enthalpy (H) ───
H = U + PV
ΔH = ΔU + Δn_g × RT (at constant T & P)
Δn_g = moles(gas products) − moles(gas reactants)
─── Hess's Law ───
If a reaction occurs in steps, total enthalpy change =
sum of enthalpy changes of individual steps.
C(s) + ½O₂(g) → CO(g) ΔH₁ = −110.5 kJ
CO(g) + ½O₂(g) → CO₂(g) ΔH₂ = −283.0 kJ
─────────────────────────────────────────────
C(s) + O₂(g) → CO₂(g) ΔH = −393.5 kJ
─── Standard Enthalpy Types ───
ΔH°f (formation) → 1 mol compound from elements
ΔH°c (combustion) → 1 mol substance burnt completely
ΔH°sub (sublimation) → solid → gas
ΔH°vap (vaporization) → liquid → gas
ΔH°fus (fusion) → solid → liquid
ΔH°sol (solution) → dissolving in solvent
ΔH°hyd (hydration) → gas ion + water → hydrated ion
ΔH°atom (atomization) → elements → gaseous atoms
Bond enthalpy → avg energy to break 1 mol bonds in gas
─── Spontaneity & Gibbs Energy ───
ΔG = ΔH − TΔS
ΔG < 0 → Spontaneous (thermodynamically favourable)
ΔG = 0 → Equilibrium
ΔG > 0 → Non-spontaneous
At equilibrium: ΔG° = −RT ln K
K = equilibrium constant
R = 8.314 J/(mol·K), T = temperature (K)
─── Entropy (S) ───
Measure of disorder/randomness
S(gas) > S(liquid) > S(solid)
ΔS° = ΣS°(products) − ΣS°(reactants)| Concept | Definition |
|---|---|
| Oxidation | Loss of electrons; increase in oxidation number |
| Reduction | Gain of electrons; decrease in oxidation number |
| Oxidizing Agent | Species that gets reduced (accepts electrons) |
| Reducing Agent | Species that gets oxidized (donates electrons) |
| Disproportionation | Same element is both oxidized and reduced |
| Comproportionation | Two species with same element in different states combine |
─── Oxidation Number Rules ───
1. Elements in free state: O.N. = 0 (Na, H₂, O₂, Cl₂, S₈)
2. Monoatomic ion: O.N. = charge (Na⁺ → +1, O²⁻ → −2)
3. Oxygen: usually −2 (except: peroxides −1, OF₂ +2, KO₂ −½)
4. Hydrogen: usually +1 (except: metal hydrides −1)
5. Halogens: usually −1 (except: with O or more electronegative halogen)
6. Alkali metals: +1; Alkaline earth: +2
7. Sum of O.N. in neutral molecule = 0; in ion = charge
─── Balancing by Oxidation Number Method ───
Step 1: Write skeletal equation
Step 2: Assign oxidation numbers
Step 3: Calculate increase & decrease in O.N.
Step 4: Balance electrons by multiplying
Step 5: Balance all atoms
Step 6: Balance O by adding H₂O (acidic) or OH⁻ (basic)
Step 7: Balance H by adding H⁺ (acidic) or H₂O (basic)
─── Ion-Electron Method (Half-Reaction) ───
Step 1: Split into oxidation & reduction half-reactions
Step 2: Balance atoms (except O and H)
Step 3: Balance O by adding H₂O
Step 4: Balance H by adding H⁺ (acidic) or OH⁻ (basic)
Step 5: Balance charge by adding electrons
Step 6: Multiply to equalize electrons, then add half-reactions
─── Equivalent Weight ───
For acids: Eq. wt = Molar mass / basicity
For bases: Eq. wt = Molar mass / acidity
For salts: Eq. wt = Molar mass / total +ve charge
For redox: Eq. wt = Molar mass / change in O.N. per formula unit| Property | Ideal Solution | Non-Ideal (+ve deviation) | Non-Ideal (−ve deviation) |
|---|---|---|---|
| ΔH_mix | Zero | Positive (endothermic) | Negative (exothermic) |
| ΔV_mix | Zero | Positive (expansion) | Negative (contraction) |
| A-B forces | A-A = B-B = A-B | A-B < A-A, B-B | A-B > A-A, B-B |
| Example | Benzene-Toluene | Ethanol-Water | Acetone-Chloroform |
─── Raoult's Law ───
For volatile liquids:
P_total = x_A × P°_A + x_B × P°_B
P° = vapour pressure of pure component
x = mole fraction in liquid phase
For non-volatile solute:
(P° − P_s) / P° = x₂ (mole fraction of solute)
→ Relative lowering of VP = x₂
─── Colligative Properties (depend on no. of particles) ───
1. Relative Lowering of Vapour Pressure:
ΔP/P° = x₂ = n₂ / (n₁ + n₂) ≈ n₂/n₁ (dilute)
2. Elevation in Boiling Point:
ΔTb = Kb × m
Kb = molal boiling point elevation constant
Kb = RT_b² × M / (1000 × ΔH_vap)
3. Depression in Freezing Point:
ΔTf = Kf × m
Kf = molal freezing point depression constant
Kf = RT_f² × M / (1000 × ΔH_fus)
4. Osmotic Pressure:
π = CRT (for dilute solutions)
π = iCRT (for electrolytes)
C = molarity (mol/L), R = 0.0821, T = K
─── Van't Hoff Factor (i) ───
i = observed value / calculated value (without dissociation)
i = 1 + (n − 1)α for association/dissociation
For NaCl: i ≈ 2 (Na⁺ + Cl⁻)
For CaCl₂: i ≈ 3 (Ca²⁺ + 2Cl⁻)
For Al₂(SO₄)₃: i ≈ 5 (2Al³⁺ + 3SO₄²⁻)
Modified colligative properties:
ΔTb = i × Kb × m
ΔTf = i × Kf × m
π = i × CRT| Order | Rate Law | Units of k | Half-life | Plot |
|---|---|---|---|---|
| Zero | r = k | mol L⁻¹ s⁻¹ | t½ = [A]₀/(2k) | [A] vs t (linear) |
| First | r = k[A] | s⁻¹ | t½ = 0.693/k | ln[A] vs t (linear) |
| Second | r = k[A]² | L mol⁻¹ s⁻¹ | t½ = 1/(k[A]₀) | 1/[A] vs t (linear) |
| Pseudo 1st | r = k'[A] | s⁻¹ | t½ = 0.693/k' | Same as 1st order |
─── First Order Kinetics (most common in NCERT) ───
Rate Law: r = k[A]
Integrated Rate Law: k = (1/t) ln([A]₀/[A])
or k = (2.303/t) log([A]₀/[A])
[A] = [A]₀ × e^(−kt)
Half-life: t½ = 0.693 / k (independent of [A]₀)
Life of reaction: t_99% = 2.303/k × log(100/1) = 4.606/k
─── Arrhenius Equation ───
k = A × e^(−Ea/RT)
ln k = ln A − Ea/(RT)
log(k₂/k₁) = Ea/(2.303R) × (1/T₁ − 1/T₂)
A = pre-exponential (frequency) factor
Ea = activation energy (J/mol)
R = 8.314 J/(mol·K)
T = temperature (K)
Plot: ln k vs 1/T → straight line
slope = −Ea/R
intercept = ln A
─── Key Concepts ───
Order vs Molecularity:
Order: experimentally determined, can be zero/fractional
Molecularity: no. of molecules colliding in elementary step,
always a whole number, applies only to elementary reactions
Rate determining step: slowest step of a complex reaction
determines the overall rate law
Activation energy: minimum energy required for reaction
Catalyst: lowers Ea without being consumed, provides
alternate pathway with lower activation energy
Temperature dependence: rate roughly doubles for every
10°C rise (van't Hoff rule)| Concept | Definition / Details |
|---|---|
| Adsorption | Accumulation of molecular species on surface (not bulk); ΔH < 0 always |
| Physisorption | Weak van der Waals forces; non-specific; reversible; low ΔH; no activation energy |
| Chemisorption | Strong chemical bonds; specific; irreversible; high ΔH; may have activation energy |
| Freundlich Isotherm | x/m = kP^(1/n) (adsorption isotherm for gases) |
| Langmuir Isotherm | θ = KP/(1 + KP) — monolayer adsorption model |
| Catalyst | Substance that alters rate without being consumed; lowers Ea via alternate pathway |
| Homogeneous | Catalyst and reactants in same phase (e.g., SO₂ + O₂ with NO gas) |
| Heterogeneous | Catalyst in different phase (e.g., Fe in Haber process, V₂O₅ in Contact process) |
| Enzymes | Biological catalysts; highly specific; work at body temp; lock-and-key mechanism |
─── Colloids ───
Size range: 1 nm to 1000 nm (between true solution and suspension)
Types of Colloids:
Dispersion Medium / Dispersed Phase
Sol: solid in liquid (e.g., paint, mud)
Gel: liquid in solid (e.g., butter, jelly)
Foam: gas in liquid (e.g., shaving cream) or solid
Emulsion: liquid in liquid (e.g., milk, cod liver oil)
Aerosol: liquid/solid in gas (e.g., fog, smoke)
─── Key Colloidal Properties ───
Tyndall Effect: Scattering of light by colloidal particles;
path of light becomes visible
Brownian Motion: Random zig-zag movement of colloidal particles
due to unbalanced bombardment by molecules
Coagulation: Destabilizing colloid to precipitate by:
- Adding electrolyte (Hardy-Schulze rule)
- Heating, mixing oppositely charged colloids
Electrophoresis: Movement of charged particles towards
opposite electrode under electric field
Dialysis: Removal of dissolved ions using semipermeable
membrane (only ions pass, colloids don't)
Emulsifying Agent: Stabilizes emulsion (e.g., soap for O/W,
heavy metal salts for W/O emulsions)| Method | Used For | Principle |
|---|---|---|
| Crystallization | Solids | Dissolve in min. hot solvent, cool to crystallize |
| Sublimation | Volatile solids | Direct solid → gas → solid (camphor, naphthalene) |
| Distillation | Liquids with different BP | Boil and condense based on boiling points |
| Fractional Distillation | Close BP liquids | Use fractionating column for better separation |
| Steam Distillation | Immiscible liquids | Lowers BP for heat-sensitive compounds (oils) |
| Chromatography | All types | Differential partition between stationary/mobile phase |
| Differential Extraction | Organic from aqueous | Use separating funnel with immiscible solvents |
─── Lassaigne's Test (Detection of N, S, Halogens) ───
Step 1: Fuse organic compound with Na metal → NaCN, Na₂S, NaX
Step 2: Extract with water (Lassaigne's extract / sodium fusion extract)
Test for Nitrogen:
Add FeSO₄, boil → add conc. H₂SO₄
Prussian blue precipitate → N present
(NaCN + FeSO₄ → Na₄[Fe(CN)₆] + Fe³⁺ → Fe₄[Fe(CN)₆]₃)
Test for Sulphur:
Add sodium nitroprusside → violet colour → S present
OR add Pb(CH₃COO)₂ → black precipitate of PbS
Test for Nitrogen AND Sulphur both:
Add FeSO₄ → boil → add conc. H₂SO₄ → blood red colour
(sodium thiocyanate formed: NaSCN + Fe³⁺ → blood red)
Test for Halogens:
Add HNO₃ (to remove CN⁻/S²⁻), then AgNO₃
Cl⁻ → white ppt (AgCl, soluble in NH₃)
Br⁻ → pale yellow ppt (AgBr, partially soluble in NH₃)
I⁻ → yellow ppt (AgI, insoluble in NH₃)
─── Quantitative Analysis ───
C, H: Liebig's combustion method
CO₂ absorbed by KOH → C estimated
H₂O absorbed by CaCl₂ → H estimated
N: Dumas method (free N₂ measured) or Kjeldahl method
(NH₃ titrated for N in organic compounds)
Halogens: Carius method (AgX precipitate weighed)
S: Carius method (BaSO₄ precipitate weighed)
P: Carius method (Mg₂P₂O₇ weighed)| Priority (Highest → Lowest) | Functional Group | Suffix / Prefix |
|---|---|---|
| 1 | Carboxylic acid | −oic acid |
| 2 | Anhydride | −oic anhydride |
| 3 | Ester | −oate |
| 4 | Acid halide | −oyl chloride |
| 5 | Amide | −amide |
| 6 | Nitrile | −nitrile / −carbonitrile |
| 7 | Aldehyde | −al |
| 8 | Ketone | −one |
| 9 | Alcohol | −ol |
| 10 | Amine | −amine |
| 11 | Alkene | −ene |
| 12 | Alkyne | −yne |
| 13 | Ether | alkoxy− (prefix) |
| 14 | Halo / Nitro | halo− / nitro− (prefix) |
─── Structural Isomerism ───
1. Chain Isomerism: Different carbon skeleton
C₄H₁₀: n-butane vs isobutane
2. Position Isomerism: Same chain, different position of FG
CH₃CH₂CH₂OH vs CH₃CH(OH)CH₃ (1-propanol vs 2-propanol)
3. Functional Group Isomerism: Different FG, same formula
C₃H₆O: propanal (aldehyde) vs propanone (ketone) vs
prop-2-en-1-ol (enol)
4. Metamerism: Different alkyl groups on either side of
same functional group (ethers, amines)
C₃H₈O: CH₃-O-C₂H₅ vs C₂H₅-O-CH₃ (same here, but
C₄H₁₀O: CH₃OC₃H₇ vs C₂H₅OC₂H₅)
5. Tautomerism: Structural isomers in dynamic equilibrium
Keto-enol tautomerism:
CH₃COCH₃ (keto) ⇌ CH₂=C(OH)CH₃ (enol)
─── Stereoisomerism ───
1. Geometrical (cis-trans / E-Z):
Due to restricted rotation around C=C or ring
cis: same side; trans: opposite side
E-Z system: priority by Cahn-Ingold-Prelog rules
2. Optical Isomerism:
Chiral molecules (no plane of symmetry)
Rotate plane-polarized light
d (+): clockwise rotation
l (−): anti-clockwise rotation
Racemic mixture: 50:50 d and l → optically inactive
Meso compound: chiral centers but optically inactive
due to internal plane of symmetry─── Alkanes (Saturated Hydrocarbons) ───
General Formula: CₙH₂ₙ₊₂
Hybridization: sp³ (tetrahedral, bond angle 109.5°)
All C-C bonds are sigma (σ) bonds
Conformations of Ethane:
Eclipsed: maximum energy, H atoms aligned
Staggered: minimum energy, H atoms offset (most stable)
Energy barrier: 12.5 kJ/mol (torsional strain)
Sawhorse & Newman projections used to represent
Isomerism in Alkanes:
C₄: 2 isomers (n-butane, isobutane)
C₅: 3 isomers
C₆: 5 isomers
C₇: 9 isomers
─── Preparation ───
1. Wurtz Reaction: 2R-Br + 2Na → R-R + 2NaBr
2. Decarboxylation: RCOONa + NaOH → RH + Na₂CO₃
3. Kolbe's electrolysis: 2RCOONa → R-R + 2CO₂ + H₂ + Na₂
4. Sabatier-Senderens: RCOOH + H₂ → RCH₃ + H₂O (Ni catalyst)
5. Reduction: R-X + 2[H] → R-H + HX (Zn/HCl, LiAlH₄)
─── Chemical Reactions ───
Halogenation (Free radical substitution):
CH₄ + Cl₂ → hv → CH₃Cl → CH₂Cl₂ → CHCl₃ → CCl₄
Mechanism: Initiation → Propagation → Termination
Combustion: CₙH₂ₙ₊₂ + (3n+1)/2 O₂ → nCO₂ + (n+1)H₂O
Pyrolysis: Thermal decomposition → cracking (alkanes → smaller)
Reforming: Straight chain → branched (better fuel octane)─── Alkenes ───
General Formula: CₙH₂ₙ
Hybridization: sp² (trigonal planar, bond angle ~120°)
Contains C=C double bond (1σ + 1π bond)
Preparation:
Dehydration: RCH₂CH₂OH → Al₂O₃/Δ → RCH=CH₂ + H₂O
Dehydrohalogenation: RCH₂CH₂Br + KOH(alc) → RCH=CH₂ + KBr + H₂O
(Saytzeff rule: more substituted alkene is major product)
Addition Reactions:
Hydrogenation: RCH=CH₂ + H₂ → Pd/C → RCH₂CH₃
Halogenation: RCH=CH₂ + X₂ → RCHX-CH₂X (anti addition)
Hydrohalogenation: RCH=CH₂ + HX → RCHX-CH₃
Markovnikov's Rule: H adds to C with more H atoms
Anti-Markovnikov (Peroxide effect/Kharasch):
HBr + peroxide → H adds to C with fewer H atoms
Hydration: RCH=CH₂ + H₂O → H⁺ → RCH(OH)CH₃
Ozonolysis: RCH=CHR' → O₃ → Zn/H₂O → RCHO + R'CHO
(used to find position of double bond)
─── Alkynes ───
General Formula: CₙH₂ₙ₋₂
Hybridization: sp (linear, bond angle 180°)
Contains C≡C triple bond (1σ + 2π bonds)
Acidic Hydrogen: Terminal alkynes (−C≡CH) are acidic
because sp carbon is more electronegative
React with: NaNH₂, Na, Grignard reagent, etc.
Preparation:
Dehydrohalogenation: vicinal dihalide → KOH(alc)/Δ → alkyne
CaC₂ + 2H₂O → Ca(OH)₂ + HC≡CH (acetylene)
Reactions:
Hydrogenation: HC≡CH + 2H₂ → Ni → H₃C−CH₃
(Lindlar's catalyst → cis-alkene; Na/liq NH₃ → trans-alkene)
Addition: HC≡CH + H₂O → HgSO₄/H₂SO₄ → CH₃CHO (ethanal)─── Benzene (C₆H₆) ───
Structure: Planar hexagonal ring, all C-C bonds equal (139 pm)
Delocalized π electrons (6e⁻, 4n+2 Huckel rule: aromatic)
Resonance energy: 150.4 kJ/mol (extra stability)
─── Electrophilic Aromatic Substitution (EAS) ───
General Mechanism:
Step 1: Electrophile (E⁺) generation
Step 2: Attack on benzene → arenium ion (σ complex)
Step 3: Loss of H⁺ to restore aromaticity
1. Nitration: C₆H₆ + HNO₃ → H₂SO₄ → C₆H₅NO₂ + H₂O
Electrophile: NO₂⁺ (nitronium ion)
2. Sulphonation: C₆H₆ + H₂SO₄(conc) → C₆H₅SO₃H + H₂O
Electrophile: SO₃H⁺ (sulphonium ion)
3. Halogenation: C₆H₆ + X₂ → FeCl₃ → C₆H₅X + HX
Electrophile: X⁺ (generated by Lewis acid catalyst)
4. Friedel-Crafts Alkylation:
C₆H₆ + RCl → AlCl₃ → C₆H₅R + HCl
Electrophile: R⁺ (carbocation)
Note: Rearrangement possible, polyalkylation
5. Friedel-Crafts Acylation:
C₆H₆ + RCOCl → AlCl₃ → C₆H₅COR + HCl
Electrophile: RCO⁺ (acylium ion)
Note: No rearrangement, mono-substituted
─── Directive Influence of Groups ───
Ortho-Para Directors (activating, except halogens):
−OH, −OCH₃, −NH₂, −NHR, −NR₂, −CH₃, −C₂H₅, −X (halogens)
Meta Directors (deactivating):
−NO₂, −CN, −CHO, −COOH, −COOR, −SO₃H, −COR, −NR₃⁺─── Haloalkanes: Classification ───
By no. of halogen atoms:
Monohalogen (R-X), Dihalogen (R-X₂), Trihalogen (CHX₃, CX₄)
By type of C attached:
1° (primary), 2° (secondary), 3° (tertiary)
By hybridization of C-X:
sp³ (alkyl halides), sp² (vinyl halides), sp (aryl halides)
─── SN1 vs SN2 Mechanisms ───
SN2 (Substitution Nucleophilic Bimolecular):
• One-step mechanism (concerted)
• Rate = k[RX][Nu⁻] (second order)
• Inversion of configuration (Walden inversion)
• Favoured by: 1° halide, strong nucleophile, polar aprotic
• Stereospecific: single transition state
SN1 (Substitution Nucleophilic Unimolecular):
• Two-step: carbocation intermediate
• Rate = k[RX] (first order)
• Racemization (partial inversion + retention)
• Favoured by: 3° halide, weak nucleophile, polar protic
• Rearrangement possible
Reactivity order for SN1: 3° > 2° > 1°
Reactivity order for SN2: CH₃ > 1° > 2° > 3°
─── E1 vs E2 Mechanisms ───
E2 (Elimination Bimolecular):
• One-step, anti-periplanar elimination
• Rate = k[RX][Base]
• Saytzeff product (more substituted alkene) is major
• Favoured by: strong base, higher temperature
E1 (Elimination Unimolecular):
• Two-step via carbocation
• Rate = k[RX]
• Saytzeff product major
• Favoured by: 3° halide, weak base
Note: Strong bulky bases favour E2 over SN2
(e.g., t-BuOK favours elimination)| Type | Formula | OH Position | Acidity | Key Feature |
|---|---|---|---|---|
| 1° Alcohol | RCH₂OH | Primary C | pKa ~16 | Oxidized to aldehyde → carboxylic acid |
| 2° Alcohol | R₂CHOH | Secondary C | pKa ~17 | Oxidized to ketone only |
| 3° Alcohol | R₃COH | Tertiary C | pKa ~18 | Resistant to oxidation |
| Phenol | C₆H₅OH | Attached to benzene | pKa ~10 | Acidic, electrophilic substitution |
| Ether | R−O−R | Bridge between C | Not acidic | Inert, Williamson synthesis |
─── Important Reactions ───
Lucas Test (Distinguishes 1°, 2°, 3° alcohols):
R-OH + HCl/ZnCl₂ → R-Cl + H₂O
3° alcohol → turbid immediately
2° alcohol → turbid in 5–10 min
1° alcohol → no turbidity at room temp
Dehydration of Alcohols:
RCH₂CH₂OH → Al₂O₃/573K → RCH=CH₂ + H₂O
(Saytzeff rule: more substituted alkene preferred)
Phenol Preparation:
Cumene process: Cumene + O₂ → cumene hydroperoxide
→ acid → phenol + acetone
Phenol Reactions:
• With Br₂/H₂O → 2,4,6-tribromophenol (white ppt)
• With NaOH → sodium phenoxide (phenol is acidic)
• Reimer-Tiemann: Phenol + CHCl₃ + NaOH → salicylaldehyde
Ether Preparation (Williamson Synthesis):
R-ONa + R'-X → R-O-R' + NaX
Use 1° alkyl halide (avoids elimination)
Cleavage of Ethers:
R-O-R' + HI → RI + R'-OH (excess HI → R-I + R'-I)
Mixed ethers: bond breaks at less hindered side─── Nucleophilic Addition Reactions ───
Aldehydes > Ketones (more reactive due to less steric hindrance
+ electronic effect)
With HCN: RCHO + HCN → RCH(OH)CN (cyanohydrin)
With NaHSO₃: RCHO + NaHSO₃ → RCH(OH)SO₃Na (white ppt)
With Grignard: RCHO + RMgX → RCH(OMgX)R → RCH(OH)R
With NH₂OH: RCHO + NH₂OH → RCH=NOH + H₂O (oxime)
With NH₂NH₂: RCHO + NH₂NH₂ → RCH=NHNH₂ (hydrazone)
With 2,4-DNP: RCHO + H₂NNHCOC₆H₃(NO₂)₂ → derivative
(yellow/orange ppt — used for identification)
─── Aldol Condensation ───
Aldehydes/ketones with α-H undergo self-condensation:
2 CH₃CHO → dil. NaOH → CH₃CH(OH)CH₂CHO (aldol)
→ heat → CH₃CH=CHCHO + H₂O (crotonaldehyde)
Cross aldol: between two different aldehydes/ketones
(one must have α-H)
─── Cannizzaro Reaction ───
Aldehydes without α-H (HCHO, C₆H₅CHO) in conc. NaOH:
2 HCHO + NaOH → CH₃OH + HCOONa
(disproportionation: one oxidized to acid, one reduced to alcohol)
─── Carboxylic Acid Reactions ───
Esterification: RCOOH + R'OH ⇌ H⁺ → RCOOR' + H₂O
Hell-Volhard-Zelinsky: RCOOH + Br₂ → PBr₃ → RCH(Br)COOH
(α-halogenation of carboxylic acids)
Decarboxylation: RCOONa + NaOH/CaO → RH + Na₂CO₃
Conversion: RCOOH → SOCl₂ → RCOCl (acid chloride)
→ RCOCl + NH₃ → RCONH₂ (amide)─── Classification ───
1°: RNH₂ (primary amine, e.g., CH₃NH₂ — methylamine)
2°: R₂NH (secondary amine, e.g., (CH₃)₂NH — dimethylamine)
3°: R₃N (tertiary amine, e.g., (CH₃)₃N — trimethylamine)
─── Basicity Order ───
In gas phase: 3° > 2° > 1° > NH₃
In aqueous phase: 2° > 1° > 3° > NH₃
(Due to steric hindrance + solvation effect)
Aromatic amines: PhNH₂ < NH₃ < alkyl amines
(−NH₂ is ortho-para director but deactivating)
─── Preparation Methods ───
1. Hofmann Bromamide Degradation:
RCONH₂ + Br₂ + 4NaOH → RNH₂ + Na₂CO₃ + 2NaBr + 2H₂O
(amide → 1° amine, loses 1 carbon)
2. Gabriel Phthalimide Synthesis:
Phthalimide + KOH → potassium phthalimide + R-X →
N-alkylphthalimide → hydrazine → primary amine
(gives pure 1° amine, no 2°/3° contamination)
3. Reduction: RNO₂ + LiAlH₄/H₂/Pd → RNH₂
RCN + LiAlH₄ → RCH₂NH₂
RCONH₂ + LiAlH₄ → RCH₂NH₂
4. Carbylamine (Isocyanide) Test:
RNH₂ + CHCl₃ + 3KOH → RNC (isocyanide) + 3KCl + 3H₂O
(offensive odour, 1° and 2° amines only)
─── Diazonium Salts ───
ArNH₂ + NaNO₂ + 2HCl → 0-5°C → ArN₂⁺Cl⁻ + NaCl + 2H₂O
(Benzene diazonium chloride)
Sandmeyer Reaction:
ArN₂⁺Cl⁻ + CuCl → ArCl (chlorobenzene)
ArN₂⁺Cl⁻ + CuBr → ArBr (bromobenzene)
ArN₂⁺Cl⁻ + CuCN → ArCN (benzonitrile)
Gattermann Reaction (similar to Sandmeyer, uses Cu powder)
ArN₂⁺Cl⁻ + KI → ArI (iodobenzene, no catalyst needed)
ArN₂⁺Cl⁻ + HBF₄ → ArF (fluorobenzene, Balz-Schiemann)
ArN₂⁺Cl⁻ + H₂O → ArOH (phenol, by warming)
ArN₂⁺Cl⁻ + H₃PO₂ → ArH (benzene, reduction)| Biomolecule | Building Blocks | Key Features | Linkage |
|---|---|---|---|
| Carbohydrates | Monosaccharides (glucose, fructose) | Polyhydroxy aldehydes/ketones, (CH₂O)ₙ | Glycosidic bond |
| Proteins | Amino acids (20 types) | Peptide bond, 4 structure levels | Peptide bond (−CO−NH−) |
| Nucleic Acids | Nucleotides (pentose + N-base + PO₄) | DNA (deoxyribose) & RNA (ribose) | Phosphodiester bond |
| Lipids | Fatty acids + glycerol | Esters, water-insoluble, energy storage | Ester bond |
| Vitamins | Organic compounds | Required in small amounts, not synthesized by body | — |
─── Carbohydrates ───
Monosaccharides: glucose (aldohexose), fructose (ketohexose)
Glucose: C₆H₁₂O₆, straight chain & cyclic (α/β) forms
Mutarotation: α ↔ β in aqueous solution
D/L configuration based on chiral C farthest from CHO
Disaccharides:
Sucrose (glucose + fructose): non-reducing (no free −CHO/−CO)
Maltose (glucose + glucose): reducing sugar
Lactose (glucose + galactose): reducing sugar
Polysaccharides:
Starch: amylose (linear, α-1,4) + amylopectin (branched, α-1,6)
Cellulose: β-1,4 glycosidic bonds (humans can't digest)
Glycogen: animal starch, highly branched
─── Proteins ───
Amino acids: H₂N−CHR−COOH (R = side chain)
Essential: Val, Leu, Ile, Met, Phe, Trp, Thr, Lys
Zwitterion: H₃N⁺−CHR−COO⁻ (at isoelectric point, pH = pI)
pI = (pKa₁ + pKa₂)/2 for neutral amino acids
Protein Structure:
Primary: Amino acid sequence (peptide bonds)
Secondary: α-helix (H-bonds) or β-pleated sheet (H-bonds)
Tertiary: 3D folding (H-bonds, ionic, disulfide, hydrophobic)
Quaternary: Multiple subunits (e.g., hemoglobin: 4 subunits)
Denaturation: Loss of 2°, 3°, 4° structure (not 1°)
Causes: heat, pH, heavy metals, organic solvents
─── Nucleic Acids ───
DNA: Deoxyribose + A,T,G,C (double helix — Watson & Crick)
A=T (2 H-bonds), G≡C (3 H-bonds)
Chargaff's rule: A=T, G=C
Complementary base pairing, antiparallel strands
RNA: Ribose + A,U,G,C (single stranded)
mRNA (messenger), tRNA (transfer), rRNA (ribosomal)| Type | Mechanism | Examples |
|---|---|---|
| Addition (Chain Growth) | Free radical / ionic | Polyethylene, PVC, Teflon, Polystyrene, PAN |
| Condensation (Step Growth) | Monomers react with loss of small molecule | Nylon-6,6, Bakelite, Polyester, Dacron |
| Elastomers | Weak intermolecular forces, elastic | Buna-S, Buna-N, Neoprene, Rubber |
| Fibres | Strong H-bonds, high tensile strength | Nylon-6,6, Polyester, Polyacrylonitrile |
| Thermoplastics | Linear, soften on heating | Polyethylene, PVC, Polystyrene |
| Thermosetting | Cross-linked, don't soften | Bakelite, Melamine, Urea-formaldehyde |
─── Important Polymers ───
Polyethylene: nCH₂=CH₂ → (−CH₂−CH₂−)ₙ
LDPE (low density): branched, flexible
HDPE (high density): linear, strong
PVC: nCH₂=CHCl → (−CH₂−CHCl−)ₙ (pipes, cables)
Teflon (PTFE): nCF₂=CF₂ → (−CF₂−CF₂−)ₙ (non-stick coating)
Polystyrene: nC₆H₅CH=CH₂ → (−CH(CH₃)−)ₙ (styrofoam)
Nylon-6,6: HOC(CH₂)₄COOH + H₂N(CH₂)₆NH₂
→ Condensation polymer with amide linkage
Bakelite: Phenol + Formaldehyde (condensation)
→ Cross-linked thermosetting polymer
Rubber (Natural): cis-1,4-polyisoprene
Vulcanization: heating with S (improves strength)
→ introduces S cross-links between chains
Buna-S: Butadiene + Styrene (copolymer)
Buna-N: Butadiene + Acrylonitrile
Neoprene: Chloroprene polymer
Biodegradable polymers:
PHBV (Poly β-hydroxybutyrate-co-β-hydroxyvalerate)
Nylon-2, Nylon-6| Bond Type | Formation | Properties |
|---|---|---|
| Ionic (Electrovalent) | Complete transfer of electrons (metal → non-metal) | High MP/BP, soluble in water, conducts when molten/dissolved, non-directional |
| Covalent | Sharing of electron pairs | Low MP/BP, insoluble in water, directional, may/may not conduct |
| Polar Covalent | Unequal sharing (ΔEN > 0 but < 1.7) | Dipole moment, intermediate properties |
| Coordinate (Dative) | Shared pair donated by one atom | Donor-acceptor bond, e.g., NH₄⁺, H₃O⁺ |
| Metallic | Sea of delocalized electrons | Good conductor, malleable, ductile, lustrous |
| Hydrogen Bond | H bonded to highly electronegative atom (F, O, N) | Weak, affects BP/MP/solubility/structure |
| Steric Number | Hybridization | Geometry | Bond Angle | Example |
|---|---|---|---|---|
| 2 | sp | Linear | 180° | BeCl₂, CO₂, C₂H₂ |
| 3 | sp² | Trigonal Planar | 120° | BF₃, BCl₃, CH₂=CH₂ |
| 4 | sp³ | Tetrahedral | 109.5° | CH₄, NH₃, H₂O, NH₄⁺ |
| 5 | sp³d | Trigonal Bipyramidal | 90°, 120° | PF₅, PCl₅ |
| 6 | sp³d² | Octahedral | 90° | SF₆, [Fe(CN)₆]³⁻ |
| 7 | sp³d³ | Pentagonal Bipyramidal | 72°, 90° | IF₇ |
─── VSEPR Shapes (with lone pairs) ───
Steric Number = bonded atoms + lone pairs on central atom
AX₂E: Bent/V-shaped (SO₂, SnCl₂) → < 120°
AX₃E: Trigonal Pyramidal (NH₃, PH₃) → 107°
AX₂E₂: Bent/V-shaped (H₂O, H₂S) → 104.5°
AX₄E: See-saw (SF₄) → varies
AX₃E₂: T-shaped (ClF₃) → 90°
AX₂E₃: Linear (XeF₂, I₃⁻) → 180°
AX₅E: Square Pyramidal (BrF₅) → < 90°
AX₄E₂: Square Planar (XeF₄) → 90°
─── Molecular Orbital (MO) Theory ───
For homonuclear diatomics (O₂, F₂: Z ≥ 8):
σ(1s) < σ*(1s) < σ(2s) < σ*(2s) < σ(2p_z)
< π(2p_x) = π(2p_y) < π*(2p_x) = π*(2p_y) < σ*(2p_z)
For N₂, C₂, B₂ (Z < 8):
σ(1s) < σ*(1s) < σ(2s) < σ*(2s) < π(2p_x) = π(2p_y)
< σ(2p_z) < π*(2p_x) = π*(2p_y) < σ*(2p_z)
Bond Order = ½(Nb − Na)
Nb = bonding electrons, Na = antibonding electrons
BO = 1: single bond (H₂, F₂)
BO = 2: double bond (O₂)
BO = 3: triple bond (N₂)
BO = 0: not stable (He₂)
O₂ is paramagnetic (2 unpaired e⁻ in π* orbitals)
N₂ is diamagnetic (all paired)
B₂ is paramagnetic (2 unpaired e⁻ in π orbitals)
─── Fajan's Rules (Covalent character in ionic bonds) ───
Covalent character increases when:
1. Cation is small (high charge density)
2. Anion is large (more polarizable)
3. High charge on cation
Example: LiCl > NaCl > KCl (increasing cation size
→ decreasing covalent character)
AlCl₃ is covalent (small, high charge Al³⁺)| Property | Group 1 (Alkali Metals) | Group 2 (Alkaline Earth) |
|---|---|---|
| General Config | ns¹ | ns² |
| Oxidation State | +1 | +2 |
| Atomic/Ionic Size | Large, increases down group | Smaller than G1, increases down |
| Ionization Energy | Low, decreases down | Higher than G1, decreases down |
| Reactivity | Very reactive, increases down | Less reactive than G1, increases down |
| Flame Colour | Li-crimson, Na-yellow, K-violet, Rb-red, Cs-blue | Ca-brick red, Sr-green, Ba-green/apple |
| Hydroxide Nature | Strongly basic | Less basic than G1 |
| Carbonates | Stable, soluble | Less stable, decompose on heating |
| Bicarbonates | Exist in solid state | Exist only in solution |
| Oxides | M₂O, M₂O₂, MO₂ (Li) | MO, MO₂, MO₃ (Ba) |
─── Anomalous Behavior of Li (Group 1) ───
Li is small with high charge density → differs from rest:
• Hardest alkali metal, highest MP/BP
• Least reactive (coated with Li₂O/Li₂CO₃)
• Forms Li₃N (others don't react with N₂)
• Li₂CO₃ decomposes on heating (others are stable)
• LiHCO₃ doesn't exist in solid state
• LiCl is deliquescent, soluble in organic solvents
• Li forms covalent organometallic compounds
• Li₂O₂ less stable than Li₂O
─── Anomalous Behavior of Be (Group 2) ───
• High ionization energy (small size, high IE)
• Amphoteric oxide (BeO) — others are basic
• Covalent compounds (BeCl₂ is polymeric)
• Doesn't react with water (others do)
• No peroxide or superoxide
─── Diagonal Relationship ───
Li ↔ Mg: Similar ionic radii (polarizing power)
• Both form nitride, carbonate decomposes
• Both form covalent organometallics
• Both carbonates decompose on heating
Be ↔ Al: Similar charge/size ratio
• Both form covalent halides
• Both oxides/hydroxides are amphoteric
• Both chlorides are Lewis acids| Group | Elements | General Formula | Key Feature |
|---|---|---|---|
| 13 (Boron) | B, Al, Ga, In, Tl | ns²np¹ | B is metalloid, +3 oxidation state |
| 14 (Carbon) | C, Si, Ge, Sn, Pb | ns²np² | Catenation (C, Si), allotropy |
| 15 (Nitrogen) | N, P, As, Sb, Bi | ns²np³ | N₂ gas, P₄ solid, Bi metallic |
| 16 (Oxygen) | O, S, Se, Te, Po | ns²np⁴ | O₂ gas, S₈ ring, allotropy |
| 17 (Halogens) | F, Cl, Br, I, At | ns²np⁵ | Highly reactive non-metals, −1 state |
| 18 (Noble Gases) | He, Ne, Ar, Kr, Xe, Rn | ns²np⁶ (except He) | Inert, Xe forms compounds |
─── Group 15 (Nitrogen Family) ───
Electronic config: ns²np³
Oxides of N: N₂O, NO, N₂O₃, N₂O₄, NO₂, N₂O₅
N₂O: Laughing gas, neutral
NO: Colourless, paramagnetic, neutral
NO₂: Brown gas, acidic, dimerizes to N₂O₄
Oxyacids of N:
HNO₂ (nitrous acid): weak, oxidizing
HNO₃ (nitric acid): strong, oxidizing
Cone. HNO₃ → brown fumes (NO₂)
Dilute HNO₃ + Cu → NO
Cone. HNO₃ + Cu → NO₂
Au + conc. HNO₃ + conc. HCl → Aqua Regia → AuCl₄⁻ + NO
Phosphorus: P₄ (white), red P, black P
PH₃ (phosphine): toxic, weaker base than NH₃
PCl₃, PCl₅: important reagents
H₃PO₄ (phosphoric acid): tribasic, weak
─── Group 16 (Oxygen Family) ───
O₂: Paramagnetic (2 unpaired e⁻), diatomic gas
O₃ (ozone): Resonance hybrid, bent, powerful oxidizing agent
S₈: Crown-shaped, rhombic (α) and monoclinic (β) forms
Oxyacids of S:
H₂SO₃ (sulphurous): weak, reducing agent
H₂SO₄ (sulphuric): strong, dehydrating, oxidizing
Concentrated H₂SO₄: charring (C → CO₂ + H₂O)
Contact process: SO₂ + ½O₂ → V₂O₅ → SO₃ → H₂SO₄
─── Group 17 (Halogens) ───
Reactivity: F₂ > Cl₂ > Br₂ > I₂ (decreases down)
Acid strength: HF < HCl < HBr < HI
Bond dissociation: F₂ < Cl₂ > Br₂ > I₂ (anomaly at F₂)
Oxidizing power: F₂ > Cl₂ > Br₂ > I₂
Interhalogen Compounds: AB, AB₃, AB₅, AB₇
ClF, BrF₃, IF₅, IF₇
─── Group 18 (Noble Gases) ───
He: used in balloons, cooling (low BP: 4.2 K)
Xe compounds: XeF₂, XeF₄, XeF₆, XeO₃
XeF₂: linear (sp³d), XeF₄: square planar (sp³d²)─── Transition Metals (d-Block, Group 3-12) ───
General Characteristics:
1. Variable Oxidation States: Due to comparable energy of
ns and (n-1)d orbitals
e.g., Mn: +2, +3, +4, +5, +6, +7
2. Formation of Coloured Ions: Due to d-d transitions
(partially filled d-orbitals absorb visible light)
Ti³⁺ (purple), V³⁺ (green), Cr³⁺ (violet/green),
Mn²⁺ (pink), Fe²⁺ (green), Fe³⁺ (yellow/brown),
Co²⁺ (pink), Ni²⁺ (green), Cu²⁺ (blue)
3. Magnetic Properties:
Paramagnetic: unpaired electrons (most TM ions)
Diamagnetic: all electrons paired
Magnetic moment: μ = √[n(n+2)] BM (spin-only formula)
n = number of unpaired electrons
4. Catalytic Activity: Due to variable oxidation states
and ability to form intermediate complexes
Fe: Haber process; V₂O₅: Contact process;
Ni: Hydrogenation; MnO₂: decomposition of H₂O₂
5. Formation of Complex Compounds:
Due to small size, high charge, vacant d-orbitals
e.g., [Fe(CN)₆]⁴⁻, [Cu(NH₃)₄]²⁺
6. Formation of Interstitial Compounds:
Small atoms (H, C, N, B) trapped in metal lattice
e.g., TiC, VH₀.₅, steel (C in Fe)
Harder but less malleable than parent metal
7. Alloy Formation: Mix of metals (homogeneous)
Brass (Cu+Zn), Bronze (Cu+Sn), Steel (Fe+C)
─── Lanthanide Contraction ───
Steady decrease in atomic/ionic radii of lanthanides
due to poor shielding by 4f electrons
Consequences:
• Zr/Hf, Nb/Ta, Mo/W have very similar properties
• Separation is difficult
• Lu is similar in size to Y
─── Actinides ───
5f series, all radioactive, Th and U most common
Show wider range of oxidation states than lanthanides
U: +3 to +6, Pu: +3 to +7─── Werner's Theory ───
1. Metals show two types of valence:
Primary (ionizable): satisfied by negative ions
Secondary (non-ionizable): satisfied by ligands
2. Secondary valence is directed in fixed positions → geometry
3. Coordination number = total secondary valence
Example: CoCl₃·6NH₃
Primary: 3 Cl⁻ (ionizable)
Secondary: 6 NH₃ (non-ionizable)
→ [Co(NH₃)₆]Cl₃
─── Important Terms ───
Ligand: Molecule/ion that donates electron pair(s)
Monodentate: 1 donor atom (H₂O, NH₃, Cl⁻, CN⁻)
Bidentate: 2 donor atoms (en = ethylenediamine)
Polydentate: >2 donor atoms (EDTA⁴⁻ = hexadentate)
Ambidentate: 2 donor atoms, can bond through either
(NO₂⁻: N or O; SCN⁻: S or N)
Coordination Number (CN):
CN 2: Linear, e.g., [Ag(NH₃)₂]⁺
CN 4: Tetrahedral e.g., [NiCl₄]²⁻
CN 4: Square Planar e.g., [Ni(CN)₄]²⁻, [PtCl₄]²⁻
CN 6: Octahedral e.g., [Fe(CN)₆]³⁻, [Cu(NH₃)₄(H₂O)₂]²⁺
─── IUPAC Nomenclature ───
Order: (alphabetical) ligands → metal → oxidation state
Ligands: H₂O (aqua), NH₃ (ammine), CO (carbonyl),
NO (nitrosyl), Cl⁻ (chlorido), OH⁻ (hydroxido),
CN⁻ (cyanido), en (ethylenediamine)
Prefix for multiple: di-, tri-, tetra-, penta-, hexa-
Examples:
[Co(NH₃)₆]Cl₃ → Hexaamminecobalt(III) chloride
K₃[Fe(CN)₆] → Potassium hexacyanidoferrate(III)
[Cu(NH₃)₄]SO₄ → Tetraamminecopper(II) sulphate
[PtCl₄]²⁻ → Tetrachloridoplatinate(II)
─── Isomerism in Coordination Compounds ───
Structural:
Ionization: [Co(NH₃)₅SO₄]Br vs [Co(NH₃)₅Br]SO₄
Linkage: [Co(NO₂)(NH₃)₅]²⁺ (nitro vs nitrito)
Coordination: [Cr(H₂O)₅Cl]Cl₂ vs [Cr(H₂O)₄Cl₂]Cl·2H₂O
Stereoisomerism:
Geometrical: cis/trans (square planar, octahedral)
cis-[Pt(NH₃)₂Cl₂] vs trans-[Pt(NH₃)₂Cl₂]
cis is anticancer drug (cisplatin)
Optical: non-superimposable mirror images
[Co(en)₃]³⁺ (two optical isomers)
─── Crystal Field Theory (CFT) ───
In octahedral field (CN=6):
d orbitals split into: t₂g (lower, 3 orbitals)
and eg (higher, 2 orbitals)
Crystal field splitting energy: Δo (or 10Dq)
Strong field ligands (large Δo): CN⁻ > en > NH₃ > H₂O
→ electrons pair first → low spin (fewer unpaired)
Weak field ligands (small Δo): I⁻ < Br⁻ < Cl⁻ < F⁻ < H₂O
→ electrons fill orbitals first → high spin (more unpaired)
CFSE (Crystal Field Stabilization Energy):
For dⁿ in octahedral:
CFSE = (−0.4 × n(t₂g) + 0.6 × n(eg)) × Δo
+ pairing energy × n_pairs
Color: Δo corresponds to visible light wavelength
→ complementary colour is absorbed
[Ti(H₂O)₆]³⁺: d¹, absorbs green → appears violet/purple
Magnetic properties:
μ_eff = √[n(n+2)] BM
High spin: more unpaired → more paramagnetic
Low spin: fewer unpaired → less paramagnetic| Feature | Detail |
|---|---|
| Periods | 7 horizontal rows; period number = highest principal quantum number (n) |
| Groups | 18 vertical columns; elements in same group have similar electronic configuration |
| s-block | Group 1 & 2; ns¹, ns² (except He); highly reactive metals |
| p-block | Group 13-18; ns²np¹⁻⁶; metals, non-metals, metalloids |
| d-block | Group 3-12; (n-1)d¹⁻¹⁰ns¹⁻²; transition metals |
| f-block | Lanthanides (4f) and Actinides (5f); placed separately |
| Total Elements | 118 (up to Og, Z=118) |
| Metals | ~80% of all elements; left and centre of table |
| Non-metals | ~17 elements; upper right of table |
| Metalloids | B, Si, Ge, As, Sb, Te; border between metals & non-metals |
| Property | Across Period (→) | Down Group (↓) | Reason |
|---|---|---|---|
| Atomic Radius | Decreases | Increases | ↑ nuclear charge → pulls electrons in; ↓ shells → larger |
| Ionic Radius | Cation < atom; Anion > atom | Increases | Loss of e⁻ → smaller; Gain of e⁻ → larger; more shells → larger |
| Ionization Energy (IE) | Generally increases | Generally decreases | ↑ nuclear charge → harder to remove; ↓ distance → easier |
| Electron Affinity (EA) | Generally increases (becomes more −ve) | Generally decreases | ↑ nuclear charge → more attraction for e⁻; ↓ distance → less attraction |
| Electronegativity | Increases | Decreases | ↑ nuclear charge → stronger pull; ↓ distance → weaker pull |
| Metallic Character | Decreases | Increases | Opposite of electronegativity trend |
| Non-metallic Character | Increases | Decreases | Follows electronegativity trend |
| Valency | Varies 1→4→3→2→1→0 | Usually constant | Based on valence electrons / octet completion |
─── Effective Nuclear Charge (Z_eff) ───
Z_eff = Z − S
Z = atomic number, S = shielding constant
Penetration effect: s > p > d > f
(s electrons shield less effectively than d/f)
─── Exceptions in IE Trend ───
Group 2 vs 13: Be(IE₁) > B(IE₁) because Be has
stable 2s² configuration
Group 15 vs 16: N(IE₁) > O(IE₁) because N has
stable half-filled 2p³ configuration
Group 12 vs 13: Zn(IE₁) > Ga(IE₁) because Zn has
stable 3d¹⁰4s² configuration
─── Electronegativity Scales ───
Pauling Scale: F = 4.0 (most electronegative)
O = 3.5, N = 3.0, Cl = 3.0, C = 2.5, H = 2.1
Metals: Na = 0.9, K = 0.8, Cs = 0.7
Mulliken Scale: EN = (IE₁ + EA) / 2
─── Types of Elements ───
Representative (s & p block): Show characteristic
properties, valence = group number (1-4) or 8-group(5-7)
Transition metals (d block): Variable OS, coloured ions,
magnetic properties, catalytic activity
Inner transition (f block): Lanthanides + Actinides
Noble gases: Fully filled outermost shell, inert
(except Xe, Kr, Rn form some compounds)
─── Diagonal Relationship ───
Elements diagonally adjacent show similar properties:
Li ↔ Mg, Be ↔ Al, B ↔ Si
Due to similar ionic radii and charge/size ratio─── Dynamic Equilibrium ───
Forward rate = Backward rate
Concentrations of reactants and products remain constant
(not zero — reactions still occur at molecular level)
Law of Mass Action (Guldberg & Waage):
For reaction: aA + bB ⇌ cC + dD
Equilibrium Constant (Kc):
[C]ᶜ [D]ᵈ
Kc = ─────────────
[A]ᵃ [B]ᵇ
Equilibrium Constant (Kp) — in terms of partial pressures:
(P_C)ᶜ (P_D)ᵈ
Kp = ───────────────
(P_A)ᵃ (P_B)ᵇ
Relation between Kp and Kc:
Kp = Kc × (RT)^Δn
Δn = (moles of gaseous products) − (moles of gaseous reactants)
R = 0.0821 L·atm/(mol·K), T = temperature in K
Δn = 0 → Kp = Kc
Δn > 0 → Kp > Kc
Δn < 0 → Kp < Kc
─── Kc and Reaction Quotient (Q) ───
Q = [C]ᶜ[D]ᵈ/[A]ᵃ[B]ᵇ (at any point, not necessarily equilibrium)
Q < Kc → Reaction proceeds forward (more products)
Q = Kc → At equilibrium
Q > Kc → Reaction proceeds backward (more reactants)
─── Le Chatelier's Principle ───
"If a system at equilibrium is subjected to a change,
the equilibrium shifts in the direction that counteracts
the change."
Change | Shift Direction
──────────────────────────────────────
↑ Concentration | Away from added species
↓ Concentration | Toward removed species
↑ Pressure | Toward fewer gas moles
↓ Pressure | Toward more gas moles
↑ Temperature | Endothermic direction
↓ Temperature | Exothermic direction
Catalyst | No shift (speeds up both equally)
Note: Adding inert gas at constant volume → no effect
Adding inert gas at constant pressure → shift
toward more gas moles side─── Equilibrium Calculations ───
For: aA + bB ⇌ cC + dD
Initial: a b 0 0
At eqm: a−α b−α cα dα (if started with a,b moles)
Degree of dissociation (α):
α = moles dissociated / total moles initially
For dissociation: P = P₀(1 + α) (where P₀ = initial pressure)
Vapour density:
D/d = 1 + α (D = theoretical, d = observed)
α = (D − d) / d
Equilibrium constant in terms of α:
For PCl₅ ⇌ PCl₃ + Cl₂:
Kp = (α²) × P / (1 − α²)
─── Thermodynamic Relation ───
ΔG° = −RT ln K
ΔG° = −2.303 RT log K
At equilibrium: ΔG = 0
ΔG° < 0 → K > 1 → products favoured
ΔG° > 0 → K < 1 → reactants favoured
Van't Hoff Equation (effect of T on K):
ln(K₂/K₁) = (ΔH°/R) × (1/T₁ − 1/T₂)
or log(K₂/K₁) = ΔH°/(2.303R) × (T₂−T₁)/(T₁×T₂)
For endothermic: K increases with T
For exothermic: K decreases with T| Concept | Formula | Value / Notes |
|---|---|---|
| pH | pH = −log[H⁺] | Neutral: pH = 7; Acidic: pH < 7; Basic: pH > 7 |
| pOH | pOH = −log[OH⁻] | pOH = 14 − pH (at 25°C) |
| Kw (Water) | Kw = [H⁺][OH⁻] | Kw = 10⁻¹⁴ at 25°C |
| Ka (Acid) | Ka = [H⁺][A⁻]/[HA] | Larger Ka = stronger acid |
| Kb (Base) | Kb = [B⁺][OH⁻]/[BOH] | Larger Kb = stronger base |
| Ka × Kb | Ka × Kb = Kw | For conjugate acid-base pair |
| pKa | pKa = −log Ka | Smaller pKa = stronger acid |
| pKb | pKb = −log Kb | Smaller pKb = stronger base |
| pKa + pKb | pKa + pKb = pKw = 14 | At 25°C |
─── Acid-Base Theories ───
Arrhenius: Acid gives H⁺, base gives OH⁻ (in water)
Brønsted-Lowry: Acid = proton donor, Base = proton acceptor
Lewis: Acid = e⁻ pair acceptor, Base = e⁻ pair donor
─── Relative Strength ───
Acid strength order (hydrohalic acids):
HF < HCl < HBr < HI
Reason: H-F bond is very strong (short, high EN of F)
and F⁻ is highly solvated
Oxyacid strength: HClO < HClO₂ < HClO₃ < HClO₄
More O atoms → easier release of H⁺
─── Common Ion Effect ───
Shifting equilibrium by adding a common ion
Example: CH₃COOH ⇌ CH₃COO⁻ + H⁺
Adding CH₃COONa (provides CH₃COO⁻):
→ equilibrium shifts left → [H⁺] decreases → pH increases
─── Buffer Solutions ───
Solutions that resist change in pH on adding small
amounts of acid/base
Acidic Buffer: Weak acid + salt of its conjugate base
e.g., CH₃COOH + CH₃COONa
Basic Buffer: Weak base + salt of its conjugate acid
e.g., NH₄OH + NH₄Cl
Henderson-Hasselbalch Equation:
For acidic buffer:
pH = pKa + log([salt]/[acid])
pH = pKa + log([A⁻]/[HA])
For basic buffer:
pOH = pKb + log([salt]/[base])
pOH = pKb + log([BH⁺]/[B])
Buffer capacity: Maximum when [salt] = [acid]
Effective range: pH = pKa ± 1
─── Hydrolysis ───
Salt + Water → Acid + Base
Salt of SA + WB (e.g., NH₄Cl):
pH < 7 (acidic), h = √(Kw/Kb·C)
Salt of WA + SB (e.g., CH₃COONa):
pH > 7 (basic), h = √(Kw/Ka·C)
Salt of WA + WB (e.g., CH₃COONH₄):
pH = 7 (neutral), Kh = Kw/(Ka·Kb)
Salt of SA + SB (e.g., NaCl):
pH = 7 (neutral), no hydrolysis
h = degree of hydrolysis
Kh = hydrolysis constant─── Solubility Product ───
For a saturated solution of sparingly soluble salt:
AₓBᵧ ⇌ xAᵧ⁺ + yBₓ⁻
Ksp = [Aᵧ⁺]ˣ × [Bₓ⁻]ʸ
Relation to solubility (s):
Ksp = (xs)ˣ × (ys)ʸ
Examples:
AgCl → Ksp = [Ag⁺][Cl⁻] = s²
PbCl₂ → Ksp = [Pb²⁺][Cl⁻]² = s(2s)² = 4s³
Ag₂CrO₄ → Ksp = [Ag⁺]²[CrO₄²⁻] = (2s)²(s) = 4s³
─── Precipitation ───
Ionic Product (IP) = [Aᵧ⁺]ˣ[Bₓ⁻]ʸ (at given conditions)
IP < Ksp → Unsaturated → No ppt, more can dissolve
IP = Ksp → Saturated → Just saturated
IP > Ksp → Supersaturated → Precipitation occurs
Selective Precipitation:
Precipitate least soluble salt first
e.g., AgCl (Ksp = 1.8×10⁻¹⁰) precipitates before
Ag₂CrO₄ (Ksp = 1.1×10⁻¹²) because [Ag⁺] needed
for AgCl is lower (compare: Ksp order is misleading,
must compare required [Ag⁺])
─── Common Ion Effect on Ksp ───
Adding common ion → decreases solubility
AgCl in NaCl solution: [Ag⁺] × [Cl⁻] = Ksp
Since [Cl⁻] increases from NaCl, [Ag⁺] must decrease
→ AgCl dissolves less| Type | Reaction | Energy Change | Example |
|---|---|---|---|
| Galvanic (Voltaic) | Spontaneous redox → electricity | Chemical → Electrical | Daniell cell (Zn-Cu) |
| Electrolytic | Electricity drives non-spontaneous | Electrical → Chemical | Electroplating, water electrolysis |
─── Standard Electrode Potential ───
E°(cell) = E°(cathode) − E°(anode)
(cathode = reduction, anode = oxidation)
For Daniell Cell: Zn(s)|Zn²⁺||Cu²⁺|Cu(s)
Anode (ox): Zn → Zn²⁺ + 2e⁻ E° = +0.76 V (as oxidation)
Cathode (red): Cu²⁺ + 2e⁻ → Cu E° = +0.34 V
E°cell = 0.34 − (−0.76) = 1.10 V
─── Nernst Equation ───
For: aA + bB → cC + dD (with n electrons transferred)
RT [C]ᶜ [D]ᵈ
E = E° − ─── ln ─────────────
nF [A]ᵃ [B]ᵇ
At 25°C:
0.059 [C]ᶜ [D]ᵈ
E = E° − ───── log ─────────────
n [A]ᵃ [B]ᵇ
For single electrode (Mⁿ⁺ + ne⁻ → M):
0.059 1
E = E° − ───── log ───
n [Mⁿ⁺]
When E(cell) = 0 → equilibrium
E° = (0.059/n) log K → log K = nE°/0.059
─── Conductivity ───
Resistance: R = ρ × l/A = (1/κ) × l/A
Conductance: G = 1/R = κ × A/l
Specific conductance (κ): Conductance of 1 cm³ solution
Unit: S cm⁻¹ (Siemens per cm)
Molar conductivity (Λm):
Λm = κ × 1000 / C (C = molarity)
Unit: S cm² mol⁻¹
Kohlrausch's Law:
Λ°m = ν₊λ°₊ + ν₋λ°₋
Λ°m = limiting molar conductivity
λ° = limiting ionic conductivity
ν = no. of ions per formula
Degree of dissociation (α):
α = Λm / Λ°m = Λm / (ν₊λ°₊ + ν₋λ°₋)
Ka (acid dissociation constant):
Ka = Cα² / (1 − α) = C × (Λm/Λ°m)² / (1 − Λm/Λ°m)─── First Law ───
Mass deposited/liberated is proportional to charge passed
w = Z × Q = Z × I × t
Z = electrochemical equivalent (mass per coulomb)
I = current (amperes), t = time (seconds)
─── Second Law ───
When same quantity of electricity is passed through
different electrolytes, masses deposited are proportional
to their equivalent weights.
w₁/w₂ = E₁/E₂
─── Faraday's Constant ───
F = 96485 C mol⁻¹ ≈ 96500 C mol⁻¹
Charge for 1 mole of electrons: 1F = 96500 C
Relations:
1 mole e⁻ = 1F = 96500 C
→ deposits 1 equivalent weight of substance
For: Mⁿ⁺ + ne⁻ → M
Mass deposited = (M × I × t) / (n × F)
M = molar mass, n = valency (no. of electrons)
Practical Applications:
• Electroplating: Coating with a thin layer of metal
• Electrorefining: Purifying metals (e.g., Cu)
• Electrolysis of water: 2H₂O → 2H₂ + O₂
• Production of Na, Al, Mg by electrolysis| Battery / Cell | Anode | Cathode | EMF (V) | Key Features |
|---|---|---|---|---|
| Dry Cell (Leclanche) | Zn (container) | MnO₂ + C | ~1.5 | Not rechargeable, NH₄Cl paste |
| Lead-Acid Battery | Pb + H₂SO₄ | PbO₂ + H₂SO₄ | ~2.0 | Rechargeable, car batteries |
| Ni-Cd | Cd | NiO(OH) | ~1.2 | Rechargeable, memory effect |
| Li-ion | LiCoO₂ (cathode) | Graphite (anode) | ~3.7 | Lightweight, high energy density |
| H₂-O₂ Fuel Cell | H₂ → 2H⁺ + 2e⁻ | O₂ + 4H⁺ + 4e⁻ → 2H₂O | 1.23 | Clean energy, used in spacecraft |
─── Corrosion ───
Electrochemical process: metal is oxidized by air/water
Rusting of Iron:
Anode: Fe → Fe²⁺ + 2e⁻
Cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂O
Further: Fe²⁺ + 2OH⁻ → Fe(OH)₂ → Fe₂O₃·xH₂O (rust)
Prevention Methods:
1. Barrier: painting, oiling, greasing, plastic coating
2. Galvanization: coating iron with zinc (Zn is more reactive)
3. Sacrificial protection: Mg/Zn blocks attached (more reactive)
4. Alloying: stainless steel (Fe + Cr + Ni)
5. Cathodic protection: connecting to more active metal─── 1. WURTZ REACTION ───
2R-X + 2Na → R-R + 2NaX
(Alkyl halides → higher alkanes)
Note: Only symmetric alkanes; works with 1° halides
Example: 2CH₃Br + 2Na → CH₃-CH₃ + 2NaBr
─── 2. FRIEDEL-CRAFTS ALKYLATION ───
C₆H₆ + R-Cl → AlCl₃ → C₆H₅-R + HCl
Electrophile: R⁺ (carbocation)
Limitations: No reaction with −NO₂, −CN, −CHO groups
(deactivated rings); rearrangement of R⁺ possible
─── 3. FRIEDEL-CRAFTS ACYLATION ───
C₆H₆ + RCOCl → AlCl₃ → C₆H₅-COR + HCl
Electrophile: RCO⁺ (acylium ion)
Advantage: No rearrangement (acylium ion is resonance
stabilized); gives mono-acylated product
─── 4. CANNIZZARO REACTION ───
2RCHO + NaOH(conc) → RCH₂OH + RCOONa
(Aldehydes without α-hydrogen undergo disproportionation)
Examples: HCHO, C₆H₅CHO, (CH₃)₃CCHO
Cross Cannizzaro: HCHO + C₆H₅CHO → PhCH₂OH + HCOONa
─── 5. ALDOL CONDENSATION ───
2CH₃CHO + NaOH(dil) → CH₃CH(OH)CH₂CHO (aldol)
→ heat → CH₃CH=CHCHO (crotonaldehyde) + H₂O
Conditions: Aldehyde/ketone with at least one α-H
Cross aldol: Between two different carbonyl compounds
─── 6. HOFMANN BROMAMIDE DEGRADATION ───
RCONH₂ + Br₂ + 4NaOH → RNH₂ + Na₂CO₃ + 2NaBr + 2H₂O
(Primary amide → primary amine, loses 1 C atom)
Mechanism: N-bromoamide → isocyanate → amine
─── 7. SANDEYER REACTION ───
ArN₂⁺Cl⁻ + CuX → Ar-X + N₂ + CuCl
(Diazonium salt → aryl halide)
CuCl → ArCl, CuBr → ArBr, CuCN → ArCN
─── 8. GATTERMANN REACTION ───
Similar to Sandmeyer but uses Cu powder instead of CuX
ArN₂⁺Cl⁻ + Cu → ArCl (chlorobenzene)
Used when Cu halides are not available
─── 9. GATTERMANN-KOCH REACTION ───
C₆H₆ + CO + HCl → AlCl₃/CuCl → C₆H₅CHO (benzaldehyde)
(Formylation of benzene)
Electrophile: HCO⁺ (formyl cation)─── 10. KOLBE ELECTROLYSIS ───
2RCOONa → electrolysis (aq) → R-R + 2CO₂ + H₂ + Na₂
(Sodium salt of carboxylic acid → alkane)
Example: 2CH₃COONa → CH₃CH₃ + 2CO₂ + H₂ + Na₂
─── 11. ROSENmund REDUCTION ───
RCOCl + H₂ → Pd/BaSO₄ (poisoned) → RCHO + HCl
(Acid chloride → aldehyde, stops at aldehyde stage)
BaSO₄ poisons Pd catalyst (Lindlar-type catalyst concept)
─── 12. CLEMMENSEN REDUCTION ───
RCOCH₃ + Zn(Hg)/HCl → RCH₂CH₃
(Carbonyl compound → methylene group)
For reducing C=O to CH₂ under acidic conditions
─── 13. WOLFF-KISHNER REDUCTION ───
RCOCH₃ + NH₂NH₂ + KOH → heat → RCH₂CH₃
(Carbonyl → methylene group, basic conditions)
Same result as Clemmensen but under basic conditions
─── 14. KOECH ALCOHOL SYNTHESIS ───
Grignard + formaldehyde → 1° alcohol
RMgX + HCHO → RCH₂OH + MgX(OH)
─── 15. REIMER-TIEMANN REACTION ───
C₆H₅OH + CHCl₃ + NaOH → o/p-OH-C₆H₄-CHO (salicylaldehyde)
(Phenol → ortho-hydroxy benzaldehyde)
Ortho product predominates
─── 16. HELL-VOLHARD-ZELINSKY (HVZ) ───
RCH₂COOH + Br₂ → PBr₃ → RCH(Br)COOH + HBr
(α-bromination of carboxylic acids)
PBr₃ first converts COOH to acyl bromide (more reactive)
─── 17. WILLIAMSON ETHER SYNTHESIS ───
R-ONa + R'-X → R-O-R' + NaX
(Alkoxide + alkyl halide → ether)
Use 1° alkyl halide (avoids elimination competition)
─── 18. GABRIEL PHTHALIMIDE SYNTHESIS ───
Phthalimide + KOH → K-phthalimide → R-X →
N-alkylphthalimide → hydrazine → RNH₂
(Produces pure 1° amine, no 2°/3° contamination)
─── 19. BALZ-SCHIEMANN REACTION ───
ArN₂⁺BF₄⁻ → heat → ArF + N₂ + BF₃
(Diazonium fluoroborate → fluorobenzene)
─── 20. BECKMANN REARRANGEMENT ───
RRC=NOH (oxime) → acid → RCONHR'
(Ketoxime → amide via migration of anti group)
Example: Cyclohexanone oxime → ε-caprolactam
(precursor to nylon-6)─── 21. FRIES REARRANGEMENT ───
Phenyl ester + AlCl₃ → ortho/para-hydroxy ketone
C₆H₅OCOCH₃ → AlCl₃ → o/p-HO-C₆H₄-COCH₃
─── 22. CARBYLAMINE TEST ───
RNH₂ + CHCl₃ + 3KOH → RNC + 3KCl + 3H₂O
(1°/2° amine → foul-smelling isocyanide)
─── 23. IODOFORM TEST ───
CH₃CHO/CH₃COR/CH₃CH(OH)R + I₂ + NaOH → CHI₃ (yellow)
+ RCOONa
(Positive for methyl ketones, acetaldehyde, ethanol)
─── 24. TOLLENS TEST ───
RCHO + 2[Ag(NH₃)₂]⁺ + 2OH⁻ → RCOO⁻ + 2Ag + 4NH₃ + H₂O
(Silver mirror — aldehyde detection)
─── 25. FEHLING'S TEST ───
RCHO + 2Cu²⁺(Fehling's) + 5OH⁻ → RCOO⁻ + Cu₂O↓ + 3H₂O
(Red precipitate of Cu₂O — aldehydes only)
─── 26. BAeyer-Villiger Oxidation ───
Ketone + peracid → ester
RCOCH₃ + mCPBA → RCOOCH₃ (migration preference: 3° > 2° > 1° > Me)
─── 27. AZO COUPLING ───
Diazonium salt + phenol/amine → azo compound (coloured)
ArN₂⁺Cl⁻ + C₆H₅OH → Ar-N=N-C₆H₄-OH (azo dye)
─── 28. STEPHEN REDUCTION ───
RCN + SnCl₂/HCl → RCH=NH·HCl → H₂O → RCHO
(Nitrile → aldehyde)| Constant | Symbol | Value | Unit |
|---|---|---|---|
| Avogadro's Number | Nₐ | 6.022 × 10²³ | mol⁻¹ |
| Gas Constant | R | 8.314 J/(mol·K) | J mol⁻¹ K⁻¹ |
| Gas Constant | R | 0.0821 L·atm/(mol·K) | L atm mol⁻¹ K⁻¹ |
| Faraday Constant | F | 96500 | C mol⁻¹ |
| Planck's Constant | h | 6.626 × 10⁻³⁴ | J·s |
| Speed of Light | c | 3 × 10⁸ | m/s |
| Electron Charge | e | 1.602 × 10⁻¹⁹ | C |
| Electron Mass | mₑ | 9.109 × 10⁻³¹ | kg |
| Boltzmann Constant | k_B | 1.381 × 10⁻²³ | J/K |
| Bohr Radius | a₀ | 0.529 × 10⁻¹⁰ | m (or 0.529 Å) |
| Rydberg Constant | R_H | 1.097 × 10⁷ | m⁻¹ |
| Atomic Mass Unit | amu | 1.661 × 10⁻²⁷ | kg |
| Standard Pressure | — | 1 atm = 1.013 × 10⁵ Pa | Pa |
| Molar Volume (STP) | — | 22.4 | L/mol |
| Water Ionic Product | Kw | 10⁻¹⁴ (at 25°C) | mol² L⁻² |
| Formula | Expression | Notes |
|---|---|---|
| Energy of orbit (H) | Eₙ = −13.6/n² eV | For hydrogen atom |
| Radius of orbit (H) | rₙ = 0.529 × n²/Z Å | Z = atomic number |
| Velocity in orbit | vₙ = 2.188 × 10⁶ × Z/n m/s | Decreases with n |
| De Broglie wavelength | λ = h/mv | Wave-particle duality |
| Heisenberg uncertainty | Δx · Δp ≥ h/4π | Cannot know both precisely |
| Rydberg equation | 1/λ = R_H(1/n₁² − 1/n₂²) | Spectral lines |
| Photoelectric effect | KE = hν − φ | φ = work function |
| Maximum electrons/orbital | 2n² | Per shell |
| Formula | Expression | Notes |
|---|---|---|
| First Law | ΔU = q + w | w = −PΔV (expansion work) |
| Enthalpy | H = U + PV | ΔH = ΔU + Δn_g RT |
| Work (isothermal) | w = −nRT ln(V₂/V₁) | Ideal gas, reversible |
| Gibbs Energy | ΔG = ΔH − TΔS | Spontaneity criterion |
| ΔG° & K relation | ΔG° = −RT ln K | K = equilibrium constant |
| Carnot efficiency | η = (T_H − T_C)/T_H | Maximum theoretical efficiency |
| CP − CV | CP − CV = R | For ideal gas (1 mol) |
| Hess's Law | ΔH_total = ΣΔH_steps | State function property |
| Formula | Expression | Notes |
|---|---|---|
| Kc | [Products]^coeff / [Reactants]^coeff | Concentration equilibrium const |
| Kp | (PP)^coeff / (PR)^coeff | Pressure equilibrium const |
| Kp & Kc relation | Kp = Kc(RT)^Δn | Δn = gas products − gas reactants |
| pH | pH = −log[H⁺] | 0–14 scale |
| pOH | pOH = −log[OH⁻] | pOH = 14 − pH at 25°C |
| Ka | Ka = [H⁺][A⁻]/[HA] | Acid dissociation constant |
| Kb | Kb = [BH⁺][OH⁻]/[B] | Base dissociation constant |
| Ka × Kb | Ka × Kb = Kw | Conjugate pair relation |
| Henderson eq (acid) | pH = pKa + log([A⁻]/[HA]) | Buffer pH calculation |
| Henderson eq (base) | pOH = pKb + log([BH⁺]/[B]) | Buffer pOH calculation |
| Ksp | [cation]^m × [anion]^n | Solubility product |
| ΔG° & K | ΔG° = −2.303 RT log K | Thermodynamic link |
| Van't Hoff | log(K₂/K₁) = ΔH°(T₂−T₁)/(2.303R T₁T₂) | K vs temperature |
| Formula | Expression | Notes |
|---|---|---|
| Cell potential | E°cell = E°cathode − E°anode | Reduction potentials |
| Nernst (25°C) | E = E° − (0.059/n) log Q | Q = reaction quotient |
| Resistance | R = ρl/A | ρ = resistivity |
| Conductance | G = 1/R | Unit: Siemens (S) |
| Conductivity | κ = 1/ρ | Specific conductance, S cm⁻¹ |
| Molar conductivity | Λm = κ × 1000/C | C = molarity |
| Kohlrausch's Law | Λ°m = ν₊λ°₊ + ν₋λ°₋ | Limiting molar conductivity |
| Degree of dissociation | α = Λm / Λ°m | Weak electrolytes |
| Faraday's 1st Law | w = ZIt = MIt/(nF) | Mass deposited |
| Faraday's 2nd Law | w₁/w₂ = E₁/E₂ | Equivalent weight ratio |
| EMF & equilibrium | E° = (0.059/n) log K | At equilibrium E = 0 |
| Formula | Expression | Notes |
|---|---|---|
| Molarity | M = n/V(L) | Changes with temperature |
| Molality | m = n/W(kg solvent) | Independent of temperature |
| Mole fraction | x₁ = n₁/(n₁+n₂) | Sum of all = 1 |
| Raoult's Law | P = x₁P₁° + x₂P²° | Volatile liquid mixture |
| Rel. lowering VP | ΔP/P° = x₂ | Non-volatile solute |
| Boiling point elevation | ΔTb = Kb × m | Kb = molal BP constant |
| Freezing point depression | ΔTf = Kf × m | Kf = molal FP constant |
| Osmotic pressure | π = iCRT | Van't Hoff: i × C × R × T |
| Van't Hoff factor | i = 1 + (n−1)α | α = degree of dissociation |
| Rate law (general) | r = k[A]ⁿ[B]ᵐ | n+m = order |
| 1st order integrated | k = (2.303/t) log([A]₀/[A]) | Plot: log[A] vs t |
| Half-life (1st order) | t½ = 0.693/k | Independent of [A]₀ |
| Half-life (zero order) | t½ = [A]₀/(2k) | Depends on [A]₀ |
| Half-life (2nd order) | t½ = 1/(k[A]₀) | Inversely proportional |
| Arrhenius equation | k = Ae^(−Ea/RT) | Rate vs temperature |
| Arrhenius (log form) | log(k₂/k₁) = Ea/(2.303R) × (1/T₁−1/T₂) | Two-temperature form |
| Formula | Expression | Notes |
|---|---|---|
| Ideal gas | PV = nRT | R = 0.0821 or 8.314 |
| Van der Waals | (P + n²a/V²)(V−nb) = nRT | Real gas correction |
| Boyle's Law | P₁V₁ = P₂V₂ | T = constant |
| Charles's Law | V₁/T₁ = V₂/T₂ | P = constant |
| Gay-Lussac's | P₁/T₁ = P₂/T₂ | V = constant |
| Combined gas | P₁V₁/T₁ = P₂V₂/T₂ | n = constant |
| RMS speed | u_rms = √(3RT/M) | M = molar mass (kg) |
| Average speed | u_avg = √(8RT/πM) | Mean molecular speed |
| Most probable speed | u_mp = √(2RT/M) | Peak of Maxwell dist. |
| Kinetic energy | KE = 3/2 kT per molecule | KE = 3/2 RT per mole |
| Graham's Law | r₁/r₂ = √(M₂/M₁) | Rate of diffusion/effusion |
| Compound | Formula | Molar Mass (g/mol) | Notes |
|---|---|---|---|
| Water | H₂O | 18 | Universal solvent |
| Sodium Hydroxide | NaOH | 40 | Strong base |
| Hydrochloric Acid | HCl | 36.5 | Strong acid, gastric acid |
| Sulphuric Acid | H₂SO₄ | 98 | King of chemicals |
| Nitric Acid | HNO₃ | 63 | Strong oxidizing agent |
| Acetic Acid | CH₃COOH | 60 | Vinegar, weak acid |
| Ammonia | NH₃ | 17 | Weak base, Lewis base |
| Calcium Carbonate | CaCO₃ | 100 | Limestone, marble |
| Sodium Chloride | NaCl | 58.5 | Table salt |
| Potassium Permanganate | KMnO₄ | 158 | Oxidizing agent (violet) |
| Potassium Dichromate | K₂Cr₂O₇ | 294 | Oxidizing agent (orange) |
| Hydrogen Peroxide | H₂O₂ | 34 | Bleaching, disinfectant |
| Ethanol | C₂H₅OH | 46 | Common alcohol |
| Methane | CH₄ | 16 | Natural gas, greenhouse |
| Glucose | C₆H₁₂O₆ | 180 | Blood sugar |
| Urea | CO(NH₂)₂ | 60 | Excretory product |
| Benzene | C₆H₆ | 78 | Aromatic hydrocarbon |
| Aspirin | C₉H₈O₄ | 180 | Pain reliever (acetylsalicylic acid) |
| Oleum | H₂S₂O₇ | 178 | H₂SO₄ + SO₃ |
| Baking Soda | NaHCO₃ | 84 | Sodium bicarbonate |
| Test | Reagent | Observation | Detected Ion |
|---|---|---|---|
| Flame test | Bunsen burner | Golden yellow flame | Na⁺ |
| Flame test | Bunsen burner | Violet flame (through cobalt glass) | K⁺ |
| Flame test | Bunsen burner | Brick red flame | Ca²⁺ |
| Flame test | Bunsen burner | Apple green flame | Ba²⁺ |
| Flame test | Bunsen burner | Crimson red flame | Li⁺ |
| Brown Ring Test | FeSO₄ + conc. H₂SO₄ | Brown ring at junction | NO₃⁻ |
| Acetate Test | FeCl₃ | Blood red colour | CH₃COO⁻ |
| Sulphate Test | BaCl₂ + HCl | White ppt (BaSO₄) | SO₄²⁻ |
| Chloride Test | AgNO₃ + HNO₃ | White ppt (AgCl) | Cl⁻ |
| Bromide Test | AgNO₃ + HNO₃ | Pale yellow ppt (AgBr) | Br⁻ |
| Iodide Test | AgNO₃ + HNO₃ | Yellow ppt (AgI) | I⁻ |
| Carbonate Test | Dilute acid | Effervescence (CO₂) | CO₃²⁻ |
| Sulphide Test | Pb(CH₃COO)₂ | Black ppt (PbS) | S²⁻ |
| Phosphate Test | Ammonium molybdate | Yellow ppt | PO₄³⁻ |
─── Unit Conversions for Chemistry ───
1 atm = 760 mmHg = 1.013 × 10⁵ Pa = 1.013 bar
1 L = 1000 mL = 10⁻³ m³
1 calorie = 4.184 J
1 eV = 1.602 × 10⁻¹⁹ J
1 Å (angstrom) = 10⁻¹⁰ m
1 nm = 10⁻⁹ m = 10 Å
1 pm = 10⁻¹² m
1 bar = 10⁵ Pa
1 torr = 1 mmHg
1 M = 1 mol/L = 1000 mol/m³
─── Standard Conditions ───
STP (Standard Temperature & Pressure): 0°C (273 K), 1 atm
1 mole gas = 22.4 L at STP
SATP (Standard Ambient): 25°C (298 K), 1 bar
NTP (Normal): 20°C (293 K), 1 atm
─── Common Acid-Base Indicators ───
Phenolphthalein: pH 8.2 (colourless) → 10.0 (pink)
Methyl Orange: pH 3.1 (red) → 4.4 (yellow)
Methyl Red: pH 4.4 (red) → 6.2 (yellow)
Litmus: pH < 5 (red), pH > 8 (blue)
Bromothymol Blue: pH 6.0 (yellow) → 7.6 (blue)
─── Important pKa/pKb Values (25°C) ───
HCl: pKa ≈ −7 (very strong)
H₂SO₄: pKa₁ ≈ −3, pKa₂ ≈ 1.99
H₃PO₄: pKa₁ = 2.12, pKa₂ = 7.21, pKa₃ = 12.68
HF: pKa = 3.17 (weak acid)
CH₃COOH: pKa = 4.76
H₂CO₃: pKa₁ = 6.35, pKa₂ = 10.33
HCN: pKa = 9.31 (very weak)
NH₄⁺: pKa = 9.25 (acidic form of NH₃)
H₂O: pKa = 15.7 (as acid)