Class 11 Chemistry ⚗️

(a) Atom: Smallest particle of an element that takes part in a chemical reaction. It may or may not exist independently.

(b) Molecule: Smallest particle of a substance that has independent existence and shows all properties of that substance.

(c) Mole: Amount of substance containing 6.022 × 10²³ (Avogadro's number) particles. Molar mass in grams = 1 mole of substance.

Moles of H₂O = 18/18 = 1 mole
Number of molecules = 1 × 6.022 × 10²³ = 6.022 × 10²³ molecules

Molar mass of CO₂ = 12 + 2×16 = 44 g/mol
% of C = (12/44) × 100 = 27.27%

Unbalanced: Fe + O₂ → Fe₂O₃
Balanced: 4Fe + 3O₂ → 2Fe₂O₃
Check: Fe: 4=4 ✓, O: 6=6 ✓

The limiting reagent is the reactant that is completely consumed first in a chemical reaction, limiting the amount of product formed.
Example: N₂ + 3H₂ → 2NH₃. If 1 mol N₂ and 2 mol H₂ are taken, H₂ is limiting (needs 3 mol but only 2 available). Amount of NH₃ = 2/3 × 2 = 1.33 mol.

Moles of NaOH = 4/40 = 0.1 mol
Volume = 500 mL = 0.5 L
Molarity = 0.1/0.5 = 0.2 M

Qualitative analysis: determines WHAT substances are present in a sample (identity of components).

Quantitative analysis: determines HOW MUCH of each substance is present (amount/concentration).

Example: Testing water for lead (qualitative). Finding how many mg/L of lead (quantitative).

Flame test: a metal compound is held in a flame on a platinum wire. The colour produced identifies the metal ion.

Na (Sodium) → Golden yellow
K (Potassium) → Violet/Lilac
Ca (Calcium) → Brick red
Cu (Copper) → Blue-green

Titration: a technique where a solution of known concentration (standard solution) is slowly added to a solution of unknown concentration until the reaction is complete.

Equivalence point: the point at which the amount of titrant added is exactly stoichiometrically equivalent to the amount of analyte in the sample. Indicated by a colour change of indicator.

Chromatography is a technique used to separate components of a mixture based on their different rates of movement through a stationary phase when carried by a mobile phase.

Principle: Different components have different affinities for the stationary and mobile phases, causing them to travel at different speeds and separate.

Paper chromatography uses filter paper as stationary phase and a solvent as mobile phase.

Process: 1) A spot of mixture is placed on paper. 2) Paper is dipped in solvent. 3) Solvent moves up by capillary action. 4) Components travel at different speeds. 5) Separated components appear as different spots.

Used to separate: ink colours, amino acids, plant pigments.

Rf (Retardation factor) = Distance travelled by component / Distance travelled by solvent front
Rf value is always between 0 and 1. It is characteristic of a substance under fixed conditions and used to identify components.

1. Electrons revolve around nucleus in fixed circular orbits called stationary states without radiating energy.
2. Only those orbits are allowed in which angular momentum = nh/2π (n = 1,2,3...)
3. Energy is emitted or absorbed only when electron jumps from one orbit to another.
ΔE = E₂ − E₁ = hν

(a) Na: 1s² 2s² 2p⁶ 3s¹ → 2,8,1
(b) Cl: 1s² 2s² 2p⁶ 3s² 3p⁵ → 2,8,7
(c) Ca: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² → 2,8,8,2

Quantum numbers describe the state and position of an electron in an atom.

1. Principal quantum number (n): energy level/shell. n = 1,2,3...
2. Azimuthal quantum number (l): subshell shape. l = 0 to n−1
3. Magnetic quantum number (m): orbital orientation. m = −l to +l
4. Spin quantum number (s): electron spin. s = +½ or −½

Pauli's Exclusion Principle: No two electrons in an atom can have the same set of all four quantum numbers. Maximum 2 electrons per orbital with opposite spins.

Hund's Rule: Electrons occupy all available orbitals of equal energy (degenerate orbitals) singly before pairing. All singly occupied orbitals have parallel spins.

Ionic bond: formed by complete transfer of electrons from metal to non-metal, creating oppositely charged ions which attract each other.

Formation of NaCl: Na (2,8,1) loses 1 electron → Na⁺ (2,8). Cl (2,8,7) gains 1 electron → Cl⁻ (2,8,8). Na⁺ and Cl⁻ attract each other forming ionic bond.

Covalent bond: formed by sharing of electrons between two non-metal atoms, each contributing one or more electrons to the shared pair.

Examples:
H₂: H−H (single bond, 1 shared pair)
O₂: O=O (double bond, 2 shared pairs)
N₂: N≡N (triple bond, 3 shared pairs)

VSEPR (Valence Shell Electron Pair Repulsion): electron pairs around central atom arrange themselves to minimize repulsion.

H₂O: 2 bond pairs + 2 lone pairs → Bent/V-shape, bond angle ≈ 104.5°
NH₃: 3 bond pairs + 1 lone pair → Trigonal pyramidal, bond angle ≈ 107°

Hydrogen bond: weak electrostatic attraction between partially positive H atom of one molecule and electronegative atom (F, O, N) of another molecule.

Water has high boiling point (100°C) because extensive hydrogen bonding between water molecules requires more energy to break. Without H-bonds, water would boil at −80°C.

Boyle's Law: At constant temperature, pressure of a gas is inversely proportional to its volume. PV = constant → P₁V₁ = P₂V₂

Charles' Law: At constant pressure, volume of a gas is directly proportional to its absolute temperature. V/T = constant → V₁/T₁ = V₂/T₂

Ideal gas equation: PV = nRT
P = pressure, V = volume, n = moles, T = temperature in Kelvin
R = Universal gas constant = 8.314 J mol⁻¹ K⁻¹

Using P₁V₁/T₁ = P₂V₂/T₂:
V₂ = P₁V₁T₂/T₁P₂ = (1×2×600)/(300×2) = 1200/600 = 2 L

Surface tension is the property of liquid surface to contract and behave like a stretched elastic membrane due to cohesive forces between molecules at the surface.
As temperature increases, surface tension decreases because kinetic energy of molecules increases and cohesive forces weaken.

Adsorption: accumulation of molecules on the surface of a substance (adsorbent). It is a surface phenomenon.

Absorption: uniform distribution of molecules throughout the bulk of a substance.

Key difference: Adsorption is surface phenomenon, absorption is bulk phenomenon. Example: silica gel adsorbs water; sponge absorbs water.

A colloid is a heterogeneous mixture where particle size is between 1nm and 1000nm. Colloidal particles are not visible to naked eye but scatter light (Tyndall effect).

Examples:
1. Milk (liquid in liquid — emulsion)
2. Fog (liquid in gas — aerosol)
3. Foam (gas in liquid)
4. Butter (liquid in solid — gel)
5. Smoke (solid in gas — aerosol)

Tyndall effect: scattering of light by colloidal particles when a beam of light passes through a colloidal solution, making the path of light visible.

Examples:
1. Sunbeam visible in smoky room
2. Headlights visible in fog
3. Blue colour of sky (Rayleigh scattering)