Quick Revision

1 & 2. Mixtures

• Definition: A physical combination of 2+ substances, each keeping its identity.
• Types:
  • Homogeneous → Uniform throughout, cannot see parts. (e.g., air, saltwater, brass).
  • Heterogeneous → Non-uniform, visible parts. (e.g., sand + water, granite, pizza).
• Separation methods: Filtration, distillation, chromatography.

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3. Properties of Elements

• Physical Properties → observed without changing identity.
  • Ex: Sodium (Na) is silvery-white, soft, low MP (97.8°C), good conductor.
• Chemical Properties → how substances react.
  • Na + H₂O → NaOH + H₂ (violent).
  • Na + Cl₂ → NaCl (salt).

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4. Isotopes & Atomic Mass

• Isotopes = same element, same protons, different neutrons.
• Formula: Avg. Atomic Mass = Σ (Mass × Abundance)
• Example (Chlorine):
  • Cl-35: 34.97 amu (75.76%)
  • Cl-37: 36.97 amu (24.24%)
  • Avg. Mass = 35.46 amu

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5. % Composition

• Formula: % Mass = (Mass of element in compound ÷ Molar mass) × 100
• Example (H₂O):
  • Molar mass = 18.02 g/mol (H = 2.02, O = 16.00).
  • %H = (2.02 ÷ 18.02) × 100 = 11.2%.
  • Means in 100 g water → 11.2 g H.

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6. Naming Compounds

1. Ionic (Metal + Nonmetal)
   • Metal name + nonmetal with “-ide”.
   • Ex: NaCl → Sodium chloride, MgF₂ → Magnesium fluoride.
2. Covalent (Nonmetals only)
   • Use prefixes (mono-, di-, tri-...).
   • Ex: N₂O₅ → Dinitrogen pentoxide, CO₂ → Carbon dioxide.

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7, 9, 10. Equilibrium

• Dynamic equilibrium = forward rate = reverse rate.
• Equilibrium constants:
  • Kc (concentration) = [products]/[reactants].
  • Kp (pressure). Relation: Kp = Kc(RT)^Δn.
• Le Châtelier’s Principle:
  • Add reactant → shifts right.
  • Add product → shifts left.
  • ↑ Pressure → shifts to fewer gas moles.
  • ↑ Temperature → shifts endothermic side.

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8. Empirical & Molecular Formulas

• Steps (Empirical):
  1. Assume 100 g sample → % = grams.
  2. Convert grams → moles.
  3. Divide by smallest mole value.
  4. Multiply to whole numbers.
• Molecular Formula: (Molar Mass ÷ EF Mass) × Empirical Formula.

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11. Acid-Base Theories

• Arrhenius: Acid = H⁺ donor, Base = OH⁻ donor.
• Brønsted-Lowry: Acid = proton donor, Base = proton acceptor.
• Lewis: Acid = e⁻ pair acceptor, Base = e⁻ pair donor.

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12. pH Calculations

• Formulas:
  • pH = -log[H⁺]
  • pOH = -log[OH⁻]
  • pH + pOH = 14
• Example: [OH⁻] = 1.0 × 10⁻³ M → pOH = 3 → pH = 11.

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13. Chapter 1 Review

Valence Electrons

• Definition: Electrons in the outermost shell used in bonding.
• How to determine:
  • Look at Group Number (1–18).
  • Skip transition metals (Groups 3–12).
  • Ex: Group 1 → 1 valence e⁻, Group 17 → 7 valence e⁻.

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Cations & Anions

• Cation (positive ion) → formed when an atom loses electrons.
• Anion (negative ion) → formed when an atom gains electrons.
• Example: Na → Na⁺, Cl → Cl⁻.

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Quantum Mechanical Model

• Electrons show dual nature (particle + wave).
• Movement is described by probability, not fixed orbits.
• This gave rise to Quantum Mechanical Model of the Atom.

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Quantum Numbers

• A set of 4 numbers describing position & energy of electrons:
  1. n = Principal (energy level).
  2. l = Azimuthal (subshell shape: s, p, d, f).
  3. mₗ = Magnetic (orbital orientation).
  4. mₛ = Spin (+½ or -½).

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Aufbau Principle

• Electrons fill orbitals in increasing energy order.
• Lowest energy orbitals are filled first.
• Meaning → 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p ...

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Pauli’s Exclusion Principle

• No two electrons in an atom can have the same 4 quantum numbers.
• At most 2 e⁻ per orbital, with opposite spins.

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Hund’s Rule

• Electrons occupy orbitals singly first, with parallel spins.
• After all orbitals are half-filled → pairing starts.

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Atomic Radius

• Distance from nucleus → outermost electron.
• Measured as covalent radius (½ distance between 2 bonded nuclei).
• Units: pm (10⁻¹² m) or Å (10⁻¹⁰ m).
• Trend: ↓ across period, ↑ down group.

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Ionization Energy

• Definition: Energy required to remove an e⁻ from an atom.
• Always endothermic (energy absorbed).
• Trends:
  • Larger atom → lower ionization energy.
  • Inversely proportional to atomic radius.
  • ↑ across period, ↓ down group.

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Electron Affinity

• Definition: Energy change when an atom gains an electron.
• Usually exothermic (energy released), but exceptions exist.
• High EA = atom easily accepts e⁻ (e.g., halogens).

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Electronegativity

• Definition: Atom’s ability to attract shared e⁻ in a bond.
• Scale: 0 → 4 (Pauling units).
• Metals: low EN (electron donors).
• Nonmetals: high EN (electron takers).
• Trend: ↑ across period, ↓ down group.

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Ionic Radius

• Cations: smaller than neutral atom (loss of shell).
• Anions: larger than neutral atom (extra e⁻ increases repulsion).
• Trend: Cations shrink, Anions expand.

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14. Elements & Electron Configuration

• Electron configuration = distribution of electrons in orbitals.
• Example (Fe, Z=26): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ or [Ar] 4s² 3d⁶.

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15. Chemical Bonding

• Lewis Structure: Shows bonds + lone pairs.
• Bond Order:
  • MO Theory → ½ (bonding e⁻ – antibonding e⁻).
  • Simple → number of bonds (O=O → bond order 2).
• EAN Rule: EAN = Atomic no. – Ox. state + (2 × Coordination no.).
  • Example: K₄[Fe(CN)₆] → EAN = 36 (Kr).

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16. Colloids

• Definition: Particles (1–1000 nm) dispersed in medium.
• Types:
  • Sol (solid in liquid) → paints.
  • Gel (liquid in solid) → jelly.
  • Foam (gas in liquid) → whipped cream.
  • Emulsion (liquid in liquid) → milk.
  • Aerosol (liquid/solid in gas) → fog, smoke.
• Properties: Tyndall effect, electrophoresis, coagulation.
• Applications: Medicines, food (milk, butter), soaps, industry.

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Extra Topic: How Soap Works

• Soap molecule = hydrophobic tail + hydrophilic head.
• At high conc. → form micelles.
• Tails trap grease inside, heads interact with water → grease washed away.