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Chapter 1: Stoichiometry
Atomic Mass Unit (amu)
1 amu = 1/12 of mass of one carbon-12 atom = 1.66 × 10⁻²⁴ g
Mole Concept
1 mole = 6.022 × 10²³ particles (Avogadro's Number)
Mole = Mass / Molar Mass
Mole = Volume / Molar Volume (22.4 L at STP)
Empirical Formula
Simplest whole number ratio of atoms in a compound.
Compound: 40% C, 6.67% H, 53.33% O
Divide by atomic mass: C=3.33, H=6.67, O=3.33
Divide by smallest: C=1, H=2, O=1
Empirical Formula: CH₂O
Molecular Formula
Molecular Formula = n × Empirical Formula
n = Molecular Mass / Empirical Formula Mass
Stoichiometric Calculations
2H₂ + O₂ → 2H₂O
From 2 moles H₂ → 2 moles H₂O
From 4g H₂ → 36g H₂O
Limiting Reagent
Reagent that gets completely consumed and limits product formation.
Chapter 2: Atomic Structure
Sub-atomic Particles
- Proton: Mass = 1 amu, Charge = +1
- Neutron: Mass = 1 amu, Charge = 0
- Electron: Mass = 1/1836 amu, Charge = -1
Bohr's Atomic Model
Electrons revolve in fixed circular orbits (energy levels).
Eₙ = -R(1/n²)
Where: R = 2.18 × 10⁻¹⁸ J, n = principal quantum number
Quantum Numbers
- Principal (n): Energy level (1, 2, 3...)
- Azimuthal (l): Sub-shell (s=0, p=1, d=2, f=3)
- Magnetic (m): Orbital orientation (-l to +l)
- Spin (s): Electron spin (+1/2 or -1/2)
Electron Configuration
Aufbau Principle: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p...
Hund's Rule: Electrons fill orbitals singly first
Pauli Exclusion: Max 2 electrons per orbital, opposite spins
Isotopes
Same atomic number (Z), different mass number (A).
Carbon: ¹²C (98.9%), ¹³C (1.1%)
Chapter 3: Chemical Bonding
Types of Bonds
1. Ionic Bond
Complete transfer of electrons (metal + non-metal).
Na → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
NaCl = Na⁺Cl⁻
2. Covalent Bond
Sharing of electrons (non-metals).
- Single: H-H
- Double: O=O
- Triple: N≡N
3. Coordinate Bond
Both electrons from one atom.
VSEPR Theory
Valence Shell Electron Pair Repulsion - electron pairs repel, determine shape.
Hybridization
- sp: 2 bonds, linear (180°)
- sp²: 3 bonds, trigonal planar (120°)
- sp³: 4 bonds, tetrahedral (109.5°)
- sp³d: 5 bonds, trigonal bipyramidal
- sp³d²: 6 bonds, octahedral
Intermolecular Forces
- Van der Waals: Weakest
- Dipole-Dipole: Moderate
- Hydrogen Bond: Strongest (H with F, O, N)
Chapter 4: States of Matter
Gases
Gas Laws
Boyle's Law: P₁V₁ = P₂V₂ (T constant)
Charles Law: V₁/T₁ = V₂/T₂ (P constant)
Avogadro's Law: V₁/n₁ = V₂/n₂ (T, P constant)
Ideal Gas Equation: PV = nRT
Kinetic Molecular Theory
- Gas particles move randomly
- No intermolecular forces
- Collisions are elastic
- Volume of particles negligible
Liquids
- Viscosity: Resistance to flow
- Surface Tension: Energy to increase surface
- Boiling Point: Vapor pressure = external pressure
Solids
- Crystalline: Regular arrangement (NaCl, diamond)
- Amorphous: Irregular (glass, rubber)
Chapter 5: Chemical Equilibrium
Reversible Reactions
aA + bB ⇌ cC + dD
Equilibrium Constant
Kc = [C]^c [D]^d / [A]^a [B]^b
Kp = P_C^c P_D^d / P_A^a P_B^b
Le Chatelier's Principle
System shifts to counteract change.
- Increase [Reactant] → shifts right
- Increase Pressure → shifts to fewer moles
- Increase Temperature → shifts endothermic
- Catalyst → no effect on K
Chapter 6: Acids, Bases, and Salts
Acid-Base Theories
Arrhenius
Acid: H⁺ donor, Base: OH⁻ donor
Bronsted-Lowry
Acid: Proton donor, Base: Proton acceptor
Lewis
Acid: Electron pair acceptor, Base: Electron pair donor
pH and pOH
pH = -log[H⁺]
pOH = -log[OH⁻]
pH + pOH = 14
Strength of Acids
Ka = [H⁺][A⁻] / [HA]
pKa = -log Ka
Buffers
Resist pH change. Mixture of weak acid and its salt (or weak base and its salt).
Chapter 7: Thermochemistry
Enthalpy (H)
Heat content of system at constant pressure.
Heat of Reaction
ΔH = Σ(ΔH_f products) - Σ(ΔH_f reactants)
Hess's Law
Total enthalpy change independent of route.
Bond Energy
ΔH = Σ(BE broken) - Σ(BE formed)
Types of Enthalpy
- ΔH⁰_f: Formation from elements
- ΔH⁰_d: Dissociation
- ΔH⁰_comb: Combustion
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