Quick Revision
Compressed notes for fast review.
Chapter 1 - Matter, Energy, and Origins
1. Classes of Matter
- Pure substances: fixed composition.
- Elements: cannot be broken chemically (O2, Fe).
- Compounds: two or more elements in fixed ratio (H2O, NaCl).
- Mixtures: combinations separated by physical means.
- Homogeneous (solutions): uniform (saltwater, air).
- Heterogeneous: non-uniform (sand + water, granite).
2. Properties of Matter
- Physical properties: observed without changing identity (color, melting point, density).
- Chemical properties: describe reactivity (flammability, reaction with acid).
3. States of Matter
- Solid: definite shape and volume, tightly packed particles.
- Liquid: definite volume, indefinite shape, flowing particles.
- Gas: no definite shape or volume, particles move randomly.
- Changes: melting, freezing, vaporization, condensation, sublimation.
Chapter 2 - Atoms, Ions, and Compounds
1. Isotopes & Average Atomic Mass
- Isotopes: same number of protons, different neutrons. Example: Fe-56 has 26 protons, 30 neutrons.
- Average atomic mass = (fraction × mass) for each isotope. Example (Neon): (0.904838 × 19.9924) + (0.002696 × 20.9940) + (0.092465 × 21.9914) = 20.1797 amu
2. Naming & Writing Formulas
- Ionic compounds (metal + nonmetal):
- NaCl = sodium chloride
- FeCl2 = iron(II) chloride
- Al3+ and O2- → Al2O3
- Molecular compounds (nonmetal + nonmetal):
- N2O5 = dinitrogen pentoxide
- P4S10 = tetraphosphorus decasulfide
3. Empirical vs Molecular Formulas
- Empirical: simplest ratio (CH2O).
- Molecular: actual numbers (C6H12O6).
- Relationship: Molecular formula = n × Empirical formula, where n = molar mass ÷ empirical mass.
Chapter 3 - Chemical Reactions & Earth's Composition
1. The Mole & Molar Mass
- 1 mole = 6.022 × 10^23 particles.
- Molar mass = atomic or molecular mass in grams. Example: CO2 = 44.01 g/mol.
2. Mass ⇄ Moles ⇄ Particles
- Mass → Moles (divide by molar mass).
- Moles → Particles (multiply by Avogadro’s number). Example: 5.32 mol CaCO3 × 6.022 × 10^23 = 3.20 × 10^24 molecules.
3. Percent Composition
% element = (mass of element in 1 mol compound ÷ molar mass) × 100 Example: Fe2O3 = (2 × 55.85 ÷ 159.70) × 100 = 69.94% Fe
4. Limiting Reactant
Steps:
- Balance equation.
- Convert masses → moles.
- Compare mole ratios.
- Smaller product = limiting reactant. Example: CH4 + 2 O2 → CO2 + 2 H2O 10.0 g CH4 → 0.623 mol 20.0 g O2 → 0.625 mol O2 limits, produces 11.3 g H2O.
Chapter 4 - Solution Chemistry & the Hydrosphere
1. Molarity
M = moles of solute ÷ liters of solution Example: 36.5 g BaCl2 ÷ 208.23 g/mol = 0.175 mol. Solution volume = 0.750 L → M = 0.175 ÷ 0.750 = 0.233 M.
2. Acid-Base Neutralization
Reaction: H+ + OH- → H2O Titration: MaVa = MbVb (for 1:1 ratio). For polyprotic acids: multiply by number of acidic protons. Example: H2SO4 + NaOH → 2 NaOH needed per H2SO4.
Chapter 5 - Thermochemistry
1. Energy
- Kinetic energy = 1/2 m v^2
- Potential energy = stored energy (bonds, position).
2. Heat Flow
- Endothermic: absorbs heat (+q, +ΔH).
- Exothermic: releases heat (−q, −ΔH).
3. First Law of Thermodynamics
ΔE = q + w w = −PΔV (for gas expansion/compression).
4. Calorimetry
q = m × c × ΔT Example: 25.0 g water, c = 4.18 J/g°C, ΔT = 15.0°C → q = 1568 J. For reactions: q_reaction = −q_calorimeter.
Chapter 6 - Properties of Gases
1. Pressure Units
1 atm = 101.3 kPa = 101325 Pa = 760 mmHg = 760 torr.
2. Ideal Gas Law
PV = nRT R = 0.08206 L·atm·mol⁻¹·K⁻¹
Example: Gas sample
- Mass = 0.495 g, V = 0.127 L, T = 371 K, P = 0.992 atm
- n = (0.992 × 0.127) ÷ (0.08206 × 371) = 0.00413 mol
- M = 0.495 ÷ 0.00413 = 120 g/mol
Chapter 7 - Chemical Bonding
1. Lewis Symbols & Octet Rule
- Lewis symbol = element symbol with dots for valence electrons.
- Atoms bond to complete octet (or duet for H).
2. Drawing Lewis Structures
- Count valence electrons.
- Place central atom (least electronegative, not H).
- Connect with single bonds.
- Complete octets with lone pairs.
- Use double/triple bonds if needed.
- Check formal charges.
Formal charge = valence electrons − (lone pairs + 1/2 bonding electrons).
3. Electronegativity & Bond Types
- EN trend: increases across period, decreases down group.
- Nonpolar covalent: ΔEN ≈ 0 (Cl2).
- Polar covalent: ΔEN between 0–2 (HCl).
- Ionic: ΔEN > 2 (NaCl).
Extra Topics & Quick Cheatsheet
Neutralization
HA + BOH → BA + H2O
pH
pH = −log[H+] pOH = −log[OH−] pH + pOH = 14 at 25°C
Example: 0.0010 M HCl → pH = 3.00
Equilibrium Constants
For aA + bB ⇌ cC + dD:
- Kc = [C]^c [D]^d ÷ [A]^a [B]^b
- Kp = Kc (RT)^Δn
Quick Trends
- Atomic radius: decreases across period, increases down group.
- Ionization energy: increases across period, decreases down group.
- Electronegativity: increases across period, decreases down group.
- Stoichiometry path: mass → moles → mole ratio → moles product → mass product.
- Gas law: always convert units (L, atm, K).
- Titration: MaVa = MbVb (adjust for polyprotic acids).
Constants & Conversions
- Avogadro's number = 6.022 × 10^23 mol⁻¹
- R = 0.08206 L·atm·mol⁻¹·K⁻¹
- 1 atm = 101.3 kPa = 101325 Pa = 760 mmHg
- Specific heat of water = 4.18 J/g°C