Chemistry - Regular Syllabus

Unit 1: Introduction to Chemistry

What is Chemistry?

  • The study of matter and how it changes.

  • Matter: anything that has mass and takes up space.

Branches of Chemistry

  • Organic, Inorganic, Physical, Analytical, Biochemistry.

Scientific Method

  • Steps: Observation → Question → Hypothesis → Experiment → Conclusion → Repeat.

Measurements

  • Use SI units (mass = grams, volume = liters, length = meters, temperature = Celsius/Kelvin).

  • Accuracy = closeness to true value; Precision = repeatability.

  • Use significant figures correctly in calculations.

Unit 2: Matter & Its Properties

States of Matter

  • Solid: definite shape and volume.

  • Liquid: definite volume, no definite shape.

  • Gas: no definite shape or volume.

Physical vs. Chemical Properties

  • Physical: observed without changing identity (e.g., melting point).

  • Chemical: describes how a substance reacts (e.g., flammability).

Physical vs. Chemical Changes

  • Physical: change in appearance, not composition (melting, cutting).

  • Chemical: new substance formed (burning, rusting).

Mixtures and Pure Substances

  • Mixture: two or more substances physically combined.

    • Homogeneous (uniform, e.g., salt water)

    • Heterogeneous (not uniform, e.g., salad)

  • Pure substance: element or compound.

Unit 3: Atomic Structure

Basic Atomic Structure

  • Atom = smallest particle of an element.

  • Proton: +1 charge, in nucleus.

  • Neutron: 0 charge, in nucleus.

  • Electron: –1 charge, outside nucleus.

Atomic Number & Mass Number

  • Atomic number = number of protons = number of electrons in neutral atom.

  • Mass number = protons + neutrons.

Isotopes

  • Atoms of the same element with different numbers of neutrons.

Atomic Models

  • Dalton (solid sphere), Thomson (electrons in atom), Rutherford (nucleus), Bohr (energy levels), Modern (electron cloud/orbitals).

Unit 4: Periodic Table

Organization

  • Elements arranged by increasing atomic number.

  • Groups/families = vertical columns (same valence electrons, similar properties).

  • Periods = horizontal rows (same energy level).

Metals, Nonmetals, Metalloids

  • Metals: shiny, conductive, malleable, left side.

  • Nonmetals: dull, brittle, poor conductors, right side.

  • Metalloids: properties of both.

Periodic Trends

  • Atomic size: increases down, decreases across.

  • Ionization energy: decreases down, increases across.

  • Electronegativity: decreases down, increases across.

Unit 5: Chemical Bonding

Why Atoms Bond

  • To become more stable (achieve full outer shell).

Types of Bonds

  • Ionic: metal + nonmetal, electron transfer.

  • Covalent: nonmetal + nonmetal, electron sharing.

  • Metallic: metal + metal, electrons shared in a “sea”.

Naming Compounds

  • Ionic: name metal, then nonmetal with “-ide” (e.g., NaCl = sodium chloride).

  • Covalent: use prefixes (mono-, di-, tri-, etc.), e.g., CO₂ = carbon dioxide.

Chemical Formulas

  • Subscripts tell how many of each atom (H₂O = 2 hydrogen, 1 oxygen).

Unit 6: Chemical Reactions

Signs of Chemical Reactions

  • Color change, gas production, temperature change, precipitate formation.

Types of Reactions

  • Synthesis: A + B → AB

  • Decomposition: AB → A + B

  • Single Replacement: A + BC → AC + B

  • Double Replacement: AB + CD → AD + CB

  • Combustion: fuel + O₂ → CO₂ + H₂O

Balancing Equations

  • Make sure the number of atoms of each element is the same on both sides.

Unit 7: The Mole & Stoichiometry

Mole Concept

  • 1 mole = 6.022 × 10²³ particles (Avogadro’s number).

  • Molar mass = mass of 1 mole of a substance (g/mol).

Conversions

  • Grams ↔ Moles ↔ Particles

  • Use molar mass, mole ratios, and Avogadro’s number.

Stoichiometry

  • Use balanced chemical equations to calculate amounts of reactants/products.

  • Limiting reactant: substance used up first.

  • Percent yield = (actual ÷ theoretical) × 100

Unit 8: States of Matter & Gas Laws

Kinetic Molecular Theory

  • Gases move randomly and constantly.

  • No attractions between particles (ideal gas assumption).

Gas Laws

  • Boyle’s Law: P₁V₁ = P₂V₂ (pressure and volume)

  • Charles’s Law: V₁/T₁ = V₂/T₂ (volume and temperature)

  • Gay-Lussac’s Law: P₁/T₁ = P₂/T₂ (pressure and temperature)

  • Combined Gas Law: P₁V₁/T₁ = P₂V₂/T₂

  • Ideal Gas Law: PV = nRT

Unit 9: Solutions

Solutions

  • Homogeneous mixture of solute (dissolved) and solvent (does the dissolving).

  • Water = universal solvent.

Solubility

  • Depends on temperature and pressure.

  • “Like dissolves like” (polar dissolves polar, etc.)

Concentration

  • Molarity (M) = moles of solute / liters of solution.

Dilution

  • M₁V₁ = M₂V₂ (to find new concentration or volume after adding solvent).

Unit 10: Acids and Bases

Properties

  • Acids: sour, pH < 7, produce H⁺ in solution.

  • Bases: bitter, slippery, pH > 7, produce OH⁻ in solution.

pH Scale

  • Ranges from 0–14.

  • 7 = neutral (pure water).

  • Lower = more acidic, higher = more basic.

Neutralization Reactions

  • Acid + Base → Salt + Water

  • Used in titrations to determine unknown concentrations.

Unit 11: Energy & Heat (Thermochemistry)

Heat vs. Temperature

  • Heat = energy transfer (measured in joules).

  • Temperature = measure of particle motion.

Heat Calculations

  • q = mcΔT

    • q = heat

    • m = mass

    • c = specific heat

    • ΔT = temperature change

Endothermic vs. Exothermic

  • Endothermic: absorbs heat (gets cold).

  • Exothermic: releases heat (gets hot).