Chemical reactions follow fixed laws. Law of Conservation of Mass (Lavoisier, 1774): In a chemical reaction, the total mass of reactants equals the total mass of products. Matter is neither created nor destroyed. Example: 2H₂ + O₂ → 2H₂O. Mass of 2H₂ + mass of O₂ = mass of 2H₂O. Law of Constant Proportions/Definite Proportions (Proust, 1799): A pure chemical compound always contains the same elements combined in the same fixed ratio by mass. Example: Water always contains H and O in 1:8 mass ratio regardless of source. Law of Multiple Proportions (Dalton): When two elements form more than one compound, masses of one element combining with fixed mass of other are in simple ratios. Example: CO and CO₂.

John Dalton (1808) proposed the first scientific atomic theory in his book "A New System of Chemical Philosophy." Postulates: (1) Matter is made of tiny, indivisible particles called atoms. (2) Atoms of the same element are identical in mass and properties. (3) Atoms of different elements have different masses and properties. (4) Atoms combine in simple whole number ratios to form compounds. (5) Atoms cannot be created, destroyed, or converted into other atoms in chemical reactions. Successes: Explained laws of conservation of mass and constant proportions. Failures: Atoms are divisible (electrons, protons, neutrons). Isotopes disprove that all atoms of an element are identical. Compounds do not always combine in simple ratios (e.g., sucrose C₁₂H₂₂O₁₁).

An atom is the smallest particle of an element that can exist independently and take part in chemical reactions. Atomic symbols: Each element has a one or two-letter symbol. First letter is always capital; second (if present) is lowercase. Examples: H (hydrogen), He (helium), C (carbon), Ca (calcium), Fe (iron — from Latin Ferrum), Na (sodium — Natrium), K (potassium — Kalium). Atomic size: atoms are extremely small — radius ~1 Å = 10⁻¹⁰ m. Atomic mass: mass of an atom. Originally measured relative to hydrogen (mass = 1). Now, carbon-12 is the standard — 1 atomic mass unit (u or amu) = 1/12th the mass of one carbon-12 atom = 1.66 × 10⁻²⁷ kg. Atomic masses: H = 1 u, C = 12 u, O = 16 u, N = 14 u, S = 32 u, Cl = 35.5 u, Na = 23 u, Fe = 56 u.

A molecule is the smallest particle of a substance (element or compound) that can exist independently and has all the properties of that substance. It is formed by the combination of two or more atoms. Molecules of elements: made of same type of atoms. Examples: H₂ (hydrogen), O₂ (oxygen), N₂ (nitrogen), Cl₂ (chlorine), P₄ (phosphorus), S₈ (sulfur). Monatomic: noble gases (He, Ne, Ar) exist as single atoms. Molecules of compounds: made of different atoms. Examples: H₂O (water — 2H, 1O), CO₂ (carbon dioxide — 1C, 2O), NH₃ (ammonia), HCl, H₂SO₄, NaCl (ionic compound). Chemical formula: representation showing types and number of atoms. Subscript = number of atoms. H₂O means 2 hydrogen atoms + 1 oxygen atom per molecule.

Ions are charged particles formed when atoms lose or gain electrons. Cation: positively charged ion (atom loses electrons). Examples: Na⁺, K⁺, Ca²⁺, Mg²⁺, Al³⁺. Anion: negatively charged ion (atom gains electrons). Examples: Cl⁻, O²⁻, N³⁻, S²⁻. Polyatomic ions: ions made of multiple atoms bonded together with a charge. Examples: OH⁻ (hydroxide), SO₄²⁻ (sulfate), NO₃⁻ (nitrate), NH₄⁺ (ammonium), CO₃²⁻ (carbonate), PO₄³⁻ (phosphate). Writing chemical formulae using valency (cross-multiplication method): (1) Write symbols of elements/ions. (2) Write their valencies below. (3) Cross-multiply valencies as subscripts (reduce to simplest ratio). Example: Calcium chloride — Ca²⁺ and Cl⁻ → CaCl₂. Magnesium oxide — Mg²⁺ and O²⁻ → MgO (ratio 2:2 = 1:1).

The mole is the unit for counting atoms/molecules (SI unit). 1 mole = 6.022 × 10²³ particles (atoms, molecules, or ions). This number is called Avogadro's number (Nₐ). Molar mass: mass of 1 mole of a substance in grams = numerically equal to atomic/molecular mass in u. Examples: 1 mole of C = 12 g; 1 mole of H₂O = 18 g; 1 mole of NaCl = 58.5 g. Calculations: Moles = given mass / molar mass. Number of particles = moles × Nₐ. Mass = moles × molar mass. Example: 36 g of water = 36/18 = 2 moles = 2 × 6.022 × 10²³ = 1.2044 × 10²⁴ molecules. Significance: The mole concept allows chemists to count atoms/molecules by weighing, since individual atoms are too small to count directly.

Molecular mass: sum of atomic masses of all atoms in one molecule of a substance. Examples: H₂O = 2(1) + 16 = 18 u. CO₂ = 12 + 2(16) = 44 u. H₂SO₄ = 2(1) + 32 + 4(16) = 98 u. NH₃ = 14 + 3(1) = 17 u. NaOH = 23 + 16 + 1 = 40 u. CaCO₃ = 40 + 12 + 3(16) = 100 u. Formula unit mass: used for ionic compounds (which do not exist as discrete molecules). Example: NaCl formula unit mass = 23 + 35.5 = 58.5 u. Relative molecular mass: ratio of molecular mass of substance to 1/12th mass of C-12. Since both are in same units, it has no unit. Note: atoms of noble gases exist as single atoms, so their atomic mass = molecular mass.

Percentage composition: the percentage by mass of each element in a compound. Formula: % of element = (mass of element in formula / molecular mass) × 100. Example: in H₂O: %H = (2/18) × 100 = 11.1%; %O = (16/18) × 100 = 88.9%. Empirical formula: simplest whole number ratio of atoms in a compound. Molecular formula: actual number of atoms in a molecule. Relationship: molecular formula = n × empirical formula (n = integer). Example: H₂O₂ has empirical formula HO, molecular formula H₂O₂. Glucose: empirical formula CH₂O, molecular formula C₆H₁₂O₆. Benzene: empirical formula CH, molecular formula C₆H₆. Knowing % composition allows calculation of empirical formula.

Atom: smallest particle of element; not independently stable for most elements. Molecule: smallest independently existing particle with chemical properties. Ion: charged particle (cation +ve, anion −ve). Atomic mass unit (u): 1/12th of C-12 mass. Mole: 6.022 × 10²³ particles. Molar mass (g/mol) = molecular/atomic mass numerically. Key atomic masses to remember: H=1, C=12, N=14, O=16, Na=23, Mg=24, Al=27, S=32, Cl=35.5, K=39, Ca=40, Fe=56, Cu=63.5, Zn=65. Laws of chemical combination: Conservation of mass (Lavoisier), Constant proportions (Proust). Dalton explained these with atomic theory. Avogadro's number: 6.022 × 10²³. Formula writing: use cross-multiplication of valencies.

One atomic mass unit (1 u) is 1/12 of the mass of one carbon-12 atom. Relative atomic mass is the mass of an atom compared with 1/12 of a carbon-12 atom. Values to remember: H = 1, He = 4, C = 12, N = 14, O = 16, F = 19, Na = 23, Mg = 24, Al = 27, Si = 28, P = 31, S = 32, Cl = 35.5, K = 39, Ca = 40, Fe = 56, Cu = 63.5, Zn = 65. Atomic masses are not exact whole numbers for some elements (like chlorine) because they are averages over isotopes.

Atomicity is the number of atoms in one molecule of an element. Monoatomic: noble gases such as He, Ne, Ar and most metals (written as Na, Fe, Cu). Diatomic: H₂, O₂, N₂, Cl₂, F₂, Br₂, I₂. Triatomic: ozone O₃. Tetra-atomic: phosphorus P₄. Poly-atomic: sulphur S₈. A molecule of a compound has atoms of different elements: H₂O has atomicity 3, CO₂ has 3, NH₃ has 4, H₂SO₄ has 7.

Cations (positive): sodium Na⁺, potassium K⁺, hydrogen H⁺, ammonium NH₄⁺, magnesium Mg²⁺, calcium Ca²⁺, zinc Zn²⁺, aluminium Al³⁺, iron(III) Fe³⁺. Anions (negative): chloride Cl⁻, hydroxide OH⁻, nitrate NO₃⁻, hydrogen carbonate HCO₃⁻, oxide O²⁻, sulphide S²⁻, carbonate CO₃²⁻, sulphate SO₄²⁻, phosphate PO₄³⁻. Polyatomic ions are groups of atoms that carry a net charge and behave as a single unit. A compound is neutral overall, so total positive charge = total negative charge.

Steps: (1) Write the symbols of the elements or ions with their charges. (2) Cross over the numerical value of the charges so that they become subscripts of the other ion. (3) Reduce to the simplest ratio and omit 1. (4) Put polyatomic ions in brackets if more than one is needed. Examples: Al³⁺ and O²⁻ give Al₂O₃. Ca²⁺ and Cl⁻ give CaCl₂. Mg²⁺ and OH⁻ give Mg(OH)₂. NH₄⁺ and SO₄²⁻ give (NH₄)₂SO₄. Ca²⁺ and PO₄³⁻ give Ca₃(PO₄)₂.

One mole is the amount of substance that contains 6.022 × 10²³ particles (atoms, molecules or ions); this number is Avogadro's constant. Number of moles = given mass ÷ molar mass = number of particles ÷ 6.022 × 10²³ = volume of a gas at STP ÷ 22.4 L. Molar mass is the mass of one mole in grams and is numerically equal to the atomic or molecular mass in u. Examples: 1 mole of water = 18 g = 6.022 × 10²³ molecules; 1 mole of CO₂ = 44 g; 1 mole of any gas at STP occupies 22.4 L.

Example 1: moles in 22 g of CO₂ = 22 ÷ 44 = 0.5 mol; molecules = 0.5 × 6.022 × 10²³ = 3.011 × 10²³. Example 2: mass of 3 moles of H₂O = 3 × 18 = 54 g. Example 3: atoms in 9 g of water: moles = 9 ÷ 18 = 0.5; molecules = 3.011 × 10²³; each molecule has 3 atoms, so atoms = 9.033 × 10²³. Example 4: mass of one atom of carbon = 12 ÷ (6.022 × 10²³) = 1.99 × 10⁻²³ g. Always convert mass to moles first, then moles to particles.