Atoms and Molecules

The law of conservation of mass states that in a chemical reaction, mass is neither created nor destroyed. The total mass of reactants equals the total mass of products.

If you burn magnesium ribbon in a closed container and carefully collect all the products (including the magnesium oxide formed and the oxygen consumed), you will find the total mass before the reaction equals the total mass after.

This law was established experimentally by Antoine Lavoisier. It forms the basis of balancing chemical equations — whatever is on the left side of the arrow must equal the right side by mass.

The law of constant proportions (also called the law of definite proportions) states that a pure compound always contains the same elements in the same fixed proportion by mass, regardless of the source or method of preparation.

Pure water from the Ganga, from a Mumbai tap, from a bore well in Rajasthan, or prepared in the lab — all contain hydrogen and oxygen in the same mass ratio of 1 : 8.

This is why compounds are different from mixtures. You can make dilute or concentrated nimbu pani (a mixture) by varying proportions. But water is always H₂O in a fixed mass ratio — you cannot make "dilute water" or "concentrated water."

John Dalton proposed his atomic theory in 1808 to explain the laws of chemical combination.

Main postulates: (1) All 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 cannot be created or destroyed in a chemical reaction. (5) Compounds are formed when atoms of different elements combine in simple whole-number ratios.

Some of Dalton's postulates have been revised since — atoms are not indivisible (they have subatomic particles), and atoms of the same element can have different masses (isotopes). But the theory was a major scientific breakthrough that explained chemical combination successfully.

An atom is the smallest particle of an element that can take part in a chemical reaction. Atoms are too small to be seen directly; their sizes are on the order of 0.1 nanometres.

Each element has a chemical symbol — a one or two-letter abbreviation. Many symbols come from the element's Latin or Greek name, which is why some symbols do not match the English name.

The actual mass of an atom is extremely small and inconvenient to use in everyday calculations. So atomic mass is defined as a relative mass — the mass of an atom relative to one-twelfth the mass of a carbon-12 atom.

This relative scale is dimensionless and denoted u (unified atomic mass unit) or amu. On this scale, hydrogen ≈ 1 u, carbon = 12 u, oxygen = 16 u, sodium = 23 u, iron = 56 u.

Atomic mass is a weighted average of the masses of all naturally occurring isotopes of the element. This is why atomic mass values are often not whole numbers (e.g., chlorine ≈ 35.5 u).

A molecule is the smallest particle of a substance that can exist independently and retains all the chemical properties of that substance.

Molecules of elements may contain one type of atom: oxygen exists as O₂, nitrogen as N₂, hydrogen as H₂, and ozone as O₃. Noble gases like helium and argon exist as single atoms.

Molecules of compounds contain atoms of different elements: water (H₂O), carbon dioxide (CO₂), glucose (C₆H₁₂O₆), and common salt is an ionic compound (NaCl). A chemical formula shows the types and numbers of atoms in a molecule.

An ion is a charged particle formed when an atom gains or loses electrons. A cation is a positively charged ion (loses electrons): Na⁺, Ca²⁺, Al³⁺. An anion is a negatively charged ion (gains electrons): Cl⁻, O²⁻, N³⁻.

Valency is the combining capacity of an atom — how many electrons it loses, gains, or shares to achieve stability. Sodium has valency 1; magnesium has valency 2; aluminium has valency 3.

Polyatomic ions are groups of atoms that carry a charge and act as a single unit in reactions: sulphate (SO₄²⁻), carbonate (CO₃²⁻), nitrate (NO₃⁻), ammonium (NH₄⁺), hydroxide (OH⁻), phosphate (PO₄³⁻).

To write the formula of a binary ionic compound: (1) write the cation first, then the anion. (2) cross the valencies — the valency of one becomes the subscript of the other. (3) simplify if needed.

Sodium oxide: Na⁺ and O²⁻. Valencies are 1 and 2 respectively. Crossing: Na₂O. Simplify: Na₂O (already in lowest ratio).

Calcium chloride: Ca²⁺ and Cl⁻. Valencies 2 and 1. Crossing: CaCl₂.

For polyatomic ions, use brackets when the subscript is more than 1. Calcium hydroxide: Ca²⁺ and OH⁻. Formula = Ca(OH)₂.

Molecular mass is the sum of atomic masses of all atoms present in one molecule of a substance.

Molecular mass of water (H₂O) = 2 × 1 + 16 = 18 u.

For ionic compounds like NaCl (which does not have individual molecules), we calculate the formula mass: Na + Cl = 23 + 35.5 = 58.5 u.

A mole is the SI unit for the amount of substance. One mole of any substance contains the same fixed number of particles — approximately 6.022 × 10²³ particles. This number is called Avogadro's number (Nₐ).

Just as "a dozen" means 12 of anything — 12 eggs, 12 mangoes, 12 atoms — "a mole" means 6.022 × 10²³ of anything. One mole of rice grains, one mole of water molecules, or one mole of electrons all contain 6.022 × 10²³ particles.

The mole bridges the microscopic world (individual atoms and molecules) to the macroscopic world (grams that can be weighed in a chemistry laboratory). This is why chemists use moles instead of trying to count atoms directly.

The molar mass of a substance is the mass of one mole of that substance, expressed in grams per mole (g/mol). It has the same numerical value as the molecular mass or atomic mass in u.

Molar mass of water = 18 g/mol (molecular mass = 18 u). So 18 g of water is exactly one mole and contains 6.022 × 10²³ water molecules.

To find the number of moles in a given mass, divide by the molar mass. To find the number of particles, multiply the number of moles by Avogadro's number.

Law of conservation of mass: total mass of reactants = total mass of products.

Law of constant proportions: fixed element ratio by mass in a pure compound.

Valency: combining capacity. Cross valencies to write formulae. Use brackets for polyatomic ions.

Molecular mass = sum of atomic masses. Molar mass has same number in g/mol.

Mole: . Avogadro's number: .

Number of particles: .