9.1 Matter, Atoms & Molecules

Everything you can touch, breathe, taste or weigh is matter — anything that has mass and occupies space. Indian Railways expects every Group D candidate to know the CBSE Class 10 picture of what matter is made of and how it is counted. This section builds that picture from the ground up, because every later chemistry question — reactions, acids, metals, carbon compounds — rests on these foundations.

States of Matter

Matter exists in three physical states, distinguished by how strongly its particles attract each other and how freely they move:

Adding heat energy lets particles overcome inter-particle forces: solid → liquid (fusion) → gas (vaporisation). Removing heat reverses the chain. The temperature at which a solid becomes liquid is its melting point; the temperature at which a liquid becomes gas (at atmospheric pressure) is its boiling point.

Atoms and Molecules

John Dalton's atomic theory (1808) proposed that all matter is made of indivisible particles called atoms. Atoms are the smallest unit of an element that participates in a chemical reaction. They combine in fixed whole-number ratios to form molecules — the smallest particle of an element or compound capable of independent existence while keeping its chemical identity.

• Element molecules: O₂ (oxygen, diatomic), O₃ (ozone, triatomic), N₂ (nitrogen), H₂ (hydrogen), P₄ (white phosphorus, tetra-atomic), S₈ (sulphur, octa-atomic). Noble gases (He, Ne, Ar) are monoatomic.
• Compound molecules: H₂O (water), CO₂ (carbon dioxide), NH₃ (ammonia), HCl, NaCl (formula unit, not discrete molecule).

The number of atoms in one molecule of an element is its atomicity. O₂ is diatomic, O₃ is triatomic, S₈ is polyatomic.

Atomic and Molecular Mass

Masses of single atoms are tiny, so chemists use the unified atomic mass unit (u), defined as exactly 1/12 the mass of a carbon-12 atom. Hydrogen ≈ 1 u, carbon = 12 u, oxygen = 16 u, nitrogen = 14 u.

• Atomic mass of an element = average mass of its atoms in u (averaged over natural isotopes).
• Molecular mass = sum of the atomic masses of all atoms in the molecule. Example: H₂O = (2 × 1) + 16 = 18 u. CO₂ = 12 + (2 × 16) = 44 u.

The Mole Concept

The mole is the chemist's counting unit. One mole of any substance contains exactly 6.022 × 10²³ particles (Avogadro's number). It bridges the invisible atomic scale with grams you can weigh on a balance.

• Mass of 1 mole of atoms (in grams) = atomic mass in grams. So 1 mole of O atoms = 16 g; 1 mole of O₂ molecules = 32 g.
• 1 mole of any gas at STP (0 °C, 1 atm) occupies 22.4 L.
• Number of particles = (mass in g / molar mass) × 6.022 × 10²³.

Example: How many molecules are in 9 g of water? Molar mass of H₂O = 18 g/mol, so 9 g = 0.5 mol → 0.5 × 6.022 × 10²³ = 3.011 × 10²³ molecules.

Structure of the Atom

An atom has a dense, positively charged nucleus (protons + neutrons) surrounded by negatively charged electrons arranged in shells (K, L, M, N …).

• Atomic number (Z) = number of protons; defines the element.
• Mass number (A) = protons + neutrons.
• Electronic configuration: shell capacities follow 2n² — K = 2, L = 8, M = 18, N = 32.
• Valence electrons = electrons in the outermost shell; they decide how an atom reacts. Sodium (2,8,1) has 1 valence electron and is highly reactive; chlorine (2,8,7) needs 1 more to complete its octet.

Isotopes are atoms of the same element (same Z) with different neutron numbers (different A). Hydrogen's isotopes — protium (¹H, no neutron), deuterium (²H, 1 neutron), tritium (³H, 2 neutrons) — are the classic exam example. Carbon-12 and carbon-14, uranium-235 and uranium-238 are others. Chemical properties of isotopes are identical (same electrons); physical properties differ slightly with mass.

RRB Exam Scenarios and Traps

• Trap: treating NaCl as a discrete molecule. NaCl is a formula unit of a giant ionic lattice; the formula mass (58.5 u) is used in mole calculations, not a molecular mass.
• Trap: confusing atomic mass with mass number. Atomic mass on the periodic table (e.g., Cl = 35.5 u) is the weighted average over isotopes; mass number is a count of nucleons in one isotope (e.g., ³⁵Cl, ³⁷Cl).
• Expect a calculation: 'How many moles in 22 g of CO₂?' Molar mass = 44 g/mol → 0.5 mol.
• Expect: 'Which is monoatomic?' Answer: noble gases.
• Expect: 'Number of neutrons in ²³Na₁₁?' → 23 − 11 = 12.

Use the table and the Avogadro shortcut (6.022 × 10²³) repeatedly until they are reflex — RRB recycles mole and isotope questions every cycle.