Atoms and Molecules — Class 9 (CBSE)

Hand a child two cups, one of salt water and one of sugar water. They look identical. Smell them — identical. Yet one is salty, one sweet. The difference is invisible: different atoms arranged into different molecules. This chapter is about understanding stuff at the level where stuff becomes itself.

1. The story — from philosophers to Dalton

In §1 of the previous chapter we met Maharishi Kanada (~600 BCE) and Democritus (~400 BCE), both of whom proposed that matter is made of tiny indivisible particles ("anu" and "atomos" — meaning uncuttable). For 2200 years that idea remained philosophy, not science. Then in 1808, English chemist John Dalton put together quantitative experimental results from Lavoisier, Proust and others, and formalised the first scientific atomic theory.

Dalton's five postulates:

1. All matter is made of tiny indivisible particles called atoms.
2. Atoms of the same element are identical in mass and chemical properties.
3. Atoms of different elements have different masses and properties.
4. Atoms combine in small whole-number ratios to form compounds.
5. In chemical reactions, atoms are neither created nor destroyed — only rearranged.

Three of these turned out to be slightly wrong:

• Atoms ARE divisible (electrons, protons, neutrons discovered in next chapter).
• Atoms of the same element CAN have different masses (isotopes).
• Nuclear reactions DO change atoms into other atoms (fusion, fission).

But for ordinary chemistry — making and breaking bonds, mixing reagents in beakers — Dalton's theory is essentially right. It's the foundation on which all of high-school chemistry is built.

2. Laws of chemical combination

Before Dalton, two laws — established by careful weighing experiments — already pointed at the existence of atoms.

Law of conservation of mass (Lavoisier, 1789)

Mass is neither created nor destroyed in a chemical reaction.

If you start with 10 g of reactants in a sealed container, you end with exactly 10 g of products. The atoms are just rearranged.

Example: $\underset{\text{2 g + 16 g = 18 g}}{\underbrace{{\text{H}}_2+\frac{1}{2}{\text{O}}_2}}\to \underset{\text{18 g}}{\underbrace{{\text{H}}_2\text{O}}}$

Law of constant proportions / definite proportions (Proust, 1799)

A pure chemical compound always contains the same elements in the same proportion by mass, regardless of where it came from.

Water from a Himalayan glacier, water from your kitchen tap, water synthesised in a lab — all have hydrogen : oxygen = 1 : 8 by mass (or H : O = 2 : 1 by atoms).

Both laws were neatly explained by Dalton's "atoms in fixed whole-number ratios" postulate — and that's why his theory caught on.

3. Atomic mass — and why it's relative

We can't put a single atom on a weighing scale; it's far too small. So instead, chemists use relative atomic masses: how heavy an atom of one element is compared with an atom of a chosen standard.

Three standards in history

1. Hydrogen (atomic mass = 1) — used initially because H is the lightest.
2. Oxygen (atomic mass = 16) — used in the late 19th and early 20th century.
3. Carbon-12 (atomic mass = 12) — current standard since 1961.

The modern definition: one atomic mass unit (u, also called dalton) is 1/12 the mass of one atom of carbon-12.

$1u=1.6605\times 10^{-27}\text{kg}$

Atomic masses you must memorise (Class 9 essentials)

Why is chlorine 35.5? Because natural chlorine is a 75:25 mix of Cl-35 and Cl-37 — its "atomic mass" is the weighted average.

4. Molecules and ions

A molecule is the smallest particle of an element or compound that can exist on its own and still retain that substance's chemical properties.

Molecules of elements

Most non-metals exist as molecules of fixed size:

• ${\text{H}}_2$, ${\text{O}}_2$, ${\text{N}}_2$, ${\text{Cl}}_2$, ${\text{F}}_2$, ${\text{Br}}_2$, ${\text{I}}_2$ — diatomic (2 atoms).
• ${\text{O}}_3$ (ozone) — triatomic.
• ${\text{P}}_4$ (white phosphorus) — tetra-atomic.
• ${\text{S}}_8$ — octa-atomic.
• Inert/noble gases (He, Ne, Ar) — monatomic.

This count is the atomicity of the element.

Molecules of compounds

Atoms of different elements bonded together: H₂O, CO₂, NH₃, CH₄, H₂SO₄, NaCl (debated — see ions below).

Ions — charged particles

When atoms gain or lose electrons, they become charged → ions.

• Cation (positive): lost electrons. Na → Na⁺ + e⁻.
• Anion (negative): gained electrons. Cl + e⁻ → Cl⁻.

Compounds like NaCl, CaCO₃, KNO₃ are made of ions held together by electrostatic forces — they don't form discrete molecules, but you'll still write "formulas" for them (these are called formula units).

Common ions to memorise

Valency rules for writing formulas

Cross-multiply the valencies and simplify:

Example 1 — sodium chloride:

• Na is +1, Cl is −1. Cross-multiply: Na₁Cl₁ → NaCl.

Example 2 — aluminium oxide:

• Al is +3, O is −2. Cross-multiply: Al₂O₃ → Al₂O₃.

Example 3 — calcium phosphate:

• Ca is +2, PO₄ is −3. Cross-multiply: Ca₃(PO₄)₂ → Ca₃(PO₄)₂.
• Use brackets when the polyatomic ion is taken more than once.

5. Molecular mass and formula unit mass

Molecular mass = sum of atomic masses of all atoms in one molecule.

Example — water, H₂O: $M({\text{H}}_2\text{O})=2\times 1+1\times 16=18u$

Example — sulphuric acid, H₂SO₄: $M=2\times 1+1\times 32+4\times 16=2+32+64=98u$

Formula unit mass is the same calculation but used for ionic compounds (which don't form discrete molecules):

Example — sodium chloride: $\text{Formula unit mass}=23+35.5=58.5u$

Example — calcium carbonate, CaCO₃: $=40+12+3\times 16=100u$

The numerical work is identical to molecular mass; the terminology distinguishes ionic from molecular compounds.

6. The mole concept — the heart of chemistry

Why we need a "mole"

In a chemistry lab you weigh out grams. But chemical reactions happen at the level of individual atoms and molecules. We need a unit that bridges "grams I can weigh" and "atoms I can count."

That bridge is the mole.

Definition

One mole of any substance contains $6.022\times 10^{23}$ particles of that substance. This is Avogadro's number ($N_A$).

Why this specific number? It's defined so that 1 mole of carbon-12 atoms weighs exactly 12 grams. Same number trick:

• 1 mole of H atoms weighs 1 g.
• 1 mole of O atoms weighs 16 g.
• 1 mole of H₂O molecules weighs 18 g.

The mass in grams of 1 mole = atomic/molecular mass expressed in u. This is the molar mass.

The three relationships you'll use everywhere

$n=\frac{m}{M}=\frac{N}{N_A}$

Where:

• $n$ = number of moles
• $m$ = mass in grams
• $M$ = molar mass (g/mol)
• $N$ = number of particles
• $N_A$ = Avogadro's number ($6.022\times 10^{23}$)

Three quantities, three formulas, easy to remember:

• moles from mass: $n=m/M$.
• moles from particle count: $n=N/N_A$.
• combine to compute one from the other: $m=(N/N_A)\times M$.

How big is Avogadro's number?

$6.022\times 10^{23}$ is unfathomably large. If you had Avogadro's number of seconds, it would last $1.9\times 10^{16}$ years — about 1.4 million times longer than the age of the universe. One mole of marbles would cover the entire Earth to a depth of 80 km.

Yet 1 mole of water is just 18 g — about a tablespoonful. Atoms really are that tiny.

7. Worked example walkthrough

Question: How many molecules are there in 36 g of water?

Step-by-step solution:

Step 1 — Find molar mass. $M({\text{H}}_2\text{O})=2(1)+16=18\text{g/mol}$

Step 2 — Find moles. $n=\frac{m}{M}=\frac{36}{18}=2\text{mol}$

Step 3 — Number of molecules. $N=n\times N_A=2\times 6.022\times 10^{23}=1.2044\times 10^{24}\text{molecules}$

Answer: $1.2\times 10^{24}$ molecules.

8. Closing thought

What this chapter did, beautifully:

• Took the invisible atom and made it weighable (via relative atomic mass).
• Took the uncountable atom and made it countable (via the mole).
• Took gross mass measurements in a lab and translated them into particle counts at atomic scale.

Every later chapter — chemical reactions, stoichiometry, gas laws, thermodynamics, electrochemistry — depends on this translation. Master the formula $n=m/M=N/N_A$ and you have the working language of chemistry.