What this chapter is about
This chapter introduces the idea that all matter is made of tiny particles called atoms, and that atoms of different elements combine to form molecules. You will learn how scientists worked out that matter is not continuous but granular — built from incredibly small units that cannot be seen with the naked eye. The chapter connects what you already know about the particle nature of matter (from earlier study of states and changes of state) to the actual building blocks: atoms.
By the end, you should be able to describe what atoms and molecules are, explain how elements differ from compounds at the atomic level, write symbols for elements and simple formulas for compounds, and calculate the relative masses of atoms and molecules. You will also meet important laws (law of conservation of mass, law of constant proportions) that guided early chemists toward the atomic theory.
Understanding atoms is essential because every topic in chemistry — reactions, acids, metals, carbon compounds — depends on knowing how atoms behave, combine and rearrange.
Key ideas
- Atoms are the smallest particles of an element that retain the chemical properties of that element; they cannot be divided further by ordinary chemical means.
- Molecules are groups of two or more atoms chemically bonded together. A molecule may contain atoms of the same element (like O₂) or different elements (like H₂O).
- Elements are pure substances made of only one kind of atom; compounds are pure substances made of two or more kinds of atoms in a fixed ratio by mass.
- The law of conservation of mass states that mass is neither created nor destroyed in a chemical reaction; total mass of reactants equals total mass of products.
- The law of constant proportions (or definite proportions) states that a given compound always contains the same elements in the same ratio by mass, regardless of source or method of preparation.
- Each element is represented by a one- or two-letter symbol (first letter capital, second lowercase). For example, sodium is Na, chlorine is Cl.
- Atomic mass is the relative mass of an atom compared with one-twelfth the mass of a carbon-12 atom, expressed in atomic mass units (u).
- Molecular mass is the sum of the atomic masses of all atoms in a molecule.
Formulas and facts to remember
1. Atomic mass unit: 1 u = 1/12 the mass of one carbon-12 atom. 2. Molecular mass = sum of atomic masses of all atoms in the formula. 3. Law of conservation of mass: Mass of reactants = Mass of products. 4. Law of constant proportions: In a compound, elements combine in a fixed mass ratio. 5. Symbol rule: First letter capital, second (if any) lowercase — e.g., Ca for calcium, Fe for iron. 6. Water (H₂O) always has hydrogen and oxygen in the mass ratio 1 : 8. 7. Carbon dioxide (CO₂) always has carbon and oxygen in the mass ratio 3 : 8. 8. Common atomic masses to recall: H = 1 u, C = 12 u, N = 14 u, O = 16 u, S = 32 u, Na = 23 u, Cl = 35.5 u.
Worked examples
### Example 1: Calculating molecular mass of water
Problem: Find the molecular mass of water, H₂O. (Atomic masses: H = 1 u, O = 16 u.)
Solution: Water has 2 hydrogen atoms and 1 oxygen atom. Mass from H = 2 × 1 u = 2 u Mass from O = 1 × 16 u = 16 u Molecular mass of H₂O = 2 u + 16 u = 18 u
### Example 2: Verifying constant proportions in carbon dioxide
Problem: Two samples of carbon dioxide weigh 22 g and 44 g. Show that carbon and oxygen are present in the same ratio in both.
Solution: Carbon dioxide has the formula CO₂. Atomic masses: C = 12 u, O = 16 u. Molecular mass = 12 + 2 × 16 = 44 u. Fraction of carbon = 12/44 = 3/11. Fraction of oxygen = 32/44 = 8/11.
For the 22 g sample: Carbon = 22 × 3/11 = 6 g; Oxygen = 22 × 8/11 = 16 g. For the 44 g sample: Carbon = 44 × 3/11 = 12 g; Oxygen = 44 × 8/11 = 32 g.
Ratio of carbon to oxygen in both = 6 : 16 = 12 : 32 = 3 : 8. This confirms the law of constant proportions.
### Example 3: Applying conservation of mass
Problem: 5 g of iron reacts completely with 2 g of sulphur to form iron sulphide. What is the mass of iron sulphide produced?
Solution: By the law of conservation of mass, total mass before reaction = total mass after. Mass of reactants = 5 g + 2 g = 7 g. Therefore, mass of iron sulphide formed = 7 g.
Common mistakes
- Writing Na as NA or na → the first letter must be capital, the second lowercase.
- Confusing atomic mass with actual mass in grams → atomic mass is a relative number in units (u), not a direct gram value.
- Adding subscripts when balancing equations instead of using coefficients → subscripts show the formula and must not be changed; only coefficients change.
- Thinking smaller samples of a compound have different element ratios → the law of constant proportions means the ratio stays the same regardless of sample size.
- Believing mass can disappear if a gas escapes during a reaction → mass is still conserved; it only seems lost because the gas left the container.
Quick revision
- An atom is the smallest unit of an element; a molecule is two or more atoms bonded together.
- Symbols: one or two letters, first capital, second lowercase.
- Molecular mass = sum of atomic masses of all atoms in the formula.
- Law of conservation of mass: mass before reaction = mass after reaction.
- Law of constant proportions: a compound always has elements in a fixed mass ratio.
- Common atomic masses: H = 1 u, C = 12 u, O = 16 u, N = 14 u.