Atoms and Molecules
CBSE Class 9 · Science · Notes, formulas and practice questions
Revision notes and practice questions for the CBSE Class 9 chapter 'Atoms and Molecules', covering laws of chemical combination, Dalton's atomic theory, atomic and molecular masses, ions, chemical formulae, the mole concept, and mass–mole–particle calculations.
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This chapter builds an understanding of matter from the smallest building blocks. It begins with two empirical laws – the law of conservation of mass and the law of constant proportions – which arise from careful measurements in chemical reactions. Dalton's atomic theory provides a particulate explanation for these laws, introducing atoms as indestructible, indivisible particles that combine in fixed ratios to form compounds.
Atoms of different elements have different masses and combining capacities, known as valency. Using valency, we can write both molecular formulae for covalent compounds and formula units for ionic compounds, including those containing polyatomic ions such as sulphate and ammonium. Balancing charges correctly is the key step in deducing a correct chemical formula.
Since atoms are too light to weigh individually, relative masses are used. The atomic mass unit (u) is defined on the carbon-12 scale, and the molecular mass of a compound is the sum of the atomic masses of its constituent atoms. For ionic compounds the same sum gives the formula unit mass. These values are needed for all later mole calculations.
The mole concept connects the microscopic world of atoms to laboratory-scale masses. One mole contains Avogadro's constant of particles, and the mass of one mole equals the molar mass in grams. The relations n = m/M and N = n × N₀ allow conversion between mass, moles and number of particles, enabling quantitative problems throughout chemistry.
Key terms
- Atom
- The smallest particle of an element that can take part in a chemical reaction. Atoms are extremely small and may or may not exist independently; they combine with other atoms to form molecules or ions.
- Molecule
- The smallest particle of an element or compound that can exist independently and shows all the properties of that substance. A molecule is formed when two or more atoms combine chemically in a fixed ratio.
- Atomic mass unit (u)
- A unit used to express the relative masses of atoms and molecules. 1 u is defined as 1/12 the mass of one carbon-12 atom, which is approximately 1.66 × 10⁻²⁴ g.
- Molecular mass
- The sum of the atomic masses of all the atoms present in a molecule, expressed in atomic mass units (u). It is calculated by multiplying each element's atomic mass by the number of its atoms in the formula and adding them.
- Ion
- An atom or a group of atoms that carries a net electric charge because it has lost or gained electrons. Positively charged ions are called cations, and negatively charged ions are called anions.
- Polyatomic ion
- A group of atoms that are covalently bonded together and carry an overall electric charge, acting as a single unit in chemical reactions. Examples include ammonium (NH₄⁺) and sulphate (SO₄²⁻).
- Mole
- The SI unit for the amount of substance. One mole contains exactly 6.022 × 10²³ particles – atoms, molecules, ions, or electrons – a number known as Avogadro's constant.
- Molar mass
- The mass in grams of one mole of a substance. It is numerically equal to the relative atomic mass, molecular mass, or formula unit mass, so for example the molar mass of O₂ is 32 g/mol.
Formula sheet
| What | Formula | Notes |
|---|---|---|
| Number of moles from mass | n = m/M | n is the amount of substance in moles, m is the mass of the sample in grams, and M is the molar mass in grams per mole. This relation is used when a known mass of a pure substance is converted to moles. |
| Number of particles from moles | N = n × N₀ | N is the total number of particles, n is the number of moles, and N₀ is Avogadro's constant equal to 6.022 × 10²³ particles per mole. The particles may be atoms, molecules or formula units depending on the substance. |
Practice questions with answers
1. State the law of conservation of mass.
During a chemical reaction, mass can neither be created nor destroyed. The total mass of all reactants equals the total mass of all products, because atoms are merely rearranged and no atoms are lost or gained.
2. State the law of constant proportions.
A pure chemical compound always contains the same elements combined together in the same proportion by mass, regardless of how the compound is prepared or obtained. For example, water from any source has hydrogen and oxygen in a fixed mass ratio of 1:8.
3. Calculate the number of moles in 11 g of carbon dioxide. (Atomic masses: C = 12 u, O = 16 u)
Molar mass of CO₂ = 12 + 2 × 16 = 44 g/mol. Moles of CO₂ = given mass / molar mass = 11 g / 44 g/mol = 0.25 mol. So 11 g of CO₂ corresponds to 0.25 mole.
4. How many molecules are present in 18 g of water? (H = 1 u, O = 16 u)
Molar mass of H₂O = 2 × 1 + 16 = 18 g/mol. Moles of H₂O = 18 g / 18 g/mol = 1 mol. The number of molecules is 1 mol × 6.022 × 10²³ molecules/mol = 6.022 × 10²³ molecules.
5. What is the mass of 0.25 mol of oxygen gas (O₂)?
Molar mass of O₂ = 2 × 16 = 32 g/mol. Mass = number of moles × molar mass = 0.25 mol × 32 g/mol = 8 g. Therefore, 0.25 mol of oxygen gas has a mass of 8 g.
6. Write the chemical formula of aluminium sulphate.
Aluminium has valency 3 (Al³⁺) and sulphate is a polyatomic ion with valency 2 (SO₄²⁻). To balance charges, two Al³⁺ ions combine with three SO₄²⁻ ions, so the formula is Al₂(SO₄)₃.
7. What is the molecular mass of calcium carbonate (CaCO₃)? (Ca = 40 u, C = 12 u, O = 16 u)
Molecular mass of CaCO₃ = 40 + 12 + 3 × 16 = 100 u. Thus a single formula unit of calcium carbonate has a relative mass of 100 u.
8. Calculate the number of atoms in 4 g of helium. (He = 4 u)
Molar mass of helium = 4 g/mol, so 4 g of helium corresponds to 1 mole. Helium exists as monatomic atoms, and 1 mole of helium contains 6.022 × 10²³ atoms.
9. Define an ion. Give one example of a cation and one example of an anion.
An ion is an atom or group of atoms that carries a net electric charge. A cation is a positively charged ion, such as Na⁺, while an anion is negatively charged, such as Cl⁻. Polyatomic ions also exist, for example NH₄⁺ and SO₄²⁻.
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