Class 9 Science Chapter 4: Structure of the Atom

Structure of the Atom is Chapter 4 of NCERT Class 9 Science. It describes the particles inside an atom, the atomic models of Thomson, Rutherford and Bohr, how electrons are arranged in shells, and the ideas of valency, atomic number, mass number, isotopes and isobars. Watch the video lesson below, then revise with these notes and solved examples.

What you will learn

• The discovery and properties of electrons, protons and neutrons
• Thomson's model, Rutherford's alpha-particle scattering experiment and Bohr's model of the atom
• The Bohr-Bury scheme for distributing electrons in shells
• How to find valency from electronic configuration
• Atomic number, mass number, isotopes and isobars, with average atomic mass

Video Lessons

Structure of the Atom - Part 1

Key Concepts and Formulas

Sub-atomic Particles

Dalton thought atoms could not be divided, but experiments with electricity showed that atoms contain charged particles.

• Electron (e⁻): discovered by J.J. Thomson; charge −1, mass about 1/2000 of a proton (taken as negligible)
• Proton (p⁺): linked to the canal rays studied by E. Goldstein; charge +1, mass about 1 u
• Neutron (n): discovered by J. Chadwick in 1932; no charge, mass almost equal to a proton; present in all atoms except ordinary hydrogen (protium)

Thomson's Model of the Atom

Thomson pictured the atom like a watermelon: a sphere of positive charge with electrons embedded in it like seeds.

• The positive and negative charges are equal in magnitude, so the atom is electrically neutral
• It could not explain the results of Rutherford's scattering experiment, so it was rejected

Rutherford's Scattering Experiment

Rutherford directed fast-moving alpha particles (helium ions with charge +2 and mass 4 u) at a very thin gold foil.

• Most alpha particles passed straight through, so most of the space inside an atom is empty
• A few were deflected by small angles, so the positive charge occupies very little space
• Very few (about 1 in 12,000) bounced back, so the positive charge and nearly all the mass are concentrated in a tiny nucleus
• The radius of the nucleus is about 10⁵ times smaller than the radius of the atom
• Drawback: an electron revolving in a circle would continuously radiate energy and spiral into the nucleus, so the atom would be unstable, which is not what we observe

Bohr's Model of the Atom

Neils Bohr removed this drawback by suggesting that electrons can move only in certain fixed orbits.

• Electrons revolve in special discrete orbits called shells or energy levels
• While moving in these orbits, electrons do not radiate energy
• Shells are named K, L, M, N or numbered n = 1, 2, 3, 4 starting from the nucleus

Distribution of Electrons in Shells

The Bohr-Bury scheme gives the rules for filling electrons in the first 18 elements.

Maximum electrons in shell n = 2n²
K (n = 1) = 2
L (n = 2) = 8
M (n = 3) = 18
N (n = 4) = 32
Outermost shell: maximum 8 electrons

• Electrons fill shells step by step; an outer shell starts filling only after the inner one is full

Valency

Electrons in the outermost shell are called valence electrons. Atoms react to achieve a stable octet (8 electrons in the outermost shell, or 2 for helium).

1 to 4 valence electrons:
  valency = number of valence electrons
5 to 8 valence electrons:
  valency = 8 − valence electrons

• Noble gases (He, Ne, Ar) have a complete outer shell, so their valency is 0

Atomic Number, Mass Number, Isotopes and Isobars

An atom is written with the mass number on top and the atomic number at the bottom, for example ²³₁₁Na.

Atomic number Z = number of protons
Mass number A = protons + neutrons
Neutrons = A − Z
Average atomic mass =
  Σ (isotope mass × fractional abundance)

• In a neutral atom, the number of electrons equals Z; protons and neutrons together are called nucleons
• Isotopes: atoms of the same element with the same Z but different A, such as ¹₁H, ²₁H, ³₁H or ³⁵₁₇Cl and ³⁷₁₇Cl; they have similar chemical properties but different physical properties
• Uses of isotopes: an isotope of uranium as nuclear reactor fuel, of cobalt for treating cancer, and of iodine for treating goitre
• Isobars: atoms of different elements with the same A but different Z, such as ⁴⁰₁₈Ar and ⁴⁰₂₀Ca
• The average atomic mass of chlorine works out to 35.5 u

Solved Examples

Example 1: An atom has Z = 15 and A = 31. Find its protons, electrons, neutrons, configuration and valency.

Protons = Electrons = 15
Neutrons = 31 − 15 = 16
Configuration = 2, 8, 5
Valency = 8 − 5 = 3
∴ Valency 3 (the element is phosphorus)

Example 2: Write the electronic configuration of aluminium (Z = 13) and find its valency.

K = 2, L = 8, M = 3
Valence electrons = 3
∴ Configuration 2, 8, 3; valency 3

Example 3: How many electrons can the M shell hold, and why does sodium have only 1 electron there?

Maximum in M = 2n² = 2 × 3² = 18
Sodium has 11 electrons: K = 2, L = 8
Left for M = 11 − 2 − 8 = 1
∴ M holds up to 18, but sodium has only 1 left

Example 4: An element has two isotopes of mass 63 u (69%) and 65 u (31%). Find its average atomic mass.

= 63 × 0.69 + 65 × 0.31
= 43.47 + 20.15
∴ Average atomic mass = 63.62 u

Example 5: An element Y has an average atomic mass of 10.8 u and isotopes ¹⁰Y and ¹¹Y. Find their percentages.

Let fraction of ¹⁰Y = x
10x + 11(1 − x) = 10.8
11 − x = 10.8
x = 0.2
∴ ¹⁰Y = 20%, ¹¹Y = 80%

Example 6: Find the number of neutrons in ³⁹₁₉K and its electronic configuration.

Neutrons = 39 − 19 = 20
19 electrons: K = 2, L = 8, M = 8, N = 1
∴ 20 neutrons; configuration 2, 8, 8, 1

Example 7: Are ²⁴₁₂Mg and ²⁴₁₁Na isotopes or isobars?

Answer: They have the same mass number (24) but different atomic numbers (12 and 11), so they are isobars.

Common Mistakes to Avoid

• Do not confuse mass number with atomic mass; mass number is always a whole number, while average atomic mass (like 35.5 u) need not be.
• Remember that potassium (Z = 19) is 2, 8, 8, 1 and not 2, 8, 9, because the outermost shell cannot hold more than 8 electrons.
• In Rutherford's experiment, link each observation to its conclusion; examiners give marks for both parts.
• Isotopes have the same number of protons, not the same number of neutrons.

Quick Revision

Electron −1, proton +1, neutron 0
Maximum electrons in a shell = 2n²
Neutrons = A − Z

• Neutron mass is almost equal to proton mass
• Rutherford: tiny, dense, positively charged nucleus; most of the atom is empty
• Bohr: electrons in fixed shells do not radiate energy
• The outermost shell holds at most 8 electrons
• Isotopes: same Z; isobars: same A

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