Atomic Structure and Isotopes

Learning Objectives (OCR Spec 2.1.1)

By the end of this lesson you should be able to:

• Describe the structure of an atom in terms of protons, neutrons and electrons
• Define atomic number (Z) and mass number (A)
• Explain what isotopes are and why they have identical chemical properties
• Determine the number of protons, neutrons and electrons in atoms and ions from isotope notation
• Recall the relative charges and relative masses of sub-atomic particles
• Appreciate the historical development of modern atomic theory

A Brief History of the Atom

The idea that matter is made of indivisible particles traces back to the ancient Greek philosopher Democritus (c. 460 BCE), who coined the word atomos meaning "uncuttable". Scientific atomic theory, however, did not emerge until John Dalton's work in 1808, which proposed that:

• Elements are composed of tiny, indivisible particles called atoms
• Atoms of a given element are identical
• Compounds form when atoms combine in simple whole-number ratios
• Chemical reactions rearrange atoms without creating or destroying them

Dalton's model was refined dramatically over the next century:

Year	Scientist	Contribution
1897	J.J. Thomson	Discovered the electron (cathode-ray experiments); "plum pudding" model
1909	Rutherford, Geiger, Marsden	Gold-foil experiment revealed a small, dense, positive nucleus
1913	Niels Bohr	Electrons in fixed energy levels (shells)
1932	James Chadwick	Discovered the neutron

Rutherford's famous gold-foil experiment, in which most alpha particles passed straight through a thin gold foil but a few bounced back at large angles, was described by Rutherford himself as "almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you". It conclusively disproved the plum-pudding model and established the nuclear atom.

The Sub-Atomic Particles

Atoms are the fundamental building blocks of matter. Despite being unimaginably small (around 10⁻¹⁰ m in diameter), they are themselves composed of three sub-atomic particles: protons, neutrons and electrons.

Particle	Symbol	Relative Mass	Relative Charge	Location
Proton	p	1	+1	Nucleus
Neutron	n	1	0	Nucleus
Electron	e⁻	1/1836 (≈ 0)	−1	Orbitals around nucleus

The actual masses and charges are:

Particle	Mass (kg)	Charge (C)
Proton	1.673 × 10⁻²⁷	+1.602 × 10⁻¹⁹
Neutron	1.675 × 10⁻²⁷	0
Electron	9.109 × 10⁻³¹	−1.602 × 10⁻¹⁹

The nucleus contains the protons and neutrons (collectively called nucleons) and holds almost all of the mass of the atom. The electrons occupy the much larger volume of space around the nucleus, arranged in orbitals.

Key Fact: If an atom were the size of a football stadium, the nucleus would be roughly the size of a pea at the centre circle. Atoms are overwhelmingly empty space.

The density of the nucleus is astonishing — around 2 × 10¹⁷ kg m⁻³. One teaspoon of nuclear matter would weigh around a trillion tonnes. This is the same density found in neutron stars.

Atomic Number and Mass Number

Two numbers fully describe any particular nuclide:

• Atomic number (Z) — the number of protons in the nucleus. This defines the element. Every carbon atom has Z = 6; anything with Z = 6 is carbon by definition.
• Mass number (A) — the total number of protons plus neutrons in the nucleus.

The standard isotope notation is:

$^{A}_{Z}\text{X}$ZA​X

For example, ²³Na (sodium-23) has:

• Z = 11 (11 protons)
• A = 23 (23 nucleons in total)
• Number of neutrons = A − Z = 23 − 11 = 12
• Number of electrons = 11 (in a neutral atom, electrons = protons)

flowchart LR
    A[Atom] --> B[Nucleus]
    A --> C[Electron cloud]
    B --> D[Protons Z]
    B --> E[Neutrons A - Z]
    C --> F[Electrons in orbitals]

Practice Exercise

For each species, work out the number of protons, neutrons and electrons:

Species	Protons	Neutrons	Electrons
¹⁶O	8	8	8
²⁰Ne	10	10	10
³⁹K	19	20	19
⁵⁶Fe	26	30	26
²³⁸U	92	146	92

Ions — Atoms with Extra or Missing Electrons

When an atom gains or loses electrons it becomes an ion. Crucially, the number of protons does not change — only the electron count changes. Chemical reactions involve electrons only, never protons.

Worked Example 1: Cations

How many protons, neutrons and electrons are in ²⁷Al³⁺?

• Z for aluminium = 13, so 13 protons
• Neutrons = 27 − 13 = 14
• The 3+ charge means 3 electrons have been lost: 13 − 3 = 10 electrons

Note that Al³⁺ is isoelectronic with neon (both have 10 electrons and the configuration 1s² 2s² 2p⁶). This is why Al prefers to form 3+ ions: doing so gives it a stable noble-gas configuration.

Worked Example 2: Anions

How many protons, neutrons and electrons are in ³²S²⁻?

• Z for sulfur = 16, so 16 protons
• Neutrons = 32 − 16 = 16
• The 2− charge means 2 electrons have been gained: 16 + 2 = 18 electrons

Again S²⁻ is isoelectronic with argon — the next noble gas after sulfur.

Worked Example 3: Transition Metal Cation

How many protons, neutrons and electrons are in ⁵⁶Fe³⁺?

• Z for iron = 26, so 26 protons
• Neutrons = 56 − 26 = 30
• 26 − 3 = 23 electrons

Note Fe³⁺ is not isoelectronic with a noble gas — transition metals achieve stable half-filled or full d-subshell configurations instead.

Isotopes

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.

Because they have the same number of protons (and in neutral atoms therefore the same number of electrons), isotopes of an element have identical chemical properties. Chemistry is governed by electron arrangement, and isotopes are electronically identical.

However, isotopes have slightly different physical properties (density, rate of diffusion, boiling point) because they have different masses.

Common Examples

Element	Isotopes	Protons	Neutrons	Natural Abundance
Hydrogen	¹H (protium)	1	0	99.985%
Hydrogen	²H (deuterium, D)	1	1	0.015%
Hydrogen	³H (tritium, T)	1	2	Trace (radioactive)
Carbon	¹²C	6	6	98.93%
Carbon	¹³C	6	7	1.07%
Carbon	¹⁴C	6	8	Trace (radioactive, used for dating)
Chlorine	³⁵Cl	17	18	75.78%
Chlorine	³⁷Cl	17	20	24.22%
Uranium	²³⁵U	92	143	0.72%
Uranium	²³⁸U	92	146	99.27%

Why Chemical Properties are Identical

Consider ³⁵Cl and ³⁷Cl. Both have 17 protons and 17 electrons arranged as 2,8,7. When chlorine reacts with sodium, both isotopes form Cl⁻ ions just as readily because the reaction involves the outer electron only. The extra two neutrons in ³⁷Cl make no difference to bonding behaviour.

Isotope Effects on Physical Properties

Although chemical behaviour is identical, physical properties differ slightly:

• Density: heavier isotopes give denser solids/liquids. D₂O ("heavy water") has density 1.11 g cm⁻³ vs H₂O at 1.00 g cm⁻³
• Rate of diffusion: lighter isotopes diffuse faster. This fact was exploited during the Manhattan Project to separate ²³⁵U from ²³⁸U by gaseous diffusion of UF₆
• Boiling points: D₂O boils at 101.4 °C vs H₂O at 100.0 °C
• Bond strength: bonds involving D are slightly stronger than those involving H — the kinetic isotope effect

Radioactive Isotopes and Their Uses

Some isotopes are unstable and decay, emitting radiation. These radioisotopes have many applications:

• ¹⁴C — radiocarbon dating of organic remains (half-life 5730 years)
• ⁶⁰Co — medical radiotherapy for cancer treatment
• ¹³¹I — diagnosis and treatment of thyroid disorders
• ²³⁵U — nuclear power and weapons
• ³²P — biological tracer in DNA/protein research

Exam Tips

• Always subtract Z from A to get neutron count, never subtract electrons
• For ions, adjust the electron count only — protons and neutrons never change
• If asked "why do isotopes have the same chemical properties" the mark scheme wants: same number of electrons / same electron configuration
• Mass number A must be a whole number (you cannot have half a neutron)
• If asked to define "isotopes" you must say: atoms of the same element with the same number of protons but different numbers of neutrons. Saying "different mass" alone is insufficient

Common Exam Mistakes

1. Confusing Z and A — candidates sometimes subtract Z from A and call the answer "electrons". The answer is neutrons.
2. Changing protons for ions — protons are fixed for an element. Na⁺ still has 11 protons, not 10.
3. Saying isotopes have different electrons — they do not. Neutral isotopes of the same element have identical electron counts.
4. Forgetting the sign of charge — a 3+ ion has lost 3 electrons; a 2− ion has gained 2.
5. Writing fractional mass numbers — mass number is always an integer. Only the relative atomic mass (see next lesson) can be a decimal.
6. Confusing isotopes with ions — isotopes differ in neutrons; ions differ in electrons. Both are distinct from "isomers" (different compounds with the same formula).
7. Giving vague definitions — phrases like "atoms with different mass" are not precise enough. Always specify "same protons, different neutrons".

Summary

• Atoms consist of a dense nucleus of protons and neutrons surrounded by electrons
• Z (atomic number) = number of protons and defines the element
• A (mass number) = protons + neutrons
• Isotopes differ only in neutron number; they have identical chemical properties because they have the same electron configuration
• Ions differ from neutral atoms in electron number only
• Atomic theory has developed over two centuries from Dalton to Chadwick and beyond
• Radioisotopes have important medical, industrial and scientific applications