Atomic Models and Subatomic Particles

This lesson covers the development of atomic models, the properties of subatomic particles, isotopes, and the concepts of mass number, atomic number, and relative atomic mass. A solid grasp of these fundamentals is essential for everything that follows in A-Level Chemistry.

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The Development of Atomic Models

Our understanding of the atom has evolved over centuries through key experiments and theoretical breakthroughs.

Dalton's Model (1803)

John Dalton proposed that all matter is made of indivisible atoms. He suggested that atoms of the same element are identical and that chemical reactions involve the rearrangement of atoms. While groundbreaking, this model had no concept of internal structure.

Thomson's Plum Pudding Model (1897)

J.J. Thomson discovered the electron using cathode ray tubes. He measured the charge-to-mass ratio of electrons and proposed the "plum pudding" model: a sphere of positive charge with negatively charged electrons embedded within it, like plums in a pudding.

Rutherford's Nuclear Model (1911)

Ernest Rutherford directed alpha particles at a thin gold foil. Most passed straight through, but a small fraction were deflected at large angles and some bounced back. This could only be explained if:

• Most of the atom is empty space (most alpha particles pass through).
• The positive charge and most of the mass are concentrated in a tiny, dense nucleus (large-angle deflections).
• The nucleus is positively charged (alpha particles, which are positive, are repelled).

Bohr's Model (1913)

Niels Bohr refined the nuclear model by proposing that electrons orbit the nucleus in fixed energy levels (shells). Electrons can move between energy levels by absorbing or emitting specific amounts of energy. This model successfully explained the line spectrum of hydrogen.

The Quantum Mechanical Model (1920s onwards)

The modern model treats electrons as existing in orbitals — regions of space where there is a high probability of finding an electron. Electrons do not follow fixed orbits but are described by wave functions. This model underpins A-Level electron configuration.

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Subatomic Particles

Atoms are made of three subatomic particles:

Particle	Relative mass	Relative charge	Location
Proton	1	+1	Nucleus
Neutron	1	0	Nucleus
Electron	1/1836 (≈ 0.00055)	−1	Orbitals around nucleus

The actual masses are:

• Proton: 1.673 × 10⁻²⁷ kg
• Neutron: 1.675 × 10⁻²⁷ kg
• Electron: 9.109 × 10⁻³¹ kg

Key Point: For most A-Level calculations, the mass of the electron is considered negligible compared to protons and neutrons. The relative mass of an electron is approximately 1/1836 that of a proton.

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Atomic Number and Mass Number

Every element is defined by two key numbers:

• Atomic number (Z): The number of protons in the nucleus. This defines the element. All atoms of the same element have the same atomic number.
• Mass number (A): The total number of protons and neutrons in the nucleus. Also called the nucleon number.

The number of neutrons = A − Z.

In a neutral atom, the number of electrons equals the number of protons. When an atom forms an ion, it gains or loses electrons but the number of protons remains unchanged.

Notation

An atom is represented as:

ᴬ_Z X

For example, sodium-23: ²³₁₁Na has 11 protons, 12 neutrons (23 − 11), and 11 electrons in the neutral atom.

Exam Tip: If asked for the number of electrons in an ion, remember to adjust: a 2+ ion has lost 2 electrons, a 2− ion has gained 2 electrons.

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Isotopes

Isotopes are atoms of the same element (same number of protons / same atomic number) that have different numbers of neutrons (different mass numbers).

Examples of Important Isotopes

Isotope	Protons	Neutrons	Mass number	Natural abundance (%)
¹H (protium)	1	0	1	99.98
²H (deuterium)	1	1	2	0.02
¹²C	6	6	12	98.9
¹³C	6	7	13	1.1
³⁵Cl	17	18	35	75.8
³⁷Cl	17	20	37	24.2

Properties of Isotopes

Isotopes of the same element have:

• Identical chemical properties — because they have the same electron configuration and the same number of electrons.
• Different physical properties — such as density, rate of diffusion, and boiling point, because they have different masses.

Common Misconception: Students sometimes say isotopes have "different chemical properties." This is incorrect. Chemical properties depend on the electron configuration, which is the same for all isotopes of an element.

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Relative Masses

Because the actual masses of atoms are incredibly small, we use a relative scale based on carbon-12.

Key Definitions

• Relative isotopic mass: The mass of one atom of an isotope compared to one-twelfth of the mass of one atom of carbon-12.
• Relative atomic mass (Aᵣ): The weighted mean mass of an atom of an element compared to one-twelfth of the mass of one atom of carbon-12.
• Relative molecular mass (Mᵣ): The weighted mean mass of a molecule compared to one-twelfth of the mass of one atom of carbon-12.
• Relative formula mass: Used for ionic compounds — the sum of the relative atomic masses of all atoms in the formula unit.

Key Definition: Relative atomic mass is a weighted mean because it takes into account the natural abundances of each isotope, not just their masses.

Calculating Relative Atomic Mass

The formula for Aᵣ is:

Aᵣ = Σ (isotopic mass × percentage abundance) / 100

Worked Example 1: Calculating Aᵣ of Chlorine

Chlorine has two isotopes: ³⁵Cl (75.8%) and ³⁷Cl (24.2%).

Aᵣ = (35 × 75.8 + 37 × 24.2) / 100

Aᵣ = (2653.0 + 895.4) / 100

Aᵣ = 3548.4 / 100

Aᵣ = 35.5 (to 1 decimal place)

Worked Example 2: Calculating Aᵣ of Boron

Boron has two isotopes: ¹⁰B (19.9%) and ¹¹B (80.1%).

Aᵣ = (10 × 19.9 + 11 × 80.1) / 100

Aᵣ = (199.0 + 881.1) / 100

Aᵣ = 1080.1 / 100

Aᵣ = 10.8 (to 1 decimal place)

Worked Example 3: Three Isotopes — Silicon

Silicon has three isotopes: ²⁸Si (92.2%), ²⁹Si (4.7%), ³⁰Si (3.1%).

Aᵣ = (28 × 92.2 + 29 × 4.7 + 30 × 3.1) / 100

Aᵣ = (2581.6 + 136.3 + 93.0) / 100

Aᵣ = 2810.9 / 100

Aᵣ = 28.1 (to 1 decimal place)

Exam Tip: Always show your working in relative atomic mass calculations. A common error is to simply average the mass numbers without weighting by abundance. The Aᵣ of chlorine is 35.5, NOT 36 (which would be the simple average of 35 and 37).

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Ions and Their Electronic Structure

When atoms gain or lose electrons, they form ions:

• Metals lose electrons to form positive ions (cations). For example, Na → Na⁺ + e⁻.
• Non-metals gain electrons to form negative ions (anions). For example, Cl + e⁻ → Cl⁻.

Worked Example 4: Particles in Ions

How many protons, neutrons, and electrons are in ⁵⁶₂₆Fe³⁺?

• Protons = 26 (the atomic number)
• Neutrons = 56 − 26 = 30
• Electrons = 26 − 3 = 23 (lost 3 electrons to form the 3+ ion)

Worked Example 5: Identifying an Unknown Ion

An ion X²⁻ has 10 electrons and 8 protons. Identify the element and write its notation.

• Atomic number = 8, so the element is oxygen.
• Number of electrons = 8 + 2 = 10 (gained 2 electrons for the 2− charge). ✓
• Mass number is not given, but the most common isotope of oxygen is ¹⁶O, so the ion is ¹⁶₈O²⁻.

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Summary

Concept	Key facts
Atomic number (Z)	Number of protons; defines the element
Mass number (A)	Protons + neutrons
Isotopes	Same Z, different A (different neutrons)
Relative atomic mass	Weighted mean mass relative to ¹²C / 12
Neutral atom	Protons = electrons
Positive ion	Fewer electrons than protons
Negative ion	More electrons than protons

Exam Tip: In multiple-choice questions, you may be asked to identify the correct definition of isotopes. The key phrase is "atoms of the same element with different numbers of neutrons." Do not say "different mass" — be specific about neutrons.