Atomic Structure — Protons, Neutrons and Electrons

This lesson covers the sub-atomic particles, atomic number, mass number and isotopes as required by AQA GCSE Combined Science Trilogy (8464, Chemistry 4.1.1). You must be able to determine the number of each sub-atomic particle from the periodic table, use atomic notation and calculate relative atomic mass.

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Sub-Atomic Particles

Atoms are made of three types of sub-atomic particle:

Particle	Relative Charge	Relative Mass	Location
Proton	+1	1	Nucleus
Neutron	0	1	Nucleus
Electron	−1	Very small ($\approx$≈ 1/1836, negligible)	Energy levels (shells) around the nucleus

Key points:

• Protons and neutrons have approximately the same mass (relative mass of 1).
• Electrons have negligible mass compared to protons and neutrons.
• In a neutral atom, the number of protons equals the number of electrons — the charges balance to zero.
• The mass of an atom is concentrated almost entirely in the nucleus.

Exam Tip: You can ignore the mass of electrons when calculating atomic mass, but you must never ignore their charge — electrons determine the charge of ions.

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Atomic Number and Mass Number

Atomic Notation

The standard way to represent an atom uses atomic notation:

$^{A}_{Z}\text{X}$$

Where:

• A = mass number (top) = total number of protons + neutrons
• Z = atomic number (bottom) = number of protons
• X = chemical symbol of the element

Calculating Sub-Atomic Particles

Quantity	How to Find It
Number of protons	= atomic number (Z)
Number of electrons	= atomic number (in a neutral atom)
Number of neutrons	= mass number − atomic number (A − Z)

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Worked Examples

Example 1: Sodium

$^{23}_{11}\text{Na}$$

• Atomic number (Z) = 11, so 11 protons and 11 electrons
• Mass number (A) = 23
• Neutrons = 23 − 11 = 12 neutrons

Example 2: Oxygen

$^{16}_{8}\text{O}$$

• Protons = 8, Electrons = 8
• Neutrons = 16 − 8 = 8 neutrons

Example 3: Iron

$^{56}_{26}\text{Fe}$$

• Protons = 26, Electrons = 26
• Neutrons = 56 − 26 = 30 neutrons

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Atomic Size

• The radius of an atom is approximately 1 × 10⁻¹⁰ m (0.1 nm).
• The radius of the nucleus is approximately 1 × 10⁻¹⁴ m — about 10,000 times smaller than the atom.
• Most of the atom is empty space.

Exam Tip: If the atom were the size of a football stadium, the nucleus would be about the size of a marble on the centre spot.

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Isotopes

Isotopes are atoms of the same element with the same number of protons but a different number of neutrons. This means they have the same atomic number but different mass numbers.

Isotope	Protons	Neutrons	Mass Number
^{12}_{6}\text{C}$ (Carbon-12)	6	6	12
^{13}_{6}\text{C}$ (Carbon-13)	6	7	13
^{14}_{6}\text{C}$ (Carbon-14)	6	8	14
^{35}_{17}\text{Cl}$ (Chlorine-35)	17	18	35
^{37}_{17}\text{Cl}$ (Chlorine-37)	17	20	37

Properties of Isotopes

• Same chemical properties — because they have the same number of electrons in the same configuration, and chemical properties depend on electron arrangement.
• Different physical properties — because they have different masses (different numbers of neutrons).
• Some isotopes are radioactive (radioisotopes) — their unstable nuclei decay, emitting radiation.

Exam Tip (AQA 8464): "Explain why isotopes of the same element have identical chemical properties" is a very common question. Answer: they have the same number of electrons in the same arrangement, and chemical properties are determined by electron configuration.

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Relative Atomic Mass ($A_r$Ar​)

Because most elements exist as a mixture of isotopes, the relative atomic mass ($A_r$Ar​) is a weighted average of the masses of all naturally occurring isotopes:

$A_r = \frac{\sum (\text{isotope mass} \times \text{percentage abundance})}{100}$Ar​=100∑(isotope mass×percentage abundance)​

Worked Example: Chlorine

Chlorine has two isotopes:

• ^{35}\text{Cl}$: 75% abundance
• ^{37}\text{Cl}$: 25% abundance

$A_r = \frac{(35 \times 75) + (37 \times 25)}{100} = \frac{2625 + 925}{100} = \frac{3550}{100} = 35.5$Ar​=100(35×75)+(37×25)​=1002625+925​=1003550​=35.5

This is why chlorine's $A_r$Ar​ on the periodic table is 35.5, not a whole number.

Worked Example: Boron

Boron has two isotopes:

• ^{10}\text{B}$: 20% abundance
• ^{11}\text{B}$: 80% abundance

$A_r = \frac{(10 \times 20) + (11 \times 80)}{100} = \frac{200 + 880}{100} = \frac{1080}{100} = 10.8$Ar​=100(10×20)+(11×80)​=100200+880​=1001080​=10.8

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Atomic Structure Diagram

graph TD
    A["Atom"] --> B["Nucleus<br/>(small, dense, positive)"]
    A --> C["Electron Shells<br/>(surround the nucleus)"]
    B --> D["Protons<br/>Charge: +1 | Mass: 1"]
    B --> E["Neutrons<br/>Charge: 0 | Mass: 1"]
    C --> F["Electrons<br/>Charge: −1 | Mass: negligible"]
    A --> G["Atomic Number = Protons"]
    A --> H["Mass Number = Protons + Neutrons"]

    style A fill:#2c3e50,color:#fff
    style B fill:#e74c3c,color:#fff
    style C fill:#3498db,color:#fff
    style D fill:#e74c3c,color:#fff
    style E fill:#95a5a6,color:#fff
    style F fill:#3498db,color:#fff

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Common Mistakes

Mistake	Correction
Saying the nucleus contains electrons	Electrons are in shells around the nucleus
Confusing atomic number and mass number	Atomic number = protons; mass number = protons + neutrons
Saying isotopes have different numbers of protons	Isotopes have the same number of protons but different neutrons
Forgetting that neutrons have no charge	Neutrons are neutral — this is a frequently tested point
Using $A_r$Ar​ values as whole numbers from the periodic table	$A_r$Ar​ is a weighted average and may not be a whole number

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Summary