# Atomic Structure

> Edexcel International GCSE Chemistry · 4CH1 2017
> Source: https://www.owlsprep.com/study/edexcel-igcse-chemistry-s1-atomic-structure/

This guide covers all core atomic structure content for Edexcel IGCSE Chemistry 4CH1 specification points 1.14–1.17, including subatomic particles, isotopes, and Ar calculations.

**Prerequisites:** [Basic knowledge of elements and the periodic table](https://www.owlsprep.com/study/edexcel-igcse-chemistry-s1-elements-periodic-table/)

## Learning objectives

- Define the terms atom and molecule
- Describe the structure of the atom, including relative masses and charges of protons, neutrons and electrons
- Define atomic number, mass number, isotopes and relative atomic mass (Ar)
- Calculate relative atomic mass from given isotopic abundance data

## Key Definitions: Atoms and Molecules

All matter is made of particles, with atoms and molecules being the two fundamental units you need to distinguish for this specification.

**Atom and Molecule** — An atom is the smallest unit of an element that can take part in a chemical reaction. A molecule is two or more atoms chemically bonded together, either of the same element (e.g. O₂) or different elements (e.g. H₂O).

**Worked example:** Classify the following as atoms or molecules: (a) Ne, (b) CO₂, (c) N₂

1. 1. Recall that an atom is a single neutral particle of an element, while a molecule has 2+ bonded atoms.
2. 2. (a) Ne is a single, unbonded atom of neon, so it is an atom.
3. 3. (b) CO₂ has 1 carbon and 2 oxygen atoms bonded, so it is a molecule.
4. 4. (c) N₂ has 2 nitrogen atoms bonded, so it is a molecule.

> **Exam tip:** When given particle diagrams in exam questions, count the number of bonded circles to quickly distinguish atoms (1 circle) from molecules (2+ circles).

## Subatomic Particle Structure

All atoms consist of a dense, positively charged central nucleus surrounded by negatively charged electrons orbiting in energy shells. The nucleus contains two types of particles: protons and neutrons.

| Subatomic Particle | Relative Charge | Relative Mass | Location |
| --- | --- | --- | --- |
| Proton | +1 | 1 | Nucleus |
| Neutron | 0 | 1 | Nucleus |
| Electron | -1 | ~1/1840 (negligible) | Electron shells outside nucleus |

> **warning**
>
> You must memorize the relative charge and mass of all three subatomic particles, as these values are not provided on the periodic table in your exam.

**Worked example:** State the number of subatomic particles in a neutral atom of lithium, which has atomic number 3 and mass number 7.

1. 1. Atomic number = number of protons: 3 protons.
2. 2. Neutral atoms have equal numbers of protons and electrons: 3 electrons.
3. 3. Number of neutrons = mass number - atomic number = 7 - 3 = 4 neutrons.

> **Exam tip:** Neutral atoms have no overall charge because the positive charge from protons exactly cancels the negative charge from electrons.

## Atomic Number, Mass Number and Isotopes

Two key numbers define the composition of an atom: atomic number (given on the periodic table for every element) and mass number (varies between isotopes of the same element).

**Isotopes** — Atoms of the same element with the same atomic number (same number of protons) but different mass numbers (different numbers of neutrons).

Isotopes have identical chemical properties because they have the same number of electrons, but different physical properties like mass, density or radioactivity.

**Worked example:** Oxygen has three isotopes: oxygen-16, oxygen-17 and oxygen-18. All have atomic number 8. State the number of neutrons in each isotope.

1. 1. Neutron count = mass number - atomic number for each isotope.
2. 2. Oxygen-16: 16 - 8 = 8 neutrons.
3. 3. Oxygen-17: 17 - 8 = 9 neutrons.
4. 4. Oxygen-18: 18 - 8 = 10 neutrons.

> **Exam tip:** Questions often ask why isotopes have the same chemical properties: always reference the equal number of electrons in your answer to get full marks.

## Calculating Relative Atomic Mass (Ar)

The relative atomic mass (Ar) of an element is a weighted average of the masses of its naturally occurring isotopes, accounting for how abundant each isotope is.

$$A_r = \frac{\text{Sum of (isotope mass} \times \text{percentage abundance)}}{100}$$

> **tip**
>
> Isotopic masses and percentage abundances will always be provided in the question, you do not need to memorize these values. Round your final answer to match the significant figures of the given data.

**Worked example:** Bromine has two isotopes: 50.7% of bromine atoms are bromine-79, 49.3% are bromine-81. Calculate the Ar of bromine to 3 significant figures.

1. 1. Calculate the weighted mass of each isotope:
2. $$(79 \times 50.7) = 4005.3; (81 \times 49.3) = 3993.3$$
3. 2. Add the weighted masses together:
4. $$4005.3 + 3993.3 = 7998.6$$
5. 3. Divide by 100:
6. $$7998.6 \times \frac{1}{100} = 79.986$$
7. 4. Round to 3 significant figures: Ar = 80.0

> **Exam tip:** Never calculate a simple average of isotope masses, as this does not account for differences in abundance and will give you an incorrect answer.

*Calculator:* allowed

## Common pitfalls

- **Wrong:** Counting electrons when calculating mass number
  - Why it fails: Electron mass is negligible, ~1/1840 of a proton mass, so it is not included in mass number calculations
  - Correct: Only add the number of protons and neutrons to calculate mass number
- **Wrong:** Assuming the number of neutrons equals the number of protons
  - Why it fails: This is only true for some light elements, and isotopes of the same element have different neutron counts
  - Correct: Calculate neutron count as mass number minus atomic number for each atom/isotope
- **Wrong:** Calculating a simple mean of isotope masses for Ar
  - Why it fails: Isotopes are not equally abundant in nature, so their masses must be weighted by their percentage abundance
  - Correct: Use the weighted mean formula: Ar = sum (isotope mass × % abundance) ÷ 100
- **Wrong:** Stating isotopes have different chemical properties
  - Why it fails: Chemical properties are controlled by the number of electrons, which is identical for all isotopes of the same element
  - Correct: Remember isotopes only differ in physical properties like mass, density or radioactivity
- **Wrong:** Using relative charge values for mass number calculations
  - Why it fails: Charge and mass are independent properties of subatomic particles
  - Correct: Only use relative mass values for any calculation of total atomic mass

## Cheatsheet

| Concept | Key Fact / Formula |
| --- | --- |
| Atom | Smallest unit of an element; nucleus + orbiting electrons |
| Molecule | 2+ atoms chemically bonded together |
| Subatomic Particles | Proton: +1 charge, mass 1 (nucleus); Neutron: 0 charge, mass 1 (nucleus); Electron: -1 charge, mass ~1/1840 (shells) |
| Atomic number (Z) | Number of protons = number of electrons (neutral atoms) |
| Mass number (A) | Number of protons + number of neutrons |
| Isotopes | Same atomic number, different mass number; same chemical properties |
| Relative Atomic Mass (Ar) | Weighted mean mass; $A_r = \frac{\text{Sum of (isotope mass} \times \text{% abundance)}}{100}$ |

## What's next

Now that you have mastered core atomic structure, you are ready to move on to electron arrangement in shells, which explains the structure of the periodic table and how elements form bonds. You will also use your understanding of relative atomic mass to calculate relative formula mass and mole quantities in upcoming quantitative chemistry topics. Practice Ar calculation questions regularly to avoid common weighting errors in your exam.

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