## Learning Outcomes

- Identify the location, relative charge and relative mass of protons, neutrons and electrons
- Explain how atomic number and mass number are used to find numbers of subatomic particles
- Distinguish clearly between atoms, ions and isotopes
- Explain why isotopes have similar chemical properties but different physical properties
- Calculate numbers of protons, neutrons and electrons in atoms and ions step by step
- Calculate relative atomic mass from isotopic abundance data
- Avoid common exam errors involving mass number, relative atomic mass and isotope definitions

## AQA A-Level Chemistry (7405) Syllabus

For the AQA A-Level Chemistry (7405), you are expected to understand subatomic particles and how they are used to describe atomic structure, with a focus on the following syllabus points:

- Know that an atom consists of a nucleus containing protons and neutrons, surrounded by electrons
- Recall the relative charge and relative mass of protons, neutrons and electrons
- Identify atomic number as the number of protons in the nucleus
- Identify mass number as the total number of protons and neutrons in the nucleus
- Determine the numbers of protons, neutrons and electrons in atoms and ions
- Explain what isotopes are and how they differ from one another
- Understand that isotopes of the same element have similar chemical properties because they have the same electron arrangement
- Calculate relative atomic mass as a weighted mean using isotopic masses and abundances

## Test Your Knowledge

_Attempt these questions before reading this article. If you find some difficult or cannot remember the answers, look more closely at that area during your revision._

1. What are the relative charges and relative masses of a proton, neutron and electron?
2. How do you calculate the number of neutrons in an atom from its mass number and atomic number?
3. What is the difference between mass number and relative atomic mass?
4. Why do isotopes of the same element have similar chemical properties?
5. How many protons, neutrons and electrons are present in $^{24}_{12}\text{Mg}^{2+}$?

## Introduction

Atomic structure looks simple at first: atoms contain protons, neutrons and electrons. But in A-level Chemistry, this topic quickly becomes the foundation for much more difficult ideas, including electron configuration, ionisation energy, bonding, redox chemistry and mass spectrometry. If your definitions here are weak, errors spread into many other parts of the course.

For Paper 3 especially, this topic often appears in mixed questions where you must move quickly from definitions to calculations and then to explanation. That means you need more than facts to memorise. You need to know what each particle does, how the numbers relate to one another, and why small differences in atomic structure matter.

> **Key Term: subatomic particles**
> The particles smaller than an atom that make up atomic structure: protons, neutrons and electrons.

### The three subatomic particles and what they tell you

An atom has a central nucleus and electrons around it. The nucleus contains **protons** and **neutrons**. **Electrons** are found outside the nucleus in shells and orbitals. Nearly all the mass of an atom is concentrated in the nucleus because protons and neutrons are much more massive than electrons.

The three particles must be known precisely.

| Particle | Relative charge | Relative mass | Location        |
| -------- | --------------: | ------------: | --------------- |
| Proton   |              +1 |             1 | Nucleus         |
| Neutron  |               0 |             1 | Nucleus         |
| Electron |              -1 |        1/1836 | Outside nucleus |

A common shortcut in chemistry is to treat electron mass as negligible. This is usually fine in chemical calculations because an electron is so much lighter than a proton or neutron. However, in a definition question, you should still give the proper relative mass as **1/1836** rather than saying just "0". "Negligible" is more accurate than "zero".

> **Key Term: proton**
> A subatomic particle found in the nucleus with relative charge +1 and relative mass 1.
>
> **Key Term: neutron**
> A subatomic particle found in the nucleus with relative charge 0 and relative mass 1.
>
> **Key Term: electron**
> A subatomic particle found outside the nucleus with relative charge -1 and relative mass 1/1836.

Why do these values matter?

- The **proton number** determines which element an atom is.
- The **electron arrangement** controls chemical behaviour.
- The **neutron number** changes mass and can change stability, but does not change the identity of the element.

This is why changing electrons gives an ion, changing neutrons gives an isotope, and changing protons gives a different element entirely.

There is also an apparent problem inside the nucleus: protons are all positively charged, so they should repel one another. The nucleus stays together because of the **strong nuclear force**, which acts between nucleons over very short distances. This force is stronger than electrostatic repulsion at that scale.

> **Key Term: nucleus**
> The dense central part of the atom that contains protons and neutrons and holds almost all of the atom's mass.
>
> **Test Tip:** If a 1-mark question asks where most of the mass of the atom is found, the answer is the **nucleus**, not "protons and neutrons" on their own. Make sure your wording matches the exact point being tested.

### Atomic number, mass number, atoms and ions

Two numbers describe the nucleus of an atom.

- **Atomic number**, $Z$, is the number of protons
- **Mass number**, $A$, is the number of protons plus neutrons

So:

$$
A = \text{protons} + \text{neutrons}
$$

and therefore:

$$
\text{neutrons} = A - Z
$$

> **Key Term: atomic number**
> The number of protons in the nucleus of an atom.
>
> **Key Term: mass number**
> The total number of protons and neutrons in the nucleus of an atom.

These numbers are often written in nuclear notation, for example:

$$
^{27}_{13}\text{Al}
$$

This tells you:

- protons = 13
- neutrons = 27 - 13 = 14
- electrons = 13 if the atom is neutral

The number of electrons depends on charge, not mass number.

For a **neutral atom**:

- electrons = protons

For a **positive ion**:

- electrons = protons - charge

For a **negative ion**:

- electrons = protons + charge size

This is where many students make careless mistakes. They find protons and neutrons correctly, then forget to adjust electrons for the ion charge.

Worked example 1:

Find the numbers of protons, neutrons and electrons in $^{24}_{12}\text{Mg}^{2+}$.

Step 1: Protons = atomic number = 12

Step 2: Neutrons = mass number - atomic number

$$
24 - 12 = 12
$$

Step 3: The ion is $2+$, so it has lost 2 electrons

$$
12 - 2 = 10
$$

Answer:

- protons = 12
- neutrons = 12
- electrons = 10

Worked example 2:

Find the numbers of protons, neutrons and electrons in $^{35}_{17}\text{Cl}^{-}$.

Step 1: Protons = 17

Step 2: Neutrons

$$
35 - 17 = 18
$$

Step 3: A $1-$ ion has gained one electron

$$
17 + 1 = 18
$$

Answer:

- protons = 17
- neutrons = 18
- electrons = 18

Why does atomic number matter more than mass number for identity? Because chemical identity depends on proton number. Every atom with 8 protons is oxygen. It can have 8, 9 or 10 neutrons and still be oxygen. But if it has 9 protons, it is fluorine.

> **Key Term: ion**
> An atom or group of atoms with an overall electric charge because electrons have been gained or lost.

Common student errors in this area:

- using mass number as the number of neutrons
- forgetting that atomic number is protons, not total particles
- changing proton number when forming ions
- subtracting charge from neutrons instead of electrons
  > **Test Tip:** In particle-count questions, always write three lines in the same order: protons, neutrons, electrons. This reduces mistakes under exam pressure.

### Isotopes and why they matter

An element can exist as different isotopes.

> **Key Term: isotope**
> Atoms of the same element with the same number of protons but different numbers of neutrons.

This definition must be exact. "Same element but different masses" is not enough for full-mark definition questions, because it does not explain why. The key point is **same number of protons, different number of neutrons**.

Take chlorine as an example:

- $^{35}\text{Cl}$: 17 protons, 18 neutrons
- $^{37}\text{Cl}$: 17 protons, 20 neutrons

Both are chlorine because both have 17 protons.

Why do isotopes have similar chemical properties?

Chemical properties depend mainly on electron arrangement, especially the outer-shell electrons involved in bonding. Since isotopes of the same element have the same number of protons, neutral atoms also have the same number of electrons. So they have the same electron configuration and react in very similar ways.

Why can isotopes have different physical properties?

Physical properties can depend on mass. Since isotopes have different numbers of neutrons, they have different masses. This can affect properties such as rate of diffusion. The chemistry stays very similar, but the mass changes.

This distinction is often tested:

- **chemical properties**: controlled by electrons
- **mass differences**: caused by neutrons

Students often wrongly say isotopes react differently because they have different neutrons. For most standard chemistry questions, that is not accepted.

A useful cause-and-effect chain is:

same protons -> same electron arrangement -> similar chemical properties

different neutrons -> different mass -> possible differences in physical properties

### Relative atomic mass and simple mass spectrometry links

The mass number of a single atom is always a whole number, because it is a count of particles in the nucleus. But the value shown in the Periodic Table is usually not a whole number. That value is the **relative atomic mass**, $A_r$.

> **Key Term: relative atomic mass**
> The weighted mean mass of an atom of an element compared with 1/12 of the mass of an atom of carbon-12.

There are two parts students often miss:

- it is a **weighted mean**
- it is compared with **1/12 of carbon-12**

If you leave either part out in a definition, you risk losing marks.

Why is $A_r$ often not an integer? Because most elements are mixtures of isotopes. The Periodic Table value is an average based on both isotope masses and abundances.

General formula:

$$
A_r = \frac{\sum(\text{isotopic mass} \times \text{abundance})}{\sum(\text{abundance})}
$$

If percentages are given, divide by 100.

Worked example 3:

Chlorine has two isotopes:

- mass 35, abundance 75%
- mass 37, abundance 25%

Calculate $A_r$.

Step 1: Multiply each mass by its abundance

$$
35 \times 75 = 2625
$$

$$
37 \times 25 = 925
$$

Step 2: Add the totals

$$
2625 + 925 = 3550
$$

Step 3: Divide by 100

$$
A_r = \frac{3550}{100} = 35.5
$$

So the relative atomic mass of chlorine is **35.5**.

Worked example 4:

An element has isotopes:

- mass 24, abundance 79
- mass 25, abundance 10
- mass 26, abundance 11

Calculate $A_r$.

Step 1: Multiply mass by abundance

$$
24 \times 79 = 1896
$$

$$
25 \times 10 = 250
$$

$$
26 \times 11 = 286
$$

Step 2: Add products

$$
1896 + 250 + 286 = 2432
$$

Step 3: Add abundances

$$
79 + 10 + 11 = 100
$$

Step 4: Divide

$$
A_r = \frac{2432}{100} = 24.32
$$

So:

$$
A_r = 24.32
$$

This type of data often comes from **mass spectrometry**, where isotopes give separate peaks.

> **Key Term: mass-to-charge ratio (m/z)**
> The value measured in mass spectrometry, equal to the mass of an ion divided by its charge.

For simple atomic spectra, the ions are often $1+$, so $m/z$ is effectively the isotopic mass. Peak height or area shows relative abundance. That lets you calculate $A_r$.

You do not need to confuse this with mass number. A peak at $m/z = 35$ for a $1+$ ion suggests an isotope of mass 35, but **relative atomic mass** is the average of all isotopes present.

Side-by-side comparison:

| Term                 | Meaning                                    | Whole number? | Applies to                    |
| -------------------- | ------------------------------------------ | ------------- | ----------------------------- |
| Mass number          | Protons + neutrons in one atom             | Yes           | One specific atom/isotope     |
| Relative atomic mass | Weighted mean mass compared with carbon-12 | Usually no    | The element as a whole sample |

This difference is one of the biggest exam traps in atomic structure.

Common student errors:

- saying $A_r$ is the mass of one atom
- forgetting to multiply by abundance before averaging
- dividing by number of isotopes instead of total abundance
- defining $A_r$ without mentioning carbon-12
- confusing relative atomic mass with mass number

How to structure an evaluative answer is less central here than in some synoptic topics, but if a longer question asks you to explain data, use this order:

1. identify what stays the same
2. identify what changes
3. link the unchanged feature to chemical behaviour
4. link the changed feature to mass or physical effect

For isotopes, that becomes:

- same proton number
- same electron arrangement
- similar chemical properties
- different neutron number
- different masses
  > **Test Tip:** If asked why isotopes have similar chemical properties, mention **electrons** or **electron configuration** explicitly. Do not say only "because they are the same element".

## Key Point Checklist

_This article has covered the following key knowledge points:_

- An atom consists of a nucleus containing protons and neutrons, with electrons outside the nucleus
- Protons have relative charge +1 and relative mass 1
- Neutrons have relative charge 0 and relative mass 1
- Electrons have relative charge -1 and relative mass 1/1836
- Atomic number is the number of protons in the nucleus
- Mass number is the total number of protons and neutrons in the nucleus
- Number of neutrons equals mass number minus atomic number
- In a neutral atom, number of electrons equals number of protons
- Ions form when electrons are gained or lost; proton number does not change
- Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons
- Isotopes have similar chemical properties because they have the same electron arrangement
- Relative atomic mass is a weighted mean based on isotopic masses and abundances and is compared with 1/12 of a carbon-12 atom

## Key Terms and Concepts

- subatomic particles
- proton
- neutron
- electron
- nucleus
- atomic number
- mass number
- ion
- isotope
- relative atomic mass
- mass-to-charge ratio (m/z)


Atomic structure - Subatomic particles