# Protons, Neutrons and Electrons

> Charges and masses of the three particles, atomic and mass numbers, isotopes and ions.

From HSC Chemistry on WhizzNotes, free study notes for NSW students: https://whizznotes.com/chemistry/subatomic-particles/. Updated 3 November 2026.
Year 11 and Year 12 boxes go deeper than the Year 10 text around them.

### Story time: The atom's ID card
At school, every student has an ID card with a number on it. No two students share a number.

Atoms have an ID number too. It is the number of **protons** in the nucleus. Every carbon atom in the universe has exactly 6 protons. Give it a seventh proton and it is no longer carbon. It has become nitrogen.

Neutrons and electrons can change a little without changing the element. Protons cannot. Once you can count the three particles, you can read almost everything about an atom from a single box on the periodic table.

**In short:** Atoms are made of protons (charge +1, in the nucleus), neutrons (no charge, in the nucleus) and electrons (charge −1, in shells). The atomic number is the number of protons, the mass number is protons plus neutrons, and isotopes are atoms of one element with different numbers of neutrons.

## The three subatomic particles

| Particle | Where | Relative charge | Relative mass |
|---|---|---|---|
| **Proton** ($p^+$) | Nucleus | $+1$ | 1 |
| **Neutron** ($n^0$) | Nucleus | $0$ | 1 |
| **Electron** ($e^-$) | Shells around the nucleus | $-1$ | $\frac{1}{1836}$ (almost 0) |

Protons and neutrons are together called **nucleons**.

Real atoms are far too light to weigh in grams easily. A carbon-12 atom has a mass of about $1.99 \times 10^{-23}$ g. So chemists use the **atomic mass unit** (amu, also written u):

$$1 \text{ amu} = \tfrac{1}{12} \text{ of the mass of a carbon-12 atom} \approx 1.66 \times 10^{-24} \text{ g}$$

A proton and a neutron each have a mass of about 1 amu. That is where the "relative mass 1" in the table comes from.

## Atomic number and mass number

- **Atomic number** ($Z$) = number of **protons**. It decides which element it is.
- **Mass number** ($A$) = number of **protons + neutrons**.

So:

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

In a **neutral atom**, the positive and negative charges balance, so

$$\text{electrons} = \text{protons} = Z$$

## Nuclear symbol notation

We write an atom like this, with the mass number on top and the atomic number underneath:

$$^{A}_{Z}\mathrm{X} \qquad \text{for example} \qquad ^{23}_{11}\mathrm{Na}$$

### Worked example: Counting particles in sodium
For $^{23}_{11}\mathrm{Na}$:
- Protons = $Z$ = 11
- Electrons = 11 (neutral atom)
- Neutrons = $A - Z = 23 - 11 = 12$

### Exam tip: The bigger number is the mass number
On many periodic tables the atomic number is a whole number and the "atomic mass" is a decimal like 22.99. The **atomic number is always the smaller** of the two. Round the atomic mass to get the most common mass number.

## Isotopes

**Isotopes** are atoms of the **same element** (same number of protons) with **different numbers of neutrons**. So they have different mass numbers.

| Isotope | Protons | Neutrons | Electrons |
|---|---|---|---|
| Carbon-12, $^{12}_{6}\mathrm{C}$ | 6 | 6 | 6 |
| Carbon-13, $^{13}_{6}\mathrm{C}$ | 6 | 7 | 6 |
| Carbon-14, $^{14}_{6}\mathrm{C}$ | 6 | 8 | 6 |

Isotopes of an element have the **same chemical properties**, because chemistry depends on electrons, and they have the same electrons. They have slightly different **physical** properties, such as mass and density. Some, like carbon-14, are **radioactive**. Scientists use carbon-14 to find the age of ancient bones and wood.

## Ions

An **ion** is an atom (or group of atoms) that has **gained or lost electrons**, so it has a charge. The protons do not change.

- **Lose** electrons: more protons than electrons, so the ion is **positive**. A positive ion is a **cation**. Metals form cations.
- **Gain** electrons: more electrons than protons, so the ion is **negative**. A negative ion is an **anion**. Non-metals form anions.

$$\text{charge} = \text{protons} - \text{electrons}$$

### Worked example: Ions
A sodium atom has 11 protons and 11 electrons. It loses 1 electron to form $\mathrm{Na^+}$: 11 protons, 10 electrons, charge $+1$.

An oxygen atom has 8 protons and 8 electrons. It gains 2 electrons to form $\mathrm{O^{2-}}$: 8 protons, 10 electrons, charge $-2$.

### Watch out: Ions change electrons, not protons
A positive ion is **not** formed by gaining protons. Protons are locked in the nucleus. Only electrons move during chemical reactions.

### Key idea: "Cat-ions are paws-itive"
Remember: a **cat**ion has **paws**, so it is **pos**itive. An anion is negative.

**Mind map**

```text
Subatomic particles
  Proton
    Charge +1, mass 1
    In nucleus
    Number = atomic number
  Neutron
    Charge 0, mass 1
    Neutrons = A − Z
  Electron
    Charge −1, mass almost 0
    Equal to protons if neutral
  Isotopes
    Same protons
    Different neutrons
  Ions
    Lose e⁻: positive cation
    Gain e⁻: negative anion
```

### Year 11: Relative atomic mass
The atomic mass on the periodic table is a **weighted average** of the masses of all the natural isotopes. That is why it is rarely a whole number.

Chlorine is 75% chlorine-35 and 25% chlorine-37.

$$A_r = \frac{(75 \times 35) + (25 \times 37)}{100} = \frac{2625 + 925}{100} = 35.5$$

The average is closer to 35 because more chlorine atoms are chlorine-35.

### Year 12: The mass spectrometer
A **mass spectrometer** measures the masses and amounts of isotopes. Atoms are turned into positive ions, sped up by an electric field, then bent by a magnetic field. Lighter ions bend more. A detector counts how many ions arrive at each mass to charge ratio ($m/z$).

The result is a **mass spectrum**: a graph of relative abundance against $m/z$. Each peak is an isotope, and the peak heights give the percentages used to calculate $A_r$. Chemists also use mass spectra to identify unknown molecules from their fragments.

## Quick summary

- Proton: $+1$, mass 1, in the nucleus. Neutron: $0$, mass 1, in the nucleus. Electron: $-1$, mass almost 0, in shells.
- Atomic number $Z$ = protons. Mass number $A$ = protons + neutrons.
- Neutrons = $A - Z$. In a neutral atom, electrons = protons.
- **Isotopes**: same protons, different neutrons. Same chemistry.
- **Ions**: atoms that lost electrons (positive cations) or gained electrons (negative anions).

## Quiz

Try these questions from the chapter quiz.

1. The atomic number of an element tells you the number of:
   - A. Protons
   - B. Neutrons
   - C. Protons plus neutrons
   - D. Shells
   - Answer: A. Protons. The atomic number is the proton number. It identifies the element.

2. Which particle has a relative charge of −1?
   - A. Proton
   - B. Neutron
   - C. Electron
   - D. Nucleus
   - Answer: C. Electron. Electrons are negative.

3. An atom has a mass number of 27 and an atomic number of 13. How many neutrons does it have?
   - A. 13
   - B. 14
   - C. 27
   - D. 40
   - Answer: B. 14. Neutrons = A − Z = 27 − 13 = 14.

The whole chapter quiz is on the page, for signed-in students: https://whizznotes.com/chemistry/subatomic-particles/#quiz