Atomic Structure: Protons, Neutrons and Electrons (KS3 Chemistry)

An atom is the smallest particle of an element that retains the chemical properties of that element. Every atom has a tiny, dense nucleus containing protons and neutrons, surrounded by electrons arranged in shells. The number of protons in the nucleus — the atomic number — defines which element the atom is.

What are protons, neutrons, and electrons?

Each type of subatomic particle has a specific charge and mass. These three properties — charge, mass, and location — are the three things you must be able to recall for KS3 and GCSE:

Particle | Location | Relative charge | Relative mass
Proton | Nucleus | +1 | 1
Neutron | Nucleus | 0 (neutral) | 1
Electron | Shells (outside nucleus) | −1 | 1/1836 (approximately 0)

The nucleus is incredibly dense: it contains almost all the mass of the atom but only about 1/10,000 of the atom's radius. If a football stadium represented a whole atom, the nucleus would be a marble on the centre spot — everything else is empty space occupied by electrons moving at high speed.

Why are atoms electrically neutral?

A neutral atom has the same number of protons (+1 each) and electrons (−1 each). The charges cancel exactly. A carbon atom has 6 protons and 6 electrons — total charge = 6(+1) + 6(−1) = 0.

If an atom gains or loses electrons, it becomes charged and is called an ion:

Losing electrons → fewer negative charges → positive ion (cation)

Gaining electrons → more negative charges → negative ion (anion)

Ion formation is covered in more depth under bonding and ionic compounds. For atomic structure at KS3, the key point is that a neutral atom always has equal numbers of protons and electrons.

What is the atomic number and the mass number?

Two numbers define an atom's identity and mass:

Atomic number (Z) = number of protons in the nucleus. This is the defining property of an element. All carbon atoms have 6 protons; all oxygen atoms have 8. You cannot change the atomic number without changing which element you have.

Mass number (A) = number of protons + number of neutrons.

From these two numbers you can work out the number of neutrons:

Number of neutrons = mass number − atomic number

Example: A carbon atom has an atomic number of 6 and a mass number of 12.

Protons = 6 (= atomic number)

Neutrons = 12 − 6 = 6

Electrons = 6 (equal to protons in a neutral atom)

How are electrons arranged in shells?

Electrons do not orbit the nucleus randomly — they occupy shells (also called energy levels) at fixed distances from the nucleus. Each shell can hold a maximum number of electrons:

Shell | Maximum electrons
1st (innermost) | 2
2nd | 8
3rd | 8 (at KS3 level)

Electrons always fill the innermost shell first before moving to the next. This order gives each element its characteristic electron configuration, which determines how it reacts chemically.

Example — sodium (atomic number 11):

1st shell: 2 electrons (full)

2nd shell: 8 electrons (full)

3rd shell: 1 electron (1 electron in the outermost shell)

Written as an electronic configuration: 2, 8, 1

The number of electrons in the outermost shell (called the valence shell) determines how reactive an element is and what type of bonds it forms.

How does electron arrangement link to the periodic table?

The periodic table is organised so that elements in the same group (column) have the same number of electrons in their outermost shell — which is why they have similar chemical properties:

Group 1 elements (e.g. lithium, sodium, potassium) all have 1 electron in their outer shell → they are all very reactive metals that form ions with a +1 charge.

Group 7 elements (the halogens: fluorine, chlorine, bromine) all have 7 electrons in their outer shell → they all need 1 more electron to fill the shell, so they form ions with a −1 charge.

Group 0 elements (noble gases: helium, neon, argon) have a full outer shell → they are very unreactive.

The period (row) an element sits in tells you how many shells its electrons occupy. Sodium is in Period 3 because it has 3 occupied shells (2, 8, 1).

What is an isotope?

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. Because they have the same atomic number (same number of protons and electrons), they are chemically identical — they react in the same way. But because their mass numbers differ, they have different masses.

Example — carbon isotopes:

Carbon-12: 6 protons, 6 neutrons (stable, makes up ~99% of natural carbon)

Carbon-13: 6 protons, 7 neutrons (stable, ~1% of natural carbon)

Carbon-14: 6 protons, 8 neutrons (radioactive, used in carbon dating)

Isotopes are written as ¹²C, ¹³C, ¹⁴C or as carbon-12, carbon-13, carbon-14.

Frequently asked questions

What is the structure of an atom in simple terms?

An atom has a very small, dense nucleus at its centre, which contains positively charged protons and neutral neutrons. Surrounding the nucleus are shells of negatively charged electrons. The number of protons defines the element, and a neutral atom always has equal numbers of protons and electrons.

How do you work out the number of neutrons in an atom?

Subtract the atomic number from the mass number: neutrons = mass number − atomic number. For example, an atom of chlorine with atomic number 17 and mass number 35 has 35 − 17 = 18 neutrons.

What is the difference between atomic number and mass number?

The atomic number is the number of protons in the nucleus and defines which element the atom is. The mass number is the total number of protons plus neutrons — it gives the approximate mass of the atom. The number of neutrons on its own is not given directly; you calculate it as mass number minus atomic number.

Why do elements in the same group have similar properties?

Elements in the same group have the same number of electrons in their outermost shell. Chemical reactions involve the outermost electrons, so elements with the same outer-shell electron count react in similar ways. For example, all Group 1 metals lose one electron to form a +1 ion, and all Group 7 halogens gain one electron to form a −1 ion.

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