Atomic Structure and Isotopes

Everything you can see, touch and breathe is built from atoms — particles so small that around a hundred million of them would sit side by side across the width of a single human hair. To understand radioactivity, which fills the second half of this topic, you first need a clear picture of what an atom is made of: a tiny central nucleus of protons and neutrons, surrounded by electrons. This lesson, part of Topic P4 (Waves and radioactivity) of OCR Gateway Combined Science A, sets out the nuclear model of the atom, the numbers that describe a nucleus, and the important idea of isotopes — atoms of the same element with different numbers of neutrons. (The story of how the atomic model developed, from Thomson to Rutherford, is covered in Topic P1; here we concentrate on the composition of the nucleus itself.)

By the end of this lesson you should be able to describe the structure of the atom and the relative charges and masses of its particles, define the atomic (proton) number and the mass (nucleon) number, use nuclide notation to work out the numbers of protons, neutrons and electrons, recall the relative sizes of an atom and its nucleus, and define an isotope.

This lesson combines AO1 understanding of atomic structure and isotopes with AO2 application when you read nuclide notation to work out the numbers of protons, neutrons and electrons in an atom or ion.

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The Structure of the Atom

An atom has two parts: a tiny central nucleus, and electrons that move around the nucleus. The nucleus contains two kinds of particle — protons and neutrons — collectively called nucleons. The electrons occupy the space around the nucleus, arranged in energy levels (also called shells).

The three sub-atomic particles have different electric charges and different masses. Because the actual values are tiny, we compare them using relative charge and relative mass, taking the proton as the reference:

Particle	Relative charge	Relative mass	Location
Proton	$+1$+1	$1$1	in the nucleus
Neutron	$0$0	$1$1	in the nucleus
Electron	$-1$−1	$\dfrac{1}{1836}$18361​ (≈ 0, negligible)	around the nucleus

Two features of this table matter enormously. First, the proton and the electron carry equal but opposite charges ($+1$+1 and $-1$−1), and a neutral atom has equal numbers of protons and electrons, so the charges cancel and the atom overall has no charge. Second, the electron's mass is only about $\dfrac{1}{1836}$18361​ of a proton's, so it is treated as negligible — almost all the mass of an atom is concentrated in its nucleus.

Exam Tip: Learn the table as a block. The two charges to remember are proton $+1$+1 and electron $-1$−1 (the neutron is neutral, $0$0); the masses are proton and neutron $= 1$=1, electron ≈ 0. Nearly all an atom's mass is in the nucleus because the electrons hardly weigh anything.

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Atomic Number and Mass Number

Two numbers describe the nucleus of any atom and are written alongside its chemical symbol:

• The atomic number (or proton number), symbol $Z$Z, is the number of protons in the nucleus. It defines which element the atom is — every carbon atom has 6 protons, every oxygen atom has 8, and so on. In a neutral atom it is also the number of electrons.
• The mass number (or nucleon number), symbol $A$A, is the total number of protons and neutrons (nucleons) in the nucleus.

The standard notation places the mass number as a superscript and the atomic number as a subscript, both to the left of the chemical symbol. For example, the most common isotope of carbon is written:

$^{12}_{\ 6}\text{C}$ 612​C

Here the top number ($12$12) is the mass number $A$A and the bottom number ($6$6) is the atomic number $Z$Z. Because the mass number counts all the nucleons, you can find the number of neutrons by subtracting:

$\text{number of neutrons} = A - Z$number of neutrons=A−Z

So carbon-12 has $12 - 6 = 6$12−6=6 neutrons. As a second example, an atom of sodium written $^{23}_{11}\text{Na}$1123​Na has $11$11 protons, $11$11 electrons (if neutral), and $23 - 11 = 12$23−11=12 neutrons.

Worked Example 1 — Reading a nuclide symbol

An atom is written $^{40}_{18}\text{Ar}$1840​Ar. State the number of protons, neutrons and electrons it contains (assume it is a neutral atom).

Step 1 — read the atomic number (bottom): $Z = 18$Z=18, so there are 18 protons.

Step 2 — a neutral atom has equal protons and electrons, so there are 18 electrons.

Step 3 — find the neutrons: $A - Z = 40 - 18 = 22$A−Z=40−18=22, so there are 22 neutrons.

Answer: $^{40}_{18}\text{Ar}$1840​Ar has 18 protons, 18 electrons and 22 neutrons.

Worked Example 2 — Another nuclide

How many protons, neutrons and electrons are in a neutral atom of $^{56}_{26}\text{Fe}$2656​Fe?

Step 1 — atomic number: $Z = 26$Z=26, so 26 protons.

Step 2 — neutral atom, so 26 electrons.

Step 3 — neutrons: $A - Z = 56 - 26 = 30$A−Z=56−26=30, so 30 neutrons.

Answer: $^{56}_{26}\text{Fe}$2656​Fe has 26 protons, 26 electrons and 30 neutrons.

Exam Tip: The bigger number is always the mass number (top); the smaller is the atomic number (bottom). To get neutrons, do top minus bottom ($A - Z$A−Z). A common misconception is to read the numbers the wrong way round — the mass number can never be smaller than the atomic number.

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The Sizes of the Atom and the Nucleus

Although almost all the mass of an atom sits in the nucleus, the nucleus is astonishingly small compared with the atom as a whole. The order of magnitude of each is worth memorising:

• the radius of a typical atom is about $1 \times 10^{-10}\text{ m}$1×10−10 m (a tenth of a nanometre);
• the radius of a nucleus is about $1 \times 10^{-14}\text{ m}$1×10−14 m — roughly ten thousand times smaller than the atom.

If the nucleus were the size of a pea sitting on the centre spot of a football stadium, the edge of the atom (where the electrons are) would be roughly at the back of the stands. This means an atom is mostly empty space: a minute, dense nucleus, a few tiny electrons, and a vast emptiness between them.

Exam Tip: Remember the two orders of magnitude: atom ≈ $1 \times 10^{-10}\text{ m}$1×10−10 m, nucleus ≈ $1 \times 10^{-14}\text{ m}$1×10−14 m. The nucleus is about $10\,000$10000 times smaller than the atom, so the atom is mostly empty space — a phrase examiners reward.

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What Is an Isotope?

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. Because they have the same number of protons (the same atomic number $Z$Z), they are the same element with the same chemical behaviour; but because they have different numbers of neutrons, they have different mass numbers $A$A.

A classic example is carbon, which has three isotopes worth knowing, all with 6 protons: