The Atom and the Nuclear Model

Everything you can see, touch and breathe is built from atoms — particles so small that around a hundred million of them would fit across the width of a single human hair. Yet the picture scientists hold of the atom today, with a tiny dense nucleus surrounded by orbiting electrons, was not obvious and did not arrive all at once. It was pieced together over more than a century, as one experiment after another overturned the model that came before it. This lesson, which opens Topic P6 (Radioactivity) of OCR Gateway Science A, sets out the structure of the atom — protons, neutrons and electrons, and the numbers that describe a nucleus — and then traces the history of the atomic model from Dalton's solid spheres to the modern nuclear atom, paying special attention to the alpha-scattering experiment that revealed the nucleus.

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, recall the relative sizes of an atom and its nucleus, and describe how the atomic model developed — especially what the alpha-scattering experiment showed.

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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 — 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.

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 slip 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 $10^{-10}\text{ m}$10−10 m (a tenth of a nanometre);
• the radius of a nucleus is about $10^{-14}\text{ m}$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. This single fact — a tiny dense centre in a mostly empty atom — is exactly what the alpha-scattering experiment revealed, as we shall see.

Exam Tip: Remember the two orders of magnitude: atom ≈ $10^{-10}\text{ m}$10−10 m, nucleus ≈ $10^{-14}\text{ m}$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 when you explain alpha scattering.

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How the Atomic Model Developed

Our modern picture of the atom is the result of a long chain of discoveries, each building on or correcting the last. You should be able to put these models in order and say what new evidence prompted each change.

graph LR
  A[Dalton<br/>solid spheres] --> B[Thomson<br/>plum pudding]
  B --> C[Rutherford<br/>nuclear atom]
  C --> D[Bohr<br/>energy levels]
  D --> E[Chadwick<br/>neutron]

Dalton (early 1800s) — solid spheres. John Dalton proposed that each element is made of tiny, identical, indivisible solid spheres called atoms, which cannot be broken down or changed. This was a huge step, but it pictured the atom as a featureless ball with no internal parts.

Thomson (1897) — the plum-pudding model. J. J. Thomson discovered the electron — a tiny, negatively charged particle far smaller than an atom — by experimenting with cathode rays. Since atoms are neutral, he reasoned the negative electrons must be balanced by positive charge, and proposed the plum-pudding model: a ball of positive charge (the "pudding") with negative electrons dotted through it (the "plums"). The atom was no longer indivisible; it had parts.

Rutherford (1909–1911) — the nuclear atom. Ernest Rutherford, with Hans Geiger and Ernest Marsden, carried out the alpha-scattering experiment (described in detail below). Its surprising results could not be explained by the plum-pudding model and led Rutherford to propose the nuclear model: the atom has a tiny, dense, positively charged nucleus at its centre, with electrons around the outside and mostly empty space in between.

Bohr (1913) — electrons in energy levels. Niels Bohr refined the nuclear model. A problem with Rutherford's atom was that orbiting electrons should spiral into the nucleus, making the atom unstable. Bohr proposed that electrons orbit the nucleus only at fixed distances, in definite energy levels (shells), and do not spiral in. This explained why atoms are stable and matched evidence from the light atoms give out.

Chadwick (1932) — the neutron. For some years the nucleus was thought to contain only protons, but its mass was too large for that. James Chadwick discovered the neutron — an uncharged particle of about the same mass as a proton — which completed the modern picture: a nucleus of protons and neutrons, surrounded by electrons in energy levels.

Scientist	Model / discovery	Key idea added
Dalton	Solid spheres	Atoms are tiny, indivisible spheres
Thomson	Plum pudding	Atoms contain electrons; positive "dough" with negative "plums"
Rutherford	Nuclear atom	A tiny dense positive nucleus; mostly empty space
Bohr	Energy levels	Electrons orbit in fixed energy levels (shells)
Chadwick	The neutron	The nucleus contains neutrons as well as protons

Exam Tip: A favourite exam task is to put these models in order and say why each changed. The trigger for the big jump from Thomson to Rutherford was the alpha-scattering experiment; Bohr's change explained why atoms are stable; Chadwick's added the neutron to account for the nuclear mass.

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The Alpha-Scattering (Geiger–Marsden) Experiment

The single most important experiment in this story is alpha scattering, carried out under Rutherford's direction. In it, a beam of alpha particles (small, positively charged particles emitted by a radioactive source) was fired at an extremely thin sheet of gold foil. A detecting screen around the foil recorded where the alpha particles went after passing through (or bouncing off) it.

The three observations — and what each one told Rutherford — are the heart of the experiment:

Observation	What was seen	What it showed
1	Most alpha particles passed straight through the foil, undeflected	The atom is mostly empty space
2	Some alpha particles were deflected through small angles	The centre of the atom is positively charged (it repelled the positive alphas)
3	A very few alpha particles bounced straight back (deflected through more than 90°)	The positive charge and mass are concentrated in a tiny, dense nucleus

The plum-pudding model predicted that the alpha particles should all pass straight through with at most very slight deflections, because the positive charge was spread thinly throughout the atom. The fact that a few alphas bounced almost straight back was, in Rutherford's words, almost as incredible as firing a shell at tissue paper and having it come back at you. The only way to explain it was that almost all the atom's mass and all its positive charge are packed into a minute central nucleus, with the rest of the atom being empty space through which most alphas sail unhindered. This is exactly the nuclear model.

Worked Example — Interpreting an alpha-scattering result

In the alpha-scattering experiment, a very small number of alpha particles bounced almost straight back off the gold foil. What does this observation tell us about the atom, and why?

Step 1 — recall the property of an alpha particle: it is positively charged and has appreciable mass.

Step 2 — explain the bounce-back: to repel a fast, positive alpha particle straight back, it must meet something positively charged (like charges repel) that is also dense and massive enough not to be pushed aside.

Step 3 — locate that charge and mass: because only a very few alphas bounce back, this region must be extremely small — a tiny nucleus holding the positive charge and most of the mass.

Answer: the bounce-back shows the atom has a tiny, dense, positively charged nucleus, because only a concentrated positive charge could repel a positive alpha particle straight back, and its rarity shows the nucleus is very small.

Exam Tip: Link each observation to its conclusion: straight through → mostly empty space; deflected → positive nucleus; bounced back → tiny, dense, massive nucleus. Quoting "like charges repel" explains why the positive nucleus deflects the positive alpha particles.

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Common Misconceptions

Misconception	The correct idea
"The electron and proton have very different charge sizes"	Their charges are equal in size but opposite: proton $+1$+1, electron $-1$−1
"Electrons make up a lot of an atom's mass"	The electron's mass is about $\dfrac{1}{1836}$18361​ of a proton's — negligible; the mass is in the nucleus
"The mass number is the number of protons"	The mass number is protons plus neutrons; the atomic number is the protons
"The nucleus fills most of the atom"	The nucleus is about $10\,000$10000 times smaller than the atom; the atom is mostly empty space
"Rutherford discovered the electron"	Thomson discovered the electron; Rutherford discovered the nucleus
"Alpha scattering proved the plum-pudding model"	It disproved plum pudding and led to the nuclear model

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Model Answer — What Alpha Scattering Revealed

Question (6 marks): Describe the alpha-scattering experiment and explain how its results led scientists to replace the plum-pudding model with the nuclear model of the atom.

Mid-band response: "They fired alpha particles at gold foil. Most went straight through, but some bounced back. This showed the atom has a nucleus in the middle, so the plum-pudding model was wrong."

Examiner-style commentary: The basic method and the headline conclusion are present, but the reasoning is thin. To climb a band, give all three observations, say what each one shows, and explain why the bounce-back means a tiny, dense, positive nucleus.

Stronger response: "Alpha particles were fired at a thin sheet of gold foil and a screen detected where they went. Most passed straight through, showing the atom is mostly empty space. Some were deflected, showing there is a concentration of positive charge. A few bounced straight back, showing this positive charge is in a tiny, dense nucleus. The plum-pudding model could not explain the particles bouncing back, so it was replaced by the nuclear model."

Examiner-style commentary: A clear, correct answer giving all three observations and their conclusions. To reach the top band, state that like charges repel (so a positive nucleus repels the positive alphas) and contrast with what the plum-pudding model predicted.

Top-band response: "In the experiment, a beam of alpha particles (small, fast, positively charged particles) was fired at a very thin gold foil, and a detector recorded the directions in which they emerged. There were three observations. (1) Most alpha particles passed straight through undeflected — showing the atom is mostly empty space. (2) Some were deflected through small angles — showing the atom contains a concentration of positive charge, which repels the positive alpha particles because like charges repel. (3) A very few bounced almost straight back — showing that the positive charge and almost all the mass are packed into a tiny, dense nucleus, since only a small, massive, highly charged region could turn a fast alpha particle right around. The plum-pudding model predicted the positive charge was spread out, so all the alphas should pass through with only slight deflection; it could not explain the back-scattering. Rutherford therefore replaced it with the nuclear model — a minute, dense, positive nucleus surrounded by mostly empty space containing the electrons."

Examiner-style commentary: Full marks. It describes the method, gives all three observations with the correct conclusion for each, uses "like charges repel", contrasts the plum-pudding prediction with the result, and names the nuclear model — a complete account of how the evidence forced the change.

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Exam Tips

• An atom is a tiny dense nucleus (protons + neutrons) surrounded by electrons in energy levels; nearly all its mass is in the nucleus.
• Relative charges: proton $+1$+1, neutron $0$0, electron $-1$−1; relative masses: proton and neutron $1$1, electron ≈ 0 ($\tfrac{1}{1836}$18361​).
• Atomic number $Z$Z = number of protons; mass number $A$A = protons + neutrons; neutrons $= A - Z$=A−Z.
• Sizes: atom ≈ $10^{-10}\text{ m}$10−10 m, nucleus ≈ $10^{-14}\text{ m}$10−14 m, so the atom is mostly empty space.
• Know the model order Dalton → Thomson → Rutherford → Bohr → Chadwick and the three alpha-scattering observations.

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Summary

• An atom is made of a central nucleus (protons and neutrons) surrounded by electrons in energy levels.
• Proton: charge $+1$+1, mass $1$1. Neutron: charge $0$0, mass $1$1. Electron: charge $-1$−1, mass ≈ $\tfrac{1}{1836}$18361​ (negligible).
• The atomic (proton) number $Z$Z counts the protons and defines the element; the mass (nucleon) number $A$A counts protons + neutrons; neutrons $= A - Z$=A−Z.
• An atom has a radius of about $10^{-10}\text{ m}$10−10 m and a nucleus about $10^{-14}\text{ m}$10−14 m, so it is mostly empty space.
• The model developed Dalton (spheres) → Thomson (plum pudding) → Rutherford (nuclear, from alpha scattering) → Bohr (energy levels) → Chadwick (neutron).
• Alpha scattering: most pass through (empty space); some deflect (positive nucleus); a few bounce back (tiny, dense, massive nucleus).

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This content is aligned with OCR Gateway Science A GCSE Physics (J249), Topic P6 Radioactivity (atomic structure; relative charge and mass of sub-atomic particles; atomic and mass number; the history of the atomic model; the alpha-scattering experiment). Refer to the official OCR specification document for the exact wording.