Fundamental particles and how the model of the atom developed
Everything in chemistry, from why sodium reacts with water to why a mass spectrometer can tell two isotopes apart, comes back to what an atom is made of and where its electrons are.
This first section sets up that picture.
Atomic structure is examined on Paper 1 and Paper 3, never on Paper 2; the amount-of-substance sections later in this unit are examined on all three papers.
What problem the atomic model solves
Chemists need a picture of the atom that explains what they measure: that elements combine in fixed ratios, that some particles carry charge, that an atom has almost all of its mass in a tiny space, and that each element has its own pattern of reactions.
No single experiment gave that picture.
You need to appreciate that knowledge and understanding of atomic structure has evolved over time: each model was kept while it explained the evidence, and replaced or refined when new evidence did not fit.
How the model changed, and why
Model (scientist, date)
What the atom was thought to be
The evidence that led to it, or that it could not explain
Solid sphere (John Dalton, early 1800s)
Tiny indivisible spheres; all atoms of one element identical
Explained fixed combining ratios, but had no charged particles inside
"Plum pudding" (J. J. Thomson, 1904)
A sphere of positive charge with negative electrons spread through it
Thomson had discovered the electron in 1897, so atoms could not be indivisible
Nuclear model (Ernest Rutherford, 1911)
A tiny, dense, positive nucleus holding almost all the mass, with electrons around it and mostly empty space
Geiger and Marsden's gold foil experiment (1909): most alpha particles passed straight through, a few were deflected through large angles
Electron shells (Niels Bohr, 1913)
Electrons in fixed energy levels (shells) around the nucleus
Explained why atoms emit light only at particular energies
Neutron added (James Chadwick, 1932)
The nucleus holds protons and neutrons
Explained why atoms are heavier than their protons alone
Orbitals (Erwin Schrödinger and others, from 1926)
Electrons in orbitals, grouped into sub-shells within each shell
Explains ionisation energy patterns and electron configurations (later in this unit)
The set-up of the gold foil experiment, and the paths the alpha particles took, are shown below.
The gold foil result is the turning point, so understand it at particle level.
Alpha particles are positive.
If the positive charge of an atom were spread thinly through it, as in the plum pudding model, every alpha particle would pass through with only a tiny deflection.
Instead a few bounced back.
That can only happen if the positive charge and almost all the mass are packed into a very small region: the nucleus.
Most alpha particles went straight through because most of the atom is empty space.
The nucleus is roughly 100 000 times smaller across than the atom.
The model you use in this course is the modern one: a nucleus of protons and neutrons, surrounded by electrons arranged in shells and sub-shells (orbitals).
Older models are not "wrong" in a useless way.
Bohr's shells still explain the big jumps in successive ionisation energies, and you will use them later in this unit.
They are simply less complete.
The three fundamental particles
You need to know protons, neutrons and electrons: relative charge and relative mass.
Particle
Relative charge
Relative mass
Where it is
proton
+1
1
nucleus
neutron
0
1
nucleus
electron
−1
1/1836 (about 0.0005)
around the nucleus, in shells
Relative values compare the particles with each other, so they have no units.
A proton and a neutron have almost the same mass; an electron is about 1836 times lighter.
Some textbooks round the electron's relative mass to 1/1840; either is accepted when you write a fraction close to 1/2000 or the decimal 0.0005.
And you need to know that an atom consists of a nucleus containing protons and neutrons surrounded by electrons.
Because the electrons are so light, almost all of an atom's mass is in the nucleus.
Because an atom is neutral, the number of electrons equals the number of protons: the +1 and −1 charges cancel exactly.
Why the arrangement matters for chemistry
The nucleus decides the identity of the element (its number of protons) and almost all of its mass.
The electrons decide its chemistry.
When atoms react, they lose, gain or share outer electrons; the nucleus is untouched.
This is why the next sections spend so long on where electrons are: shells and sub-shells explain ionisation energies, the shape of the Periodic Table and the formulae of ions.
It is also why mass spectrometry, which measures the mass of the nucleus plus its electrons, can tell isotopes apart, while ordinary reactions cannot.
- Nucleus: the very small, dense, positively charged centre of an atom, containing protons and neutrons, which holds almost all of the atom's mass.
- Proton: a particle in the nucleus with relative charge +1 and relative mass 1.
- Neutron: a particle in the nucleus with relative charge 0 and relative mass 1.
- Electron: a particle around the nucleus with relative charge −1 and relative mass 1/1836.
- Each mark is a marking point (M1, M2, M3 ...), and you earn each one separately.
- A correct final answer on its own gets full marks, unless the question says "Show your working" or asks you to justify it. With a wrong answer, you still get a process mark for each correct step.
- Errors are carried forward (ECF): a later step that uses your wrong earlier value correctly still scores, unless the slip makes the rest of the problem trivial.
- Precision (decimal places, significant figures, an integer) costs you a mark only when the question asks for it.
- "Right + wrong = wrong": a wrong or contradictory extra answer cancels a right one, so commit to one answer.
- In a 6-mark levels item, the stages you cover put you in one of three levels, and how clearly you link them decides the higher or lower mark inside the level.
In the gold foil experiment, a beam of alpha particles was fired at a very thin sheet of gold.
Most passed straight through; a very small number were deflected through large angles.
Explain how these results showed that the plum pudding model was wrong. [3]
1. M1: most alpha particles passed straight through, so most of the atom is empty space.
2. M2: a few were deflected through large angles (some back towards the source), so they met a small region of concentrated positive charge that repelled them.
3. M3: this region also holds most of the mass; in the plum pudding model the positive charge and mass are spread out, so no alpha particle would be deflected so far.
Answer: the atom has a tiny, dense, positive nucleus surrounded by mostly empty space, which the plum pudding model does not have.
Each mark links one observation to one conclusion: "most passed through" gives "empty space", "a few deflected" gives "small, positive, dense nucleus".
A student lists three features of atoms: (i) electrons occupy orbitals grouped into sub-shells; (ii) atoms of one element are identical, indivisible spheres; (iii) electrons are embedded in a sphere of positive charge.
Identify which feature belongs to the modern model of the atom, and state the model each of the other two features belongs to. [2]
1. M1: feature (i) belongs to the modern model.
2. M2: (ii) is Dalton's solid-sphere model; (iii) is Thomson's plum pudding model.
Answer: (i) modern; (ii) Dalton; (iii) Thomson.
Recent papers have asked about historical models of the atom in this way.
Name the model each feature belongs to: an older model's name written on its own, with no link to the feature, does not answer the question.
A ¹²C atom contains 6 protons, 6 neutrons and 6 electrons.
Use the relative masses of the particles to calculate the percentage of the mass of the atom that is due to its electrons. [2]
1. M1: mass of electrons = 6 × (1/1836) = 0.003268; total mass = 12 + 0.003268 = 12.003268
2. M2: percentage = (0.003268 / 12.003268) × 100 = 0.0272%
Answer: 0.0272% (0.027% to 2 s.f.).
Over 99.97% of the mass is in the nucleus.
This is why the mass number, which counts only protons and neutrons, is such a good measure of an atom's mass.
Where this leads
The particle numbers you meet here are the basis of the next section (mass number, atomic number and isotopes).
The idea that electrons sit in energy levels around a positive nucleus explains ionisation energies, and the idea that the nucleus holds the mass is the basis of mass spectrometry.
Electron arrangement also explains bonding in the next unit.
When a question asks about the development of the atomic model, link each observation to the feature of the model it supports, and say which model you mean.
For the nuclear model, "most alpha particles passed through" goes with "mostly empty space", and "a few were deflected through large angles" goes with "a small, dense, positive nucleus".
- Calling the electron's mass zero. It is tiny (1/1836), not zero. A relative mass of "0" for the electron loses the mark. It is tempting because the electron's mass is often ignored in sums, but it is still a real mass.
- Putting electrons in the nucleus. The nucleus holds protons and neutrons only. Electrons are around it, in shells. Students sometimes mix this up because the nucleus "contains the charge"; it contains only the positive charge.
- Saying the plum pudding model had a nucleus. It did not: its positive charge was spread through the whole atom. The nucleus was Rutherford's idea, from the gold foil results.
- Giving charges in coulombs. Relative charges are +1, 0 and −1, with no units. The real charge of a proton, 1.60 × 10⁻¹⁹ C, is not what "relative charge" asks for.
- Treating an older model as useless. Bohr's shells still explain large jumps in ionisation energies. Say what a model explains and what it cannot.
Check your understanding
1. Which feature of the atom was first proposed in the nuclear model? [1]
- A atoms contain negatively charged electrons
- B electrons occupy fixed energy levels
- C most of the mass of an atom is in a very small positive centre
- D electrons occupy orbitals within sub-shells
2. State the relative charge and the relative mass of a neutron. [2]
3. Complete the table. [3]
Particle
Relative charge
Relative mass
Position in the atom
proton
electron
4. This question is about the gold foil experiment.
4.1 State the observation that showed most of an atom is empty space. [1]
4.2 Explain why the plum pudding model cannot account for the alpha particles that were deflected through large angles. [2]
5. An atom of ²³Na contains 11 protons, 12 neutrons and 11 electrons. Calculate the percentage of the mass of the atom that is due to its electrons. [2]
6. The neutron was discovered about 15 years after the proton was identified. Suggest why the neutron was harder to detect. [2]
7. A student says: "Bohr's model of the atom is wrong, so it is no longer useful in chemistry." Explain why the student's statement is not correct. [2]
8. State why an atom has no overall charge. [1]
9. Tick (✓) one box to show the particle with the smallest mass. [1]
- proton
- neutron
- electron
10. Give one feature of the modern model of the atom that is not part of Bohr's model. [1]
Answers
1. C. Rutherford's nuclear model put almost all the mass and positive charge into a tiny nucleus. A (electrons) is Thomson's discovery, already in the plum pudding model. B (fixed energy levels) is Bohr's later refinement. D (orbitals in sub-shells) is the modern model. [1]
2. M1: relative charge 0. M2: relative mass 1. [2]
3. Proton: +1, 1, nucleus. Electron: −1, 1/1836 (allow 1/1840 or 0.0005), around the nucleus (in shells or orbitals). [3: any five correct cells 2, all six 3; do not allow 0 for the electron's mass]
4. 4.1 Most alpha particles passed straight through the gold foil (without being deflected). [1]
4.2 M1: in the plum pudding model the positive charge (and the mass) is spread throughout the atom, so it is too thinly spread to repel an alpha particle strongly.
M2: a large deflection needs a small region of concentrated positive charge and mass (a nucleus), which the model does not have. [2]
5. M1: mass of electrons = 11 × (1/1836) = 0.005991; total = 23 + 0.005991 = 23.005991. M2: (0.005991 / 23.005991) × 100 = 0.026%. [2]
6. M1: the neutron has no charge. M2: so it is not deflected by electric or magnetic fields (and does not ionise matter directly), which is how charged particles were detected. [2]
7. M1: Bohr's model still explains some evidence well, for example the large jumps in successive ionisation energies (electrons in shells, or energy levels). M2: it is replaced only where it fails, for example it does not include sub-shells or orbitals, which the modern model needs. [2]
8. The number of electrons equals the number of protons, so the negative and positive charges cancel. [1]
9. Electron. [1]
10. Electrons occupy orbitals (or sub-shells such as s, p and d) within the shells; or the nucleus contains neutrons as well as protons. [1; Bohr's model had shells but not sub-shells]