What you'll be able to do by the end
- Describe the structure of an atom and the properties of its particles
- Use atomic number and mass number to work out particle numbers
- Explain what isotopes are and calculate relative atomic mass
- Describe how the model of the atom changed over time
- Write electronic structures for the first 20 elements
- Explain how the periodic table is organised
- Describe the properties of Group 1, Group 7 and Group 0 elements
- Explain ionic, covalent and metallic bonding
- Link structure to properties for each bonding type
- Explain the properties of giant covalent structures and allotropes of carbon
Rates, energy, organic chemistry, electrolysis — none of it makes sense without atomic structure and bonding. A student who is secure here finds the rest of the course manageable; one who isn't will struggle for two years.
Part 1 — Atomic structure
The three subatomic particles
| Particle | Relative charge | Relative mass | Where |
|---|---|---|---|
| Proton | +1 | 1 | Nucleus |
| Neutron | 0 | 1 | Nucleus |
| Electron | −1 | Very small (1/1836) | Shells around the nucleus |
Atoms have no overall charge because the number of protons equals the number of electrons.
The size of an atom
- Atomic radius: about 0.1 nanometres (1 × 10⁻¹⁰ m)
- Nuclear radius: about 1 × 10⁻¹⁴ m — around 1/10,000 of the atom
Almost all of an atom is empty space, and almost all of its mass is in the nucleus.
If an atom were the size of a football stadium, the nucleus would be a pea on the centre spot.
Atomic number and mass number
| Term | What it tells you |
|---|---|
| Atomic number | Number of protons (and therefore electrons in a neutral atom) |
| Mass number | Number of protons + neutrons |
Neutrons = mass number − atomic number
Change the number of protons and you have a different element. Nothing else does that.
Change the neutrons → same element, different isotope. Change the electrons → same element, but now an ion.
Isotopes
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
| Isotope | Protons | Neutrons | Electrons |
|---|---|---|---|
| Chlorine-35 | 17 | 18 | 17 |
| Chlorine-37 | 17 | 20 | 17 |
Isotopes have identical chemical properties because chemical behaviour depends on electrons, and they have the same number.
Calculating relative atomic mass
RAM = (mass₁ × %₁ + mass₂ × %₂) ÷ 100
(35 × 75) + (37 × 25) = 2625 + 925 = 3550
3550 ÷ 100 = 35.5
That's why the periodic table shows chlorine as 35.5 — it's a weighted average, not the mass of any single atom.
Part 2 — How the model of the atom developed
This is a history of science topic, and it tests whether students understand that models change when evidence changes.
| Date | Model | Evidence and change |
|---|---|---|
| Pre-1800s | Atoms as tiny indivisible spheres | Dalton's atomic theory |
| 1897 | Plum pudding — a ball of positive charge with electrons embedded | J.J. Thomson discovered the electron, so atoms must contain smaller particles |
| 1909 | Nuclear model | Rutherford's alpha scattering experiment |
| 1913 | Bohr model — electrons in fixed shells | Calculations agreed with experimental observations |
| Later | Protons named; neutrons discovered | Chadwick, 1932 |
The alpha scattering experiment
What they did: fired positively charged alpha particles at a very thin sheet of gold foil.
What they expected: on the plum pudding model, all the particles should pass straight through with minor deflection.
What they observed:
| Observation | Conclusion |
|---|---|
| Most passed straight through | The atom is mostly empty space |
| Some were deflected | The centre has a positive charge |
| A very few bounced straight back | The mass is concentrated in a tiny nucleus |
It's the clearest example in the whole specification of evidence overturning a model.
Learn the three observations and the three conclusions as pairs. Questions almost always ask you to link them.
Part 3 — Electronic structure
Electrons occupy shells around the nucleus, filling from the inside out.
| Shell | Maximum electrons |
|---|---|
| 1st | 2 |
| 2nd | 8 |
| 3rd | 8 |
Writing electronic structures
| Element | Atomic number | Electronic structure |
|---|---|---|
| Hydrogen | 1 | 1 |
| Helium | 2 | 2 |
| Carbon | 6 | 2,4 |
| Oxygen | 8 | 2,6 |
| Sodium | 11 | 2,8,1 |
| Chlorine | 17 | 2,8,7 |
| Calcium | 20 | 2,8,8,2 |
What the structure tells you
The number of electrons in the outer shell = the group number.
The number of shells = the period number. Sodium has 3 shells → Period 3.
This is why the periodic table is arranged as it is. It isn't an arbitrary chart — it's a map of electronic structure.
Part 4 — The periodic table
How it's organised
- Elements arranged in order of increasing atomic number
- Groups are the vertical columns — elements with the same number of outer electrons, so similar chemical properties
- Periods are the horizontal rows — elements with the same number of shells
Mendeleev
Mendeleev arranged the elements by atomic weight but left gaps where he thought elements were missing, and swapped some pairs so that elements with similar properties lined up.
He then predicted the properties of undiscovered elements — and when they were found, they matched.
Why the table works better now: we order by atomic number, not atomic weight. This resolves the pairs Mendeleev had to swap, and it was only possible once protons were discovered.
Metals and non-metals
| Metals | Non-metals | |
|---|---|---|
| Position | Left and centre | Right |
| Electrons | Lose electrons to form positive ions | Gain electrons to form negative ions |
| Conductivity | Good conductor | Poor conductor (except graphite) |
| Appearance | Shiny, malleable | Dull, brittle when solid |
| Melting point | Generally high | Generally low |
Part 5 — The groups
Group 1 — the alkali metals
Lithium, sodium, potassium, rubidium, caesium, francium. One electron in the outer shell, which is lost easily to form a 1+ ion.
Properties:
- Soft, can be cut with a knife
- Low density — lithium, sodium and potassium float on water
- React vigorously with water, producing hydrogen and a hydroxide
Reactivity increases down the group.
Why: going down the group, atoms have more shells, so the outer electron is further from the nucleus and shielded by inner shells. It is therefore more easily lost, so the element is more reactive.
| Metal | Reaction with water |
|---|---|
| Lithium | Fizzes steadily |
| Sodium | Melts into a ball, moves rapidly, fizzes |
| Potassium | Ignites with a lilac flame |
Group 7 — the halogens
Fluorine, chlorine, bromine, iodine, astatine. Seven electrons in the outer shell, so they gain one to form a 1− ion. They exist as diatomic molecules — Cl₂, Br₂, I₂.
| Halogen | State at room temperature | Colour |
|---|---|---|
| Fluorine | Gas | Pale yellow |
| Chlorine | Gas | Green |
| Bromine | Liquid | Orange-brown |
| Iodine | Solid | Grey (purple vapour) |
Melting and boiling points increase down the group — larger molecules have stronger intermolecular forces.
Reactivity decreases down the group.
Why: going down, atoms have more shells, so the outer shell is further from the nucleus and shielded. It is therefore harder to attract an extra electron, so the element is less reactive.
Group 1: reactivity increases down the group — losing an electron gets easier.
Group 7: reactivity decreases down the group — gaining an electron gets harder.
Same reason, opposite outcome. Distance and shielding. Once you see that, you don't have to memorise two facts.
Displacement reactions
A more reactive halogen displaces a less reactive one from its compound.
Group 0 — the noble gases
Helium, neon, argon, krypton, xenon, radon. Full outer shells — helium has 2, the rest have 8. Therefore:
- Unreactive — they have no need to gain, lose or share electrons
- Monatomic — they exist as single atoms
- Boiling point increases down the group, as atoms get larger and intermolecular forces increase
Part 6 — Bonding
There are three types, and each one is about what atoms do with their outer electrons to achieve a full shell.
| Bonding | Between | Electrons |
|---|---|---|
| Ionic | Metal + non-metal | Transferred |
| Covalent | Non-metal + non-metal | Shared |
| Metallic | Metal + metal | Delocalised |
Ionic bonding
Electrons are transferred from a metal to a non-metal. The metal loses electrons and becomes a positive ion. The non-metal gains electrons and becomes a negative ion. The oppositely charged ions attract by strong electrostatic forces.
Chlorine: 2,8,7 → gains one electron → Cl⁻ with 2,8,8
Both now have full outer shells.
Common ion charges:
| Group | Ion formed |
|---|---|
| 1 | 1+ |
| 2 | 2+ |
| 6 | 2− |
| 7 | 1− |
Properties of ionic compounds
| Property | Explanation |
|---|---|
| High melting and boiling points | Strong electrostatic forces between ions require a lot of energy to overcome |
| Do not conduct when solid | Ions are held in fixed positions and cannot move |
| Do conduct when molten or dissolved | Ions are free to move and carry charge |
| Often soluble in water | Water molecules can separate the ions |
| Brittle | Shifting the layers brings like charges together, and they repel |
Not "the bonds are strong."
All eight words. Mark schemes are specific about this one.
Covalent bonding
Electrons are shared between non-metal atoms. Each shared pair is one covalent bond.
| Molecule | Bonds |
|---|---|
| H₂ | One single bond |
| Cl₂ | One single bond |
| H₂O | Two single bonds |
| NH₃ | Three single bonds |
| CH₄ | Four single bonds |
| O₂ | One double bond |
| CO₂ | Two double bonds |
Simple molecular substances
Small molecules with strong covalent bonds inside them, but weak forces between them.
| Property | Explanation |
|---|---|
| Low melting and boiling points | The weak intermolecular forces between molecules are easily overcome — the covalent bonds are not broken |
| Do not conduct electricity | No free electrons and no ions |
| Often gases or liquids at room temperature | Little energy needed to separate molecules |
"Simple molecular substances have low melting points because the covalent bonds are weak."
Wrong. Covalent bonds are very strong. What's weak are the intermolecular forces between molecules — and those are what break when a substance melts. The covalent bonds inside each molecule stay intact.
Examiners flag this every year. Say "the intermolecular forces are weak" and never "the bonds are weak."
Metallic bonding
Metal atoms lose their outer electrons, which become delocalised. The structure is a lattice of positive metal ions in a sea of delocalised electrons, held together by strong electrostatic attraction.
| Property | Explanation |
|---|---|
| Good conductors of electricity | Delocalised electrons are free to move and carry charge |
| Good conductors of heat | Delocalised electrons transfer energy quickly |
| High melting points | Strong attraction between ions and delocalised electrons |
| Malleable and ductile | Layers of ions can slide over each other without breaking the bonding |
Alloys
An alloy is a mixture of a metal with another element. Alloys are harder than pure metals because the different-sized atoms distort the layers, so they cannot slide over each other easily.
| Alloy | Made from | Used for |
|---|---|---|
| Steel | Iron + carbon | Construction, tools |
| Bronze | Copper + tin | Statues, bearings |
| Brass | Copper + zinc | Instruments, fittings |
Part 7 — Giant covalent structures
Some covalent substances form giant structures rather than small molecules — every atom bonded to others in a continuous network.
Diamond
- Each carbon atom bonded to four others
- Very high melting point — many strong covalent bonds must be broken
- Very hard — used in cutting tools
- Does not conduct — no free electrons
Graphite
- Each carbon atom bonded to three others, in layers of hexagonal rings
- One delocalised electron per atom
- Conducts electricity — the delocalised electrons can move
- Soft and slippery — the layers have only weak forces between them and can slide, so it's used as a lubricant and in pencils
- Very high melting point — the covalent bonds within layers are strong
Same element. Completely different properties.
Diamond: four bonds per atom, no free electrons, hard, insulator. Graphite: three bonds per atom, one free electron, soft, conductor.
The difference is entirely structural. That's the point of the question.
Graphene
- A single layer of graphite, one atom thick
- Excellent conductor
- Extremely strong for its mass
- Potential uses in electronics and composites
Fullerenes
- Hollow molecules of carbon, usually in rings of five or six atoms
- Buckminsterfullerene (C₆₀) — a hollow sphere, the first discovered
- Nanotubes — cylindrical, very high length-to-diameter ratio
- Uses: drug delivery, lubricants, catalysts, reinforcing materials
Silicon dioxide
- Giant covalent structure, similar to diamond
- Very high melting point, hard, does not conduct
Exam-style questions
Answer all questions. Marks are shown in brackets. Total: 70 marks.
Question 1
An atom of magnesium has atomic number 12 and mass number 24.
(Total 10 marks)
Question 2
(Total 8 marks)
Question 3
Chlorine exists as two isotopes: 75% chlorine-35 and 25% chlorine-37.
(Total 5 marks)
Question 4
(Total 13 marks)
Question 5
(Total 9 marks)
Question 6
(Total 8 marks)
Question 7
(Total 8 marks)
Question 8
Diamond and graphite are both made only of carbon atoms.
(Total 9 marks)
TOTAL FOR PAPER: 70 MARKS
Mark scheme
Question 1
| (a) 3. Protons 12 | 1 |
| Electrons 12 | 1 |
| Neutrons 24 − 12 = 12 | 1 |
(b) 1 — 2,8,2.
| (c) 3. Group 2 — because it has 2 electrons in its outer shell | 1 + 1 |
| Period 3 — because it has 3 shells | 1 |
| (d) 3. Isotopes are atoms of the same element with the same number of protons | 1 |
| But different numbers of neutrons | 1 |
| Magnesium-25 has 25 − 12 = 13 neutrons | 1 |
Question 2
| (a) 2. A ball or sphere of positive charge | 1 |
| With negatively charged electrons embedded in it | 1 |
| (b) 4. Most passing straight through shows the atom is mostly empty space (observation linked + conclusion) | 2 |
| A few bouncing back shows there is a small, dense, positively charged nucleus containing most of the mass | 2 |
| (c) 2. New experimental evidence was obtained | 1 |
| Scientists modify or replace models when existing ones cannot explain new observations | 1 |
Question 3
| (a) 3. (35 × 75) + (37 × 25) | 1 |
| = 2625 + 925 = 3550 | 1 |
| 3550 ÷ 100 = 35.5 | 1 |
| (b) 2. Chemical properties depend on the number of electrons, particularly outer electrons | 1 |
| Isotopes have the same number of electrons, so react identically | 1 |
Question 4
| (a) 2. They have the same number of electrons in their outer shell | 1 |
| And chemical reactions involve the outer electrons | 1 |
| (b) 4. Group 1 metals react by losing their outer electron | 1 |
| Going down the group, atoms have more shells so the outer electron is further from the nucleus | 1 |
| It is also shielded by more inner shells | 1 |
| So there is less attraction from the nucleus and the electron is lost more easily, making the element more reactive | 1 |
| (c) 3. Group 7 elements react by gaining an electron | 1 |
| Going down the group, the outer shell is further from the nucleus and more shielded | 1 |
| So it is harder to attract an extra electron, making the element less reactive | 1 |
(d) 2 — potassium + water → potassium hydroxide + hydrogen (1 for reactants, 1 for products).
| (e) 2. They have full outer shells | 1 |
| So they do not need to gain, lose or share electrons | 1 |
Question 5
| (a) 3. Sodium loses one electron from its outer shell | 1 |
| Chlorine gains that electron | 1 |
| Both achieve full outer shells, forming Na⁺ and Cl⁻ | 1 |
| (b) 3. There is a giant ionic lattice | 1 |
| With strong electrostatic forces of attraction between oppositely charged ions | 1 |
| A large amount of energy is needed to overcome these forces | 1 |
| (c) 3. To conduct, charged particles must be free to move | 1 |
| When solid, the ions are held in fixed positions in the lattice | 1 |
| When molten or dissolved, the ions are free to move and carry charge | 1 |
Question 6
| (a) 2. A shared pair of electrons | 1 |
| Between two non-metal atoms | 1 |
| (b) 3. Methane is a simple molecular substance | 1 |
| The intermolecular forces between molecules are weak | 1 |
| Little energy is needed to overcome them, so the boiling point is low | 1 |
| (c) 3. Covalent bonds are strong, not weak | 1 |
| When a simple molecular substance melts, the covalent bonds are not broken | 1 |
| It is the weak intermolecular forces between molecules that are overcome | 1 |
Question 7
| (a) 3. A giant structure / lattice of positive metal ions | 1 |
| In a sea of delocalised electrons | 1 |
| Held together by strong electrostatic attraction between the ions and the electrons | 1 |
| (b) 2. The delocalised electrons are free to move through the structure | 1 |
| And carry charge | 1 |
| (c) 3. Pure metals have layers of atoms that can slide over each other | 1 |
| Alloys contain atoms of different sizes | 1 |
| These distort the layers so they cannot slide as easily, making the alloy harder | 1 |
Question 8
| (a) 3. Each carbon atom is covalently bonded to four others | 1 |
| Forming a giant covalent structure / rigid three-dimensional lattice | 1 |
| Many strong covalent bonds must be broken to break it | 1 |
| (b) 3. Each carbon is bonded to three others, forming layers | 1 |
| There are only weak forces between the layers | 1 |
| So the layers can slide over each other easily | 1 |
| (c) 3. In graphite each carbon forms three bonds, leaving one delocalised electron per atom | 1 |
| These electrons are free to move and carry charge | 1 |
| In diamond each carbon forms four bonds, so there are no delocalised electrons and it cannot conduct | 1 |
Common mistakes in this topic
"The covalent bonds are weak." They are very strong. The intermolecular forces are weak. This is the single most penalised error in GCSE Chemistry.
"The bonds are strong" for ionic compounds. Say "strong electrostatic forces of attraction between oppositely charged ions."
Saying ionic compounds conduct because "electrons move." They conduct because ions move. Ionic compounds have no free electrons.
Confusing atomic number and mass number. Atomic number is protons. Mass number is protons plus neutrons.
Saying isotopes have different chemical properties. They have identical chemistry — only mass differs.
Explaining group trends without mentioning shielding and distance. "It's more reactive because it's bigger" scores nothing. Say further from the nucleus and more shielding.
Forgetting units and state symbols where a question asks for them.
Teaching notes
Electronic structure is the key that unlocks the periodic table. Once students see that group number = outer electrons and period number = number of shells, the table stops being a chart to memorise and becomes a map that explains itself. Teach it before the groups, not after.
The two group trends should be taught together, not separately. Group 1 gets more reactive down the group; Group 7 gets less. Same cause — distance and shielding — opposite outcome. Taught together it's one idea; taught apart it's two facts to confuse.
The "weak covalent bonds" error must be attacked directly. It is intuitive, it is wrong, and it is worth marks on almost every paper. Say it out loud as a class: the bonds are strong, the forces between molecules are weak.
Diamond and graphite is the best structure-and-properties question available. Same element, opposite properties, entirely explained by bonding. If a student can do this comparison they have understood the whole topic.
Alpha scattering rewards the observation–conclusion pairing. Teach it as three pairs, not six facts.
The exact wording matters more in Chemistry than in most subjects. "Strong electrostatic forces of attraction between oppositely charged ions" is eight words that appear verbatim in mark schemes. Make students write the full phrase every time until it's automatic.
Written for GCSE Chemistry, all major UK boards. Check your specification — fullerenes, graphene and nanoparticle detail vary between boards, and some content is Higher Tier only.