Atomic Structure & Periodic Table

Everything around you — the air you breathe, the chair you sit on, your own body — is built from atoms. An atom is the smallest particle of an element that still behaves like that element. Atoms are tiny: their radius is approximately $1 \times 10^{-10}\,\text{m}$ (one ten-billionth of a metre). A row of 10 million atoms placed side-by-side would still only stretch 1 mm.

An element is a substance made of only one type of atom. The 118 known elements are listed in the periodic table, each represented by a one- or two-letter chemical symbol. The first letter is always a capital; the second (if there is one) is always lowercase. AQA mark schemes penalise sloppy capitalisation — "CO" means carbon and oxygen (i.e. carbon monoxide), but "Co" is the element cobalt.

Examples of element symbols you must know:

Element	Symbol	Element	Symbol
Hydrogen	H	Sodium	Na
Carbon	C	Magnesium	Mg
Nitrogen	N	Aluminium	Al
Oxygen	O	Iron	Fe
Sulfur	S	Copper	Cu
Chlorine	Cl	Zinc	Zn
Calcium	Ca	Silver	Ag
Potassium	K	Gold	Au

Some symbols look unrelated to the English name — these come from the original Latin (e.g. Fe = ferrum for iron; Na = natrium for sodium; Au = aurum for gold). British GCSE students are expected to recognise these.

Why so few elements? Although there are only about 90 naturally occurring elements, their atoms can combine in countless arrangements. Just six elements (carbon, hydrogen, oxygen, nitrogen, calcium and phosphorus) make up over 99% of the human body — yet the variety of molecules they form gives rise to all of life.

When atoms of two or more different elements combine chemically, they form a compound. The atoms are held together by chemical bonds (covered in topic 5.2). Crucially:

• A compound has fixed proportions — water is always 2 hydrogen atoms to 1 oxygen atom, never 3:1 or 1:1.
• A compound has different properties from its starting elements — hydrogen is an explosive gas, oxygen helps things burn, but water (H₂O) puts fires out.
• A compound can only be separated back into its elements by a chemical reaction — physical methods like filtering or distilling will not split it.

Reading a chemical formula. The formula uses element symbols with small subscript numbers:

• $\text{H}_2\text{O}$ — 2 hydrogen atoms + 1 oxygen atom per molecule of water.
• $\text{CO}_2$ — 1 carbon atom + 2 oxygen atoms per molecule of carbon dioxide.
• $\text{NH}_3$ — 1 nitrogen atom + 3 hydrogen atoms per molecule of ammonia.
• $\text{H}_2\text{SO}_4$ — 2 H + 1 S + 4 O per molecule of sulfuric acid.

Brackets in formulae. Brackets mean "everything inside, multiplied by the subscript outside":

• $\text{Ca(OH)}_2$ — 1 Ca + 2 O + 2 H (calcium hydroxide).
• $\text{Mg(NO}_3)_2$ — 1 Mg + 2 N + 6 O (magnesium nitrate; the 3 inside × 2 outside = 6 oxygen atoms).

Counting atoms inside brackets is one of the highest-frequency exam questions for this topic. Practise until it's second nature.

Chemists use equations to describe reactions. There are two kinds you need:

Word equations — use the names of the substances. Reactants on the left, products on the right, separated by an arrow (→). The plus sign (+) means "and".

sodium + chlorine → sodium chloride

magnesium + oxygen → magnesium oxide

methane + oxygen → carbon dioxide + water

The arrow is read "produces" or "reacts to form". It is never an equals sign — chemical reactions are one-way changes (unless they are reversible, in which case ⇌ is used; see topic 5.6).

Symbol equations — use chemical formulae. They must show the same number of each kind of atom on both sides, because atoms are not created or destroyed in a chemical reaction (the law of conservation of mass). Compare:

Word	Unbalanced symbol	Balanced symbol
hydrogen + oxygen → water	H₂ + O₂ → H₂O	2H₂ + O₂ → 2H₂O
methane + oxygen → CO₂ + water	CH₄ + O₂ → CO₂ + H₂O	CH₄ + 2O₂ → CO₂ + 2H₂O
sodium + water → NaOH + H₂	Na + H₂O → NaOH + H₂	2Na + 2H₂O → 2NaOH + H₂

How to balance an equation (step-by-step).

1. Write the correct formulae (do NOT change subscripts — only change the big numbers in front).
2. Count atoms on each side.
3. Add multipliers (coefficients) in front of formulae to even up.
4. Re-count after each change. Tackle the most "complicated" element first; leave hydrogen and oxygen until last.
5. Add state symbols if asked: (s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in water).

Worked example. Balance: $\text{Al} + \text{O}_2 \to \text{Al}_2\text{O}_3$.

• Aluminium: 1 on left, 2 on right → put 2 in front of Al: $2\text{Al} + \text{O}_2 \to \text{Al}_2\text{O}_3$.
• Oxygen: 2 on left, 3 on right. Lowest common multiple is 6 → put 3 in front of O₂ and 2 in front of Al₂O₃: $2\text{Al} + 3\text{O}_2 \to 2\text{Al}_2\text{O}_3$.
• Re-check Al: now 4 on the right, only 2 on the left → put 4 in front of Al: $4\text{Al} + 3\text{O}_2 \to 2\text{Al}_2\text{O}_3$.
• Final check — 4 Al on both sides, 6 O on both sides. Balanced.

It is essential to distinguish chemical from physical changes.

Chemical change (chemical reaction) — one or more new substances are formed; usually difficult to reverse; energy change (heat in/out, light, sound, gas given off, colour change, precipitate forming). Examples: burning, rusting, cooking, neutralisation, photosynthesis.

Physical change — no new substance; usually easy to reverse; only the form of the substance changes (state, shape, dissolved/undissolved). Examples: melting, boiling, dissolving, evaporating, freezing, cutting, mixing sand and gravel.

Quick test. Ask yourself: if I undid this change, would I get the original substance back without using a chemical reaction? If yes → physical. If no → chemical.

Change	Type	Why
Ice melting	Physical	Still water; freezing reverses it.
Iron rusting	Chemical	New substance (iron oxide); cannot reverse by warming.
Salt dissolving in water	Physical	Can evaporate water to get salt back.
Wood burning	Chemical	Forms CO₂, H₂O, ash — can't 'unburn'.
Magnesium ribbon burning bright white	Chemical	Forms magnesium oxide (white powder); energy released as light.
Wax melting on a candle	Physical	Liquid wax becomes solid wax again on cooling.

Note: dissolving feels chemical (it looks like it disappears) but is actually physical — the salt is still there, just spread between water molecules. You can recover it by evaporation.

A mixture is two or more elements or compounds not chemically combined together. Important features:

• The substances retain their individual chemical properties (salt is still salt, sand is still sand).
• The proportions can vary (air can be 21% or 20.5% oxygen).
• Components can be separated by physical methods (no chemical reaction needed).
• No energy change accompanies the formation of a mixture.

Compound versus mixture — the contrast:

Property	Compound	Mixture
Bonding	Chemically bonded	Not bonded
Composition	Fixed (e.g. H₂O is always 2:1)	Variable (air can be 21–20% O₂)
Properties	Different from elements (water ≠ H + O)	Same as components (salty + sandy water tastes salty and looks sandy)
Separation	Chemical reaction only	Physical methods
Energy change when formed	Yes (usually exothermic)	No

Examples of mixtures around you:

• Air — mostly N₂ and O₂, with Ar, CO₂, water vapour, traces of other gases.
• Sea water — water + dissolved sodium chloride and other salts.
• Brass — a mixture of copper and zinc (this kind of metal mixture is called an alloy).
• Crude oil — a mixture of hundreds of hydrocarbons.
• Blood — water, dissolved salts, proteins, suspended cells.

The choice of separation technique depends on the physical state of the components, whether they are soluble, and (for liquids) their boiling points.

Filtration separates an insoluble solid from a liquid. The mixture is poured through a folded filter paper in a funnel. The solid is trapped (the residue) and the liquid passes through (the filtrate).

Examples:

• Separating sand from sandy water.
• Separating tea leaves from brewed tea.
• Removing copper(II) carbonate from a reaction mixture.

Crystallisation is used to recover a soluble solid from its solution. The solvent (often water) is evaporated, leaving the solute behind as solid crystals. Two ways:

1. Slow crystallisation — leave the solution in a warm place to evaporate slowly. Larger, purer crystals form.
2. Heated crystallisation — gently warm the solution in an evaporating basin over a water bath. Stop when the solution is saturated (a crystal floats on top, or you see crystals forming on a glass rod dipped in). Allow the remainder to crystallise as it cools.

Slow crystallisation is required if the solute decomposes when strongly heated (e.g. copper sulfate would lose water of crystallisation if boiled).

Simple distillation separates a liquid from a solution (e.g. pure water from salt water). The solution is heated; the water evaporates, passes into a condenser, cools to liquid water again and is collected as the distillate. The salt remains in the flask.

Method:

1. Place solution in a flask with anti-bumping granules.
2. Heat using a Bunsen or heating mantle.
3. Steam rises into a Liebig condenser (cold water flowing in the outer jacket — water in at the bottom, out at the top, to maintain a temperature gradient).
4. Condensed liquid drips into a collecting beaker.
5. The thermometer at the top of the flask shows the boiling point of whatever is currently vaporising.

Fractional distillation is used when you have two (or more) miscible liquids with different boiling points. Examples: ethanol from water (used in distilleries; bp 78 °C vs 100 °C); crude oil fractions (covered in topic 5.7).

A fractionating column is added between the flask and the condenser. It is filled with glass beads or rings; vapour rises and partially condenses at different levels. Liquids with lower boiling points reach the top first.

Worked example. A mixture of ethanol (bp 78 °C) and water (bp 100 °C):

• Heat to ~78 °C: ethanol vaporises and reaches the top of the column; water vapour condenses back into the flask.
• Ethanol vapour enters the condenser, becomes liquid and is collected.
• When all ethanol is gone, the thermometer reading rises sharply towards 100 °C.

Paper chromatography separates a mixture of soluble, coloured substances (e.g. food dyes, ink pigments). It works because the substances move at different rates with the solvent across the paper.

Apparatus and method:

1. Draw a horizontal pencil baseline ~1 cm from the bottom of a strip of chromatography paper. (Pencil — because pen ink would dissolve.)
2. Place a tiny spot of each test sample on the baseline; leave to dry; repeat 2–3 times for a concentrated spot.
3. Stand the paper in a beaker containing a thin layer of solvent (water for water-soluble dyes; ethanol or propanone for non-polar dyes). The solvent must be BELOW the pencil line.
4. Cover the beaker (so the solvent doesn't evaporate).
5. The solvent moves up the paper by capillary action, carrying the dyes with it.
6. When the solvent reaches near the top, remove the paper and immediately mark the solvent front in pencil.
7. Allow the paper to dry. Each substance has moved a different distance.

The Rᶠ value (retention factor) tells you how far a component moved relative to the solvent:

$R_f = \frac{\text{distance moved by spot}}{\text{distance moved by solvent front}}$

Both distances are measured from the original baseline. Rᶠ is always between 0 and 1 and has no units. A given substance always has the same Rᶠ in the same solvent — so Rᶠ is used to identify components by comparing with known reference values.

Worked example. A dye spot moves 3.6 cm; the solvent front moves 8.0 cm. What is the Rᶠ?

$R_f = \frac{3.6}{8.0} = 0.45$

Note: the stationary phase is the paper itself; the mobile phase is the solvent that travels up the paper. A pure substance shows only ONE spot in any solvent — a useful test for purity.

John Dalton (1803) proposed that all matter is made of tiny indivisible solid spheres called atoms. Each element had its own kind of atom; atoms of different elements differed in mass and properties. This is sometimes called the billiard ball model.

J. J. Thomson (1897, Cambridge) discovered the electron by experimenting with cathode rays in vacuum tubes. He showed that electrons were:

• Negatively charged.
• Much smaller than any atom (~1/2000 the mass of a hydrogen atom).
• Identical regardless of which gas was in the tube — so present in all atoms.

This contradicted Dalton's "indivisible" claim. Thomson proposed a new model: the plum-pudding model. The atom was a sphere of positive charge (the 'pudding') with negative electrons (the 'plums') embedded throughout.

Ernest Rutherford (1909, working in Manchester with Geiger and Marsden) fired a beam of alpha particles (positively charged, relatively heavy) at a thin sheet of gold foil. According to the plum-pudding model, the αs should pass through with only tiny deflections, because the positive charge was spread thinly throughout the atom.

Observations:

1. Most α particles passed straight through.
2. A small number were deflected by large angles.
3. A very few (~1 in 8000) bounced straight back.

Conclusions:

1. The atom is mostly empty space (most particles unaffected).
2. There is a small region of concentrated positive charge — the nucleus — strong enough to deflect a fast α.
3. The nucleus contains most of the mass of the atom (only mass could bounce the α back).

This was the birth of the nuclear model: tiny dense positive nucleus, mostly empty space, electrons distributed somewhere outside.

Niels Bohr (1913) refined the nuclear model. Calculations showed Rutherford's electrons would spiral into the nucleus, but real atoms are stable. Bohr proposed that electrons orbit at fixed distances (energy levels or shells) and can only jump between them by absorbing or emitting specific amounts of energy. This explains why atoms emit light at characteristic frequencies (the basis of flame tests and spectroscopy).

Later, James Chadwick (1932) discovered the neutron — an uncharged particle in the nucleus that explained why isotopes (atoms of the same element with different masses) exist. This completed the modern picture: nucleus = protons + neutrons; electrons in shells.

Every atom is built from three kinds of subatomic particles:

Particle	Symbol	Relative charge	Relative mass	Location
Proton	p	+1	1	Nucleus
Neutron	n	0	1	Nucleus
Electron	e⁻	−1	~1/2000 (≈ 0)	Shells around nucleus

The relative masses are not in grams — they are relative to the mass of a proton. The actual mass of a proton is about $1.67 \times 10^{-27}$ kg; AQA does not require you to remember the exact value, but you must know that protons and neutrons have the same mass and electrons are much lighter.

Why atoms are neutral. In any neutral atom:

$\text{number of protons} = \text{number of electrons}$

Equal numbers of +1 and −1 charges cancel exactly. So the overall charge of an atom is 0.

Atomic number (Z) = number of protons. This defines the element. Every carbon atom has 6 protons; every oxygen has 8.

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

The standard way to write a nuclide:

$_Z^A X$

For example, $_6^{12}\text{C}$ is a carbon atom with 6 protons, 12 − 6 = 6 neutrons and 6 electrons (in the neutral atom).

Ions. Atoms can lose or gain electrons to form charged particles called ions:

• Lose 1 electron → +1 cation (e.g. Na⁺).
• Lose 2 electrons → +2 cation (e.g. Mg²⁺).
• Gain 1 electron → −1 anion (e.g. Cl⁻).
• Gain 2 electrons → −2 anion (e.g. O²⁻).

The number of protons and neutrons doesn't change when an ion forms — only the electron count.

Atomic radius is about $1 \times 10^{-10}\,\text{m}$. To picture this: a row of 10 million atoms placed side-by-side stretches just 1 mm. Different elements vary a little — hydrogen is smaller ($0.5 \times 10^{-10}$ m); caesium is larger ($2.6 \times 10^{-10}$ m) — but $10^{-10}$ m is the AQA rule of thumb.

Nuclear radius is about $1 \times 10^{-14}\,\text{m}$ — that's $\tfrac{1}{10{,}000}$ of the atomic radius.

To get a sense of the difference: imagine an atom enlarged to the size of Wembley Stadium (~200 m across). The nucleus, at the centre of the pitch, would be about the size of a marble.

Despite being so small, the nucleus contains almost all of the atom's mass — because protons and neutrons are ~2000× heavier than electrons (recall 5.1.1.4).

An element is defined by its proton (atomic) number. But atoms of the same element can have different numbers of neutrons — these versions are called isotopes.

Isotope	Protons	Neutrons	Electrons
¹H (protium)	1	0	1
²H (deuterium)	1	1	1
³H (tritium)	1	2	1

Isotope	Protons	Neutrons	Mass #
³⁵Cl	17	18	35
³⁷Cl	17	20	37

Key features of isotopes:

• Same chemical properties (same electron arrangement → reactions identical).
• Slightly different physical properties (density, melting/boiling point — heavier atoms move more slowly).
• Naturally occur in fixed proportions for a given element on Earth.

Definition (AQA wording): the relative atomic mass (Aᵣ) of an element is the average mass of the atoms of that element, taking into account the relative abundances of each isotope, compared with $\tfrac{1}{12}$ the mass of a ¹²C atom.

The standard formula:

$A_r = \frac{\sum (\text{isotope mass} \times \text{\% abundance})}{100}$

Worked example — chlorine. Chlorine has two isotopes: ³⁵Cl (75%) and ³⁷Cl (25%).

$A_r = \frac{(35 \times 75) + (37 \times 25)}{100} = \frac{2625 + 925}{100} = \frac{3550}{100} = 35.5$

That matches the value shown on the periodic table — non-whole because of the isotope mixture.

Worked example — copper. Copper has two isotopes: ⁶³Cu (69.2%) and ⁶⁵Cu (30.8%).

$A_r = \frac{(63 \times 69.2) + (65 \times 30.8)}{100} = \frac{4359.6 + 2002.0}{100} = 63.6$

Round to 1 d.p. — the periodic-table value is 63.5.

Each electron in an atom occupies one of a series of shells (energy levels). The shells fill from the closest to the nucleus outwards. AQA expects you to know the first three shells:

Shell number	Maximum electrons
1 (innermost)	2
2	8
3	8 (for AQA GCSE up to Ca)

So the maximum number of electrons before a new shell is needed is $2 + 8 + 8 = 18$ — which is why the third row of the periodic table contains the elements sodium (Na) through argon (Ar).

Writing electronic structures. Use numbers separated by commas, starting from the innermost shell.

Element	Z	Structure
H	1	1
He	2	2
Li	3	2,1
C	6	2,4
N	7	2,5
O	8	2,6
Ne	10	2,8
Na	11	2,8,1
Mg	12	2,8,2
Al	13	2,8,3
Si	14	2,8,4
P	15	2,8,5
S	16	2,8,6
Cl	17	2,8,7
Ar	18	2,8,8
K	19	2,8,8,1
Ca	20	2,8,8,2

Drawing dot-shell diagrams. Draw the nucleus in the centre, then concentric circles for the shells. Place dots (or crosses) for each electron on the appropriate shell. Spread electrons evenly around the shell — first put one in each of the top/bottom/left/right positions, then pair them up.

Linking to the periodic table:

• Group number = number of outer-shell electrons (for main-group elements 1–7).
• Period number = number of occupied shells.
• Group 0 (noble gases): full outer shell → very unreactive.

Example: chlorine (2,8,7) — 7 outer electrons → group 7; 3 shells → period 3.

The modern periodic table lists the 118 known elements in order of increasing atomic (proton) number. The arrangement is not by atomic mass (although the two orders nearly always coincide).

Groups (vertical columns). There are 8 main groups, labelled 1, 2, then a gap for the transition metals, then 3, 4, 5, 6, 7, 0. Elements in the same group share:

• The same number of outer-shell electrons (for groups 1–7).
• Similar chemical properties (because reactivity depends on the outer shell).

Periods (horizontal rows). There are 7 periods. Within a period, the atomic number increases by 1 from left to right. Each new period starts when a new electron shell begins to fill.

Group	Outer-shell electrons	Family name	Example
1	1	Alkali metals	Sodium (Na)
2	2	Alkaline earth metals	Magnesium (Mg)
3	3	(Boron group)	Aluminium (Al)
4	4	(Carbon group)	Carbon (C)
5	5	(Nitrogen group)	Nitrogen (N)
6	6	(Oxygen group)	Oxygen (O)
7	7	Halogens	Chlorine (Cl)
0	8 (full; 2 for He)	Noble gases	Argon (Ar)

Metals on the left, non-metals on the right. A 'staircase' line (running roughly from boron down to polonium) separates them. Elements right next to the line (e.g. silicon, germanium) are metalloids — they share properties of both metals and non-metals.

Reading an element box. The box for each element typically shows:

• The element symbol (large, in the centre).
• The atomic number (smaller, bottom-left).
• The relative atomic mass (smaller, top-left).

Example for sodium:

23
Na
11

Na — symbol; 11 — atomic number; 23 — Aᵣ.

Before the modern atomic theory, chemists tried to find an order in the ~60 elements known by the 1860s.

John Newlands (1864, London) — arranged elements in order of atomic mass. He noticed that every 8th element had similar properties — he called this the Law of Octaves (by analogy with musical octaves). The arrangement worked for the lightest elements (e.g. Li, Na, K appeared as octaves) but failed for heavier elements: when transition metals appeared, they didn't fit and Newlands' system grouped chemically very different elements together. Most chemists at the time rejected his work.

Dmitri Mendeleev (1869, Russia) — also ordered by atomic mass, but with two key insights:

1. He left gaps for undiscovered elements. When the next element by mass didn't fit the chemical pattern of a group, Mendeleev assumed a heavier unknown element belonged in that slot and the lighter one belonged to the next group along.
2. He reversed some pairs. When ordering strictly by mass put elements in the wrong group based on properties, Mendeleev swapped them — e.g. tellurium (Te, mass 128) and iodine (I, mass 127). He believed atomic masses were sometimes inaccurate.

Mendeleev used his table to predict the properties of missing elements. He left gaps below silicon (predicting an element he called eka-silicon) and below aluminium (eka-aluminium). When germanium (1886) and gallium (1875) were later isolated, their properties closely matched Mendeleev's predictions. This was a striking confirmation of his arrangement.

Why does the modern table avoid these issues? Atomic numbers (protons) — not atomic masses — order the table. This was made possible by the discovery of protons (Rutherford, ~1917) and isotopes. Isotopes explain why some mass orderings disagree with chemical groupings: argon (Aᵣ 40) has more neutrons on average than potassium (Aᵣ 39), so by mass alone Ar would come after K, but by proton count Ar (18) correctly comes before K (19).

Metals (left of the staircase):

Property	Typical value
Appearance	Shiny when polished
State at room temperature	Solid (except mercury, liquid)
Malleable & ductile?	Yes (can be bent and drawn into wires)
Conductor of electricity?	Yes (in solid and molten state)
Conductor of heat?	Yes
Melting/boiling point	Usually high (Na 98 °C — but most much higher; Fe 1538 °C)
Density	Usually high (Fe 7.9 g/cm³; gold 19.3 g/cm³)
Ion formation	Lose electrons → positive ions (cations)

Non-metals (right of the staircase):

Property	Typical value
Appearance	Dull (sometimes coloured: yellow sulfur, green Cl₂, brown Br₂)
State at room temperature	Solid, liquid (Br₂) or gas (H₂, N₂, O₂, halogens, noble gases)
Malleable?	No — solid non-metals are brittle
Conductor?	No (except graphite — a special form of carbon)
Melting/boiling point	Often low
Density	Usually low
Ion formation	Gain electrons → negative ions (anions) — OR share electrons covalently

Why the difference? Metals have few outer-shell electrons (1, 2 or 3) and lose them easily to achieve a full outer shell. Non-metals have many outer-shell electrons (5, 6 or 7) and find it easier to gain a few more to reach a full shell, rather than losing them all.

Predicting ion charges from group:

Group	Outer electrons	Ion charge	Example
1	1	+1	Na⁺
2	2	+2	Ca²⁺
3	3	+3	Al³⁺
5	5	−3	N³⁻
6	6	−2	O²⁻
7	7	−1	Cl⁻
0	8 (full)	none	Ar (no ion)