HookThe experiment where the bullets bounced back
In 1909, in a laboratory at the University of Manchester, Hans Geiger and Ernest Marsden fired a beam of alpha particles at a sheet of gold foil only a few thousand atoms thick. Ernest Rutherford expected them all to sail straight through, because the accepted picture of the atom — J. J. Thomson's "plum pudding" — imagined positive charge smeared thinly through the whole atom, far too diffuse to deflect anything. Almost all of them did pass through. But roughly one alpha particle in every eight thousand bounced sharply back towards the source. Rutherford later said it was "as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you." The only way to explain it was to abandon the pudding: the atom's positive charge and nearly all its mass had to be crammed into a minute, dense core — the nucleus — surrounded by electrons and empty space.
That single result is the foundation the whole of C1 is built on. An atom is almost entirely empty space with a tiny central nucleus of protons and neutrons, and it is the number of protons that decides which element you are looking at and where it sits in the periodic table. The electrons arranged in shells around that nucleus decide how the atom bonds and reacts. Learn to read an atom — its protons, neutrons, electrons and their arrangement — and the periodic table stops being a wall of symbols and becomes what Dmitri Mendeleev built it to be: a map on which an element's position tells you, in advance, how it will behave.
ModelWhat an atom is — and the three particles inside it
An element is a substance made of only one type of atom; there are about 100 of them, each with a one- or two-letter symbol on the periodic table. A compound is two or more elements chemically bonded together in a fixed ratio, and it can only be split apart by chemical reactions or electrolysis, never by simply picking it apart. Chemical bonding always involves the electrons: they are transferred or shared when atoms join.
Every atom is built from three particles. Protons and neutrons sit in the central nucleus; electrons occupy shells around it. Their relative charges are the numbers you must know cold: a proton is \(+1\), an electron is \(-1\), and a neutron is \(0\). Their relative masses matter too — a proton and a neutron each count as \(1\), while an electron is about \(\frac{1}{1836}\) of that, so small it is treated as negligible. An atom is neutral overall because it has equal numbers of protons and electrons. The atomic (proton) number is the number of protons, and the mass number is the number of protons plus neutrons.
Atoms are astonishingly small. A typical atomic radius is about \(0.1\) nm, or \(1\times10^{-10}\) m, while the nucleus has a radius of roughly \(1\times10^{-14}\) m — about one ten-thousandth of the atom, yet it holds virtually all the mass. If an atom were the size of a football stadium, the nucleus would be a pea on the centre spot.
Take sodium, written \(^{23}_{11}\text{Na}\). The bottom number is the atomic number, \(11\), so every sodium atom has \(11\) protons — that is what makes it sodium and nothing else. Because a neutral atom has equal protons and electrons, it also has \(11\) electrons. The top number is the mass number, \(23\), the total of protons plus neutrons, so the number of neutrons is \(23 - 11 = 12\). Three facts — 11 protons, 11 electrons, 12 neutrons — read straight off two digits. Change the proton number and you change the element; change only the neutron number and you have an isotope of the same element.
MechanismPulling mixtures apart — four separation techniques
A mixture is two or more substances that are not chemically bonded; each keeps its own properties, so a mixture can be separated by physical means without any chemical reaction. The technique you choose depends on which physical difference you can exploit.
Filtration separates an insoluble solid from a liquid: the solid stays on the filter paper as the residue, the liquid passes through as the filtrate — sand from water, for instance. Crystallisation recovers a soluble solid from its solution by evaporating the solvent slowly so that crystals grow — this is how copper sulfate crystals are made. Distillation separates a liquid from a solution (simple distillation) or two mixed liquids with different boiling points (fractional distillation) by evaporating and then condensing the vapour. Chromatography separates dissolved substances that travel at different speeds through paper because they have different solubilities — the technique behind separating the dyes in an ink.
Rock salt is a mixture of salt (soluble) and sand (insoluble), so no single technique will do — you chain two. First add water and stir: the salt dissolves, the sand does not. Filter the mixture: the sand is trapped as the residue, and the salt solution runs through as the filtrate. Finally, gently evaporate the filtrate so the water leaves and the salt crystallises out. Two techniques, each chosen for a specific physical difference — solubility, then volatility — which is exactly the reasoning a required-practical question is testing.
CaseFrom billiard balls to shells — how the model of the atom grew
The atom you have just met was not handed down whole; it was rebuilt four times as new evidence arrived. John Dalton, in the early 1800s, pictured atoms as tiny solid spheres that could not be divided, different for each element. Then in 1897 J. J. Thomson discovered the electron and proposed the plum-pudding model: a ball of positive charge with negative electrons dotted through it. Rutherford's gold-foil scattering (1909–1911) demolished that and gave us the nuclear model — a tiny, dense, positive nucleus with electrons around it.
Two refinements followed. Niels Bohr (1913) showed that electrons orbit the nucleus in fixed energy levels, or shells, at set distances — this fixed a fatal flaw, because electrons in random orbits should have spiralled into the nucleus. Later work established that the nucleus itself contains positively charged protons, and in 1932 James Chadwick provided the evidence for the neutron. The pattern is the point examiners test: each change was forced by experimental evidence, not opinion — cathode rays, scattered alpha particles, atomic spectra — which is how a scientific model earns the right to replace the one before it.
ModelIsotopes and relative atomic mass
Isotopes are atoms of the same element — so the same number of protons — but with different numbers of neutrons, which means the same atomic number and different mass numbers. Chlorine has chlorine-35 and chlorine-37; carbon has carbon-12 and carbon-14. Because isotopes have identical electron arrangements, their chemistry is identical; only physical properties that depend on mass, such as density, differ slightly.
A real sample of an element is a mixture of its isotopes, so we quote a weighted average called the relative atomic mass \(A_r\): the average mass of the atoms, weighted by how abundant each isotope is, on a scale where a carbon-12 atom is exactly 12. The formula is \(A_r = \dfrac{\sum(\text{isotope mass}\times\text{abundance})}{\text{total abundance}}\). This is why chlorine's relative atomic mass is 35.5 rather than a whole number — the average lands between the two isotopes.
Chlorine is \(75\%\) chlorine-35 and \(25\%\) chlorine-37. Weight each mass by its abundance, add, and divide by the total abundance: \[ A_r = \frac{(75 \times 35) + (25 \times 37)}{100} = \frac{2625 + 925}{100} = \frac{3550}{100} = 35.5 \] The classic mistake is to average 35 and 37 to get 36. That ignores that chlorine-35 is three times more common, which drags the true average down towards 35 — giving 35.5. Always multiply each isotope mass by its abundance before dividing.
ModelElectronic structure — filling the shells
Electrons occupy shells (energy levels) around the nucleus, filling the lowest-energy shell nearest the nucleus first. At GCSE the capacities are: 2 electrons in the first shell, then up to 8 in the second, then 8 in the third. We write the arrangement as a string of numbers — sodium, with 11 electrons, is \(2,8,1\).
This arrangement is the engine of the periodic table. The number of electrons in the outer shell equals the group number (for the main groups), and the number of occupied shells equals the period. A full outer shell is very stable, which is why the noble gases are unreactive; every other element reacts by losing, gaining or sharing electrons to reach a full outer shell. So the electron arrangement predicts both where an element sits and how it will behave.
Chlorine has 17 electrons, arranged \(2,8,7\). Three shells are occupied, so chlorine is in period 3; there are 7 electrons in the outer shell, so it is in group 7. That outer shell is one electron short of a full set of 8, which tells you chlorine's whole personality: it readily gains a single electron to complete the shell, forming a \(-1\) ion, \(\text{Cl}^-\). Reading the configuration has told you the position and the chemistry at once.
CaseBuilding the table — Newlands, Mendeleev and the gaps
The modern periodic table lists the elements in order of atomic (proton) number, arranged in periods (rows) and groups (columns). Elements in the same group have the same number of outer-shell electrons, which is why they react in similar ways — the table is organised around chemistry, not coincidence.
Getting there took decades. In 1864 John Newlands spotted that every eighth element repeated similar properties — his "law of octaves" — but he forced elements into rows with no gaps, so once the pattern broke down he was ridiculed. Mendeleev's 1869 table succeeded because of two bold moves: he left gaps for elements not yet discovered, and he even swapped a few pairs out of strict mass order so that they fell into the group whose properties they matched. Crucially, he used those gaps to predict the properties of missing elements. The order of the elements only truly made sense once Henry Moseley (1913) showed that they should be arranged by atomic number, not mass — which explained Mendeleev's swaps, because isotopes make atomic mass an imperfect guide.
Mendeleev left a gap below silicon and predicted an element he called "eka-silicon": a grey metal with a relative atomic mass of about 72 that would form an oxide of formula \(\text{XO}_2\). In 1886 germanium was discovered, with a relative atomic mass of 72.6 and an oxide \(\text{GeO}_2\) — almost exactly as predicted. A table that can forecast the properties of an element nobody has yet found is doing far more than listing; that single confirmed prediction is what convinced chemists Mendeleev was right.
MechanismGroups and trends — why reactivity runs the way it does
Metals sit on the left and centre of the table, non-metals in the top right, split by a diagonal staircase. Metals lose electrons to form positive ions; non-metals gain electrons to form negative ions, or share them. How readily an atom does this depends on how far its outer shell is from the nucleus and how strongly the nucleus can hold or attract electrons.
Group 1, the alkali metals (lithium, sodium, potassium…), each have one outer electron, lost easily to form \(+1\) ions; they react with water to give a metal hydroxide plus hydrogen. Reactivity increases down the group, because the single outer electron is further from the nucleus and better shielded, so it is lost more easily — lithium fizzes, sodium skates across the surface, potassium ignites with a lilac flame.
Group 7, the halogens (fluorine, chlorine, bromine, iodine), each have seven outer electrons and gain one to form \(-1\) ions; they exist as diatomic molecules such as \(\text{Cl}_2\). Reactivity decreases down the group, because the incoming electron is added to a shell further from the nucleus and more shielded, so it is harder to attract. A more reactive halogen will displace a less reactive one from a solution of its salt. Group 0, the noble gases, have full outer shells, so they neither gain nor lose electrons and are almost completely unreactive; their boiling points rise down the group as the atoms get larger.
Add chlorine water to a colourless solution of potassium bromide and it turns orange. Chlorine is above bromine in group 7, so it is more reactive and takes the place of bromine: \[ \text{Cl}_2 + 2\text{KBr} \rightarrow 2\text{KCl} + \text{Br}_2 \] The orange colour is the displaced bromine. Because the more reactive halogen displaces the less reactive one, the reaction itself ranks them — direct evidence that reactivity falls as you go down group 7, the opposite of group 1.
DataTransition metals — everything Group 1 is not
The block of metals in the centre of the table, the transition metals, are the everyday metals — iron, copper, nickel, chromium. Compared with the group 1 alkali metals they are almost opposites: much harder and stronger, far denser, with higher melting points, and much less reactive, reacting only slowly, if at all, with water.
They also share a set of typical properties the alkali metals lack. Many are useful catalysts — iron in the Haber process, nickel in hardening vegetable oils. They form coloured compounds: copper(II) compounds are blue, iron(II) compounds green, iron(III) compounds orange-brown, which is why their solutions are so vivid. And they show variable oxidation states, forming ions with more than one charge — iron as \(\text{Fe}^{2+}\) and \(\text{Fe}^{3+}\), copper as \(\text{Cu}^{+}\) and \(\text{Cu}^{2+}\). Set that against sodium, which is soft enough to cut with a knife, low in density, violently reactive, forms only \(\text{Na}^{+}\), and whose compounds are white — and the contrast is the whole comparison an exam question wants.
VocabularyKey terms the mark scheme pays for
TrapsMisconceptions that cost marks
ExamWhat examiners want
When you describe atomic structure, give the three particles with both their relative charge and their relative mass, and always tie the proton number to the identity of the element — that link is where the marks concentrate. For relative atomic mass calculations, write the weighted-average formula, substitute, and show the division by the total abundance; simply averaging the isotope masses is the error examiners look for, and it loses the method marks as well as the answer.
On "how the model of the atom changed" questions, credit is for linking each change to the evidence that forced it — Thomson's electrons, Rutherford's scattering, Bohr's fixed shells — not for reciting dates. For group-trend questions, do not stop at "more reactive" or "less reactive": explain it through the distance of the outer shell from the nucleus, shielding, and how easily an electron is lost or gained. Use the precise words — "atomic number" not "atomic mass" for ordering, "isotopes" not "ions" — and always give an ion's charge with its sign.