What you'll learn
Atomic structure and the periodic table is the opening topic of AQA GCSE Chemistry, and every later topic rests on it. Bonding, reactivity, electrolysis and quantitative chemistry all become straightforward once you can say confidently what an atom contains and how its electrons are arranged. This topic also carries the course's clearest example of how science works: the model of the atom was revised four times in under a century, each time because an experiment produced a result the existing model could not explain. By the end of this guide you should be able to describe the atom's structure and scale, work out numbers of protons, neutrons and electrons, calculate relative atomic mass from isotopic abundances, write electronic structures and relate them to position in the periodic table, explain the historical development of the atomic model from the evidence, and account for the trends in Groups 0, 1 and 7 and the properties of the transition metals.
Key terms and definitions
Atom — the smallest part of an element that can exist, with a radius of about 0.1 nanometres
Element — a substance made of only one type of atom
Atomic number — the number of protons in an atom, which defines the element
Mass number — the total number of protons and neutrons in an atom
Isotope — an atom of the same element with the same number of protons but a different number of neutrons
Relative atomic mass — the weighted mean mass of an element's atoms, accounting for isotopic abundance
Electronic structure — the arrangement of electrons in shells, written for example as 2,8,1
Group — a vertical column of the periodic table; members share the number of outer electrons
Period — a horizontal row; members share the number of occupied electron shells
Transition metal — an element from the central block of the periodic table, typically forming coloured compounds and acting as a catalyst
Halogen — a Group 7 element, existing as diatomic molecules
Noble gas — a Group 0 element, unreactive because its outer shell is full
Core concepts
The structure and scale of the atom
An atom has a radius of about 0.1 nanometres, which is 1 times 10 to the power minus 10 metres. The nucleus lies at the centre and has a radius less than one ten-thousandth of that, around 1 times 10 to the power minus 14 metres. Almost the entire mass of the atom is concentrated in this tiny nucleus, and the rest of the atom is empty space occupied by electrons.
Protons carry a relative charge of plus 1 and have a relative mass of 1. Neutrons are uncharged with a relative mass of 1. Electrons carry a relative charge of minus 1 and have a relative mass of very nearly zero.
Because the number of protons equals the number of electrons, atoms have no overall charge. If an atom loses one or more outer electrons it becomes a positive ion; if it gains electrons it becomes a negative ion.
The atomic number is the number of protons and identifies the element absolutely — change it and you have a different element. The mass number is protons plus neutrons, so the neutron count is found by subtracting the atomic number from the mass number.
Isotopes and relative atomic mass
Isotopes are atoms of the same element with different numbers of neutrons. Because chemical behaviour depends on the arrangement of electrons, and isotopes have identical electron arrangements, isotopes of an element react identically. Only physical properties that depend on mass, such as density and rate of diffusion, differ.
Elements exist as mixtures of isotopes, so the relative atomic mass quoted in the periodic table is a weighted mean. Multiply each isotope's mass number by its percentage abundance, sum the products, and divide by 100.
A useful check on any answer: the result must lie between the mass numbers of the lightest and heaviest isotopes, and closer to whichever is more abundant.
Electronic structure
Electrons occupy shells at increasing distances from the nucleus, filling the lowest available energy level first. The first shell holds up to two electrons, and the second and third shells hold up to eight for the elements you need at GCSE.
Electronic structure is written as a series of numbers separated by commas. Sodium, with eleven electrons, is 2,8,1. Chlorine, with seventeen, is 2,8,7.
This notation maps directly onto the periodic table. The number of occupied shells gives the period; the number of electrons in the outer shell gives the group. Sodium's structure of 2,8,1 places it in period 3, group 1 — which is exactly where it sits.
This relationship is the single most useful idea in the topic, because it means the periodic table is not an arbitrary arrangement but a direct display of electronic structure.
The development of the atomic model
The historical sequence is explicitly examinable, and what matters is linking each change to the evidence that forced it.
Before the electron was discovered, John Dalton described atoms as tiny solid spheres that could not be divided, with different spheres making up different elements.
J. J. Thomson's discovery of the electron showed that atoms contain smaller charged particles, so the solid sphere model could not stand. The plum pudding model replaced it: a ball of positive charge with negative electrons embedded within it.
The alpha particle scattering experiment then overturned that. Alpha particles were fired at a very thin sheet of gold foil. Most passed straight through, showing that the atom is mostly empty space. A small number were deflected through large angles and a very few rebounded, showing that the positive charge and nearly all the mass are concentrated in a very small central volume. This produced the nuclear model.
Niels Bohr adapted the nuclear model by proposing that electrons orbit at specific fixed distances in shells, and his theoretical calculations agreed with experimental observations.
Later experiments showed that the nuclear positive charge consists of particles of equal charge, named protons. About twenty years after that, James Chadwick provided the experimental evidence for neutrons, completing the model used today.
The general principle examiners want stated is that new experimental evidence may cause a scientific model to be changed or replaced, and that the new model must be checked by other scientists before it is accepted.
The periodic table and its history
Elements are arranged in order of atomic number in rows, so that elements with similar properties fall into the same vertical column.
Early attempts arranged the elements by atomic weight, which placed some elements in groups whose properties they did not share. Dmitri Mendeleev overcame this by leaving gaps for elements he predicted had not yet been discovered, and in places by changing the strict order of atomic weight. The later discovery of elements matching his predictions confirmed the approach, and once isotopes were understood it became clear why ordering by atomic weight sometimes failed — isotopes mean atomic weight does not always increase with atomic number.
Metals lie towards the left and bottom of the table and react by losing electrons to form positive ions. Non-metals lie towards the right and top and do not form positive ions.
Group 0, the noble gases
Group 0 elements are unreactive and exist as single atoms rather than molecules, because their outer shell is full — eight electrons, except helium which has two. There is no tendency to gain, lose or share electrons, so they form almost no compounds.
Boiling points increase down the group as relative atomic mass increases, because the intermolecular forces between larger atoms are stronger.
Group 1, the alkali metals
Group 1 elements have one electron in the outer shell, which they lose readily to form ions with a charge of plus 1. Their compounds are white solids dissolving to give colourless solutions.
Reacting with water produces hydrogen and a metal hydroxide, which is alkaline. Lithium fizzes steadily, sodium melts into a ball and moves rapidly, and potassium ignites with a lilac flame.
Reactivity increases down the group. The explanation has three parts and all three are usually needed: the outer electron is in a shell further from the nucleus, it is shielded by more inner shells, so the attraction between nucleus and outer electron is weaker and the electron is lost more easily.
Group 7, the halogens
The halogens are non-metals existing as molecules of two atoms, with seven outer electrons. They form ionic compounds with metals in which the halide ion has a charge of minus 1, and covalent compounds with other non-metals.
Going down the group, relative molecular mass, melting point and boiling point all increase, but reactivity decreases. The reasoning mirrors Group 1 in reverse: the outer shell is further from the nucleus and more shielded, so an electron is attracted and gained less readily.
A more reactive halogen displaces a less reactive one from an aqueous solution of its salt. Chlorine displaces both bromine and iodine; bromine displaces iodine only.
Transition metals
The transition metals occupy the central block and are examined on the separate Chemistry course as a contrast with Group 1.
Compared with the alkali metals they have higher melting points and higher densities, they are stronger and harder, and they are much less reactive — they do not react vigorously with oxygen or water.
They also have three characteristic properties: many form ions with different charges, their compounds are often coloured, and they and their compounds are frequently useful as catalysts. Iron in the Haber process and nickel in hydrogenation are the standard examples.
Worked examples
Example 1: Determining particle numbers (3 marks)
An atom of an element has a mass number of 56 and an atomic number of 26. State the numbers of protons, neutrons and electrons.
The atomic number gives the number of protons directly, so there are 26 protons. In a neutral atom the number of electrons equals the number of protons, so there are 26 electrons. The number of neutrons is the mass number minus the atomic number, which is 56 minus 26, giving 30 neutrons.
Example 2: Calculating relative atomic mass (3 marks)
Copper consists of 69 per cent of an isotope with mass number 63 and 31 per cent of an isotope with mass number 65. Calculate the relative atomic mass to one decimal place.
Multiply each mass number by its abundance and add the products: 63 multiplied by 69 gives 4,347, and 65 multiplied by 31 gives 2,015. The total is 6,362. Dividing by 100 gives a relative atomic mass of 63.6. As a check, the answer lies between 63 and 65 and is closer to 63, which is correct because the lighter isotope is the more abundant.
Example 3: Explaining a Group 7 trend (4 marks)
Explain why chlorine is more reactive than bromine.
When a halogen reacts it gains one electron into its outer shell. In bromine, the outer shell is further from the nucleus than in chlorine, and there are more inner shells shielding it from the positive nuclear charge. The force of attraction between the nucleus and an incoming electron is therefore weaker in bromine, so bromine gains an electron less readily than chlorine. Chlorine is consequently the more reactive, which is why chlorine displaces bromine from a solution of a bromide but bromine will not displace chlorine from a chloride.
Common mistakes and how to avoid them
The most frequent error is subtracting in the wrong direction when finding neutrons. The atomic number is always the smaller value, so neutrons equal mass number minus atomic number.
Students regularly claim that isotopes have different chemical properties. They do not, because they have identical electron arrangements; only mass-dependent physical properties differ.
In group trend explanations, many answers state only that the atom is larger. Size alone is not the mark: distance from the nucleus, shielding by inner shells, and the resulting change in attraction are three separate points.
Another common slip is explaining the noble gases' unreactivity by saying they are stable, which is circular. The reason is a full outer shell, so there is no tendency to gain, lose or share electrons.
Finally, candidates often describe Mendeleev as arranging elements by atomic number. He used atomic weight; atomic number was not known at the time, which is precisely why he had to change the order in places.
Exam technique for "Atomic Structure and the Periodic Table"
For questions on the atomic model, always pair the observation with the deduction. Stating that most alpha particles passed through earns the mark only when you add that this shows the atom is mostly empty space.
When writing electronic structures, check that the digits sum to the atomic number before moving on. This one check catches most errors.
For group trend questions, write a three-part explanation every time: distance from the nucleus, shielding, and the resulting strength of attraction. The same skeleton serves Group 1 and Group 7 with the direction reversed.
In relative atomic mass calculations, use the check that the answer must lie between the two mass numbers and nearer the more abundant isotope. It catches arithmetic slips instantly.
Quick revision summary
An atom has a radius of about 0.1 nanometres with a nucleus less than one ten-thousandth of that holding nearly all the mass. Protons are plus 1 with mass 1, neutrons neutral with mass 1, electrons minus 1 with negligible mass. Atomic number is the proton count and identifies the element; mass number is protons plus neutrons. Isotopes differ only in neutrons and react identically, and relative atomic mass is the abundance-weighted mean. Electrons fill shells holding 2, 8 and 8, and the electronic structure gives period from the number of shells and group from the outer electrons. The model developed from Dalton's spheres through the plum pudding model, overturned by alpha scattering, to the nuclear model, then Bohr's shells, protons and finally neutrons. Mendeleev ordered by atomic weight, leaving gaps and changing the order where needed. Group 0 is unreactive through full outer shells; Group 1 reactivity increases down the group and Group 7 decreases, both explained by distance, shielding and attraction. Transition metals are harder, denser, higher melting and less reactive than Group 1, form ions of varying charge, give coloured compounds and act as catalysts.