What you'll learn
Atomic structure and the periodic table is the opening unit of AQA GCSE Combined Science: Trilogy chemistry, and every later unit depends on it. It explains what atoms are made of, how our model of the atom was built up and revised over a century of experiments, how electrons are arranged, and why that arrangement explains the shape of the periodic table and the behaviour of whole groups of elements. By the end you should be able to describe the structure of an atom in terms of protons, neutrons and electrons with their relative masses and charges, work out the number of each particle from an atomic number and mass number, explain what isotopes are and calculate relative atomic mass from their abundances, write electronic structures for the first twenty elements, explain why the periodic table is arranged as it is, and describe the trends in Group 1, Group 7 and Group 0. This unit is assessed on Chemistry Paper 1.
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
Compound — a substance containing two or more elements chemically combined in fixed proportions
Mixture — two or more elements or compounds not chemically combined, so each keeps its own properties
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 mean mass of the atoms of an element, taking into account the abundance of each isotope
Electronic structure — the arrangement of electrons in shells around the nucleus, written for example as 2,8,1
Group — a vertical column of the periodic table; elements in the same group have the same number of outer electrons
Period — a horizontal row of the periodic table
Core concepts
Atoms, elements, compounds and mixtures
An atom is the smallest part of an element that can exist. Elements contain only one type of atom, and there are about 100 of them, each with its own symbol shown in the periodic table.
Compounds are formed from elements by chemical reactions, which always involve the formation of one or more new substances and often an observable energy change. Compounds contain two or more elements chemically combined in fixed proportions, and they can only be separated into those elements by another chemical reaction.
Mixtures are different. The components are not chemically combined, so each substance keeps its own physical and chemical properties, and the mixture can be separated by physical processes: filtration, crystallisation, simple distillation, fractional distillation and chromatography. None of these involves a chemical reaction, which is precisely why they work on mixtures but not on compounds.
The development of the model of the atom
The model of the atom changed as new experimental evidence emerged, and this history is explicitly examinable.
Before the electron was discovered, John Dalton described atoms as solid spheres, with different spheres making up the different elements. When J. J. Thomson discovered the electron, the solid sphere model had to change, and the plum pudding model was suggested: a ball of positive charge with negative electrons embedded in it.
The alpha particle scattering experiment overturned this. When alpha particles were fired at thin gold foil, most passed straight through, but a small number were deflected and a very few bounced back. Since most passed through, the atom had to be mostly empty space; since a few were strongly deflected, the positive charge and nearly all the mass had to be concentrated in a tiny nucleus. This led to the nuclear model.
Niels Bohr then adapted the model by proposing that electrons orbit the nucleus at specific distances in shells, which agreed with experimental observations. Later work showed that the positive charge of the nucleus was divided into particles with equal charge, named protons, and about twenty years later James Chadwick provided evidence for neutrons.
Sub-atomic particles
An atom has a radius of about 0.1 nanometres, and the nucleus is about one ten-thousandth of that, roughly 1 times 10 to the power minus 14 metres. Almost all the mass is in the nucleus, and the rest of the atom is empty space containing electrons.
Protons have a relative charge of plus 1 and a relative mass of 1. Neutrons have no charge and a relative mass of 1. Electrons have a relative charge of minus 1 and a relative mass of very nearly zero. Atoms have no overall charge because the number of protons equals the number of electrons.
The atomic number is the number of protons and is what identifies the element. The mass number is protons plus neutrons, so the number of neutrons is found by subtracting the atomic number from the mass number.
Isotopes and relative atomic mass
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. Because chemical behaviour depends on electrons, and the number of electrons is unchanged, isotopes of an element react identically. Their physical properties, which depend on mass, differ slightly.
Because an element exists as a mixture of isotopes, its relative atomic mass is a weighted mean. Multiply each isotope's mass number by its percentage abundance, add the results, and divide by 100.
Electronic structure
Electrons occupy the lowest available energy levels, or shells. The first shell holds up to two electrons, and the second and third shells each hold up to eight for the elements you need. Electronic structure is written as a series of numbers separated by commas, so sodium with 11 electrons is 2,8,1.
This notation explains the periodic table immediately. The number of shells in use gives the period; the number of electrons in the outer shell gives the group.
The periodic table
Elements are arranged in order of atomic number, in rows, so that elements with similar properties line up in columns called groups. Elements in the same group have the same number of electrons in their outer shell, which is why they have similar chemical properties.
Early attempts to organise the elements ordered them by atomic weight, which put some elements in the wrong group. Dmitri Mendeleev solved this by leaving gaps for elements he believed had not yet been discovered and, in places, changing the order. The gaps were later filled by elements with properties matching his predictions, and once isotopes were understood it became clear why ordering by atomic weight sometimes failed.
Metals are found to the left and towards the bottom of the table; non-metals to the right and towards the top. Metals react to form positive ions, because they lose their outer electrons; non-metals generally do not form positive ions.
Group 0, the noble gases
The noble gases are unreactive and exist as single atoms because they have a full outer shell of electrons — eight, except for helium, which has two. They therefore have no tendency to gain, lose or share electrons.
Boiling points increase going down the group as the relative atomic masses increase.
Group 1, the alkali metals
Group 1 elements all have one electron in their outer shell, which makes them very reactive and gives them similar properties. They react with non-metals to form ionic compounds in which the metal ion carries a charge of plus 1, and the compounds are white solids that dissolve to give colourless solutions.
Reacting them with water produces hydrogen and a metal hydroxide, which is alkaline. Reactivity increases going down the group. The reason is that the outer electron is in a shell further from the nucleus and is shielded by more inner shells, so it is more easily lost.
Group 7, the halogens
The halogens are non-metals that exist as molecules of two atoms. They have seven electrons in their outer shell. They form ionic compounds with metals in which the halide ion carries a charge of minus 1, and they form covalent compounds with other non-metals.
Going down the group, relative molecular mass, melting point and boiling point all increase, but reactivity decreases. The explanation mirrors Group 1: the outer shell is further from the nucleus and more shielded, so an electron is gained less readily.
A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt. Chlorine will displace bromine from potassium bromide solution, but bromine will not displace chlorine from potassium chloride.
Worked examples
Example 1: Counting sub-atomic particles (3 marks)
An atom of potassium has atomic number 19 and mass number 39. State the number of protons, neutrons and electrons.
The atomic number is the number of protons, so there are 19 protons. In a neutral atom the number of electrons equals the number of protons, so there are 19 electrons. The number of neutrons is the mass number minus the atomic number, which is 39 minus 19, giving 20 neutrons.
Example 2: Calculating relative atomic mass (3 marks)
Chlorine exists as two isotopes: 75 per cent has a mass number of 35 and 25 per cent has a mass number of 37. Calculate the relative atomic mass.
Multiply each mass number by its abundance and add: 35 multiplied by 75 gives 2,625, and 37 multiplied by 25 gives 925. The total is 3,550. Divide by 100 to give a relative atomic mass of 35.5. Notice the answer is closer to 35 than to 37, which makes sense because the lighter isotope is more abundant — a useful check.
Example 3: Explaining reactivity in Group 1 (4 marks)
Explain why potassium is more reactive than sodium.
Both metals have one electron in their outer shell, and reacting involves losing that electron. In potassium the outer electron is in a shell further from the nucleus than in sodium, and there are more inner shells shielding it from the positive nuclear charge. The force of attraction between the nucleus and the outer electron is therefore weaker, so the electron is lost more easily and potassium reacts more vigorously.
Common mistakes and how to avoid them
The most frequent error is confusing atomic number with mass number when working out neutrons. The atomic number is always the smaller of the two on the periodic table, and neutrons are found by subtraction.
Students often state that isotopes have different chemical properties. They do not; chemical properties depend on the outer electrons, which are the same. Only physical properties such as density differ.
In Group 1 and Group 7 explanations, many answers say only that the atom is bigger. Size alone is not the mark; the marks are for distance from the nucleus, shielding by inner shells, and the resulting weaker attraction.
Another common slip is describing the noble gases as unreactive because they are stable. That is circular. The reason is a full outer shell, so there is no tendency to gain, lose or share electrons.
Finally, students frequently write that a mixture can be separated by a chemical reaction. Mixtures are separated physically; that is the definition.
Exam technique for "Chemistry: Atomic Structure and the Periodic Table"
Questions on the development of the atomic model want the evidence linked to the change. Saying that Rutherford discovered the nucleus scores little; saying that most alpha particles passed through, showing the atom is mostly empty space, and a few were deflected back, showing a small dense positive nucleus, scores well.
For electronic structure, always check the total number of electrons equals the atomic number before writing the answer. It takes two seconds and catches most errors.
Trend questions in Groups 1, 7 and 0 follow a fixed structure: state the trend, then explain it in terms of the outer shell's distance from the nucleus and shielding. Learn that explanation once and it serves both groups, with the direction reversed.
When asked to compare a metal with a non-metal, refer to ion formation — metals form positive ions by losing electrons — rather than to vague properties such as shininess.
Quick revision summary
Atoms have a radius of about 0.1 nanometres with a tiny dense nucleus of protons and neutrons surrounded by electrons in shells. Protons are plus 1 with mass 1, neutrons are neutral with mass 1, electrons are minus 1 with negligible mass. The atomic number is the proton count and identifies the element; the mass number is protons plus neutrons. Isotopes differ only in neutron number and react identically, and relative atomic mass is a weighted mean of their abundances. The atomic model developed from Dalton's spheres through the plum pudding model to the nuclear model after alpha scattering, then Bohr's shells, protons and finally neutrons. Electrons fill shells holding 2, 8 and 8, and the electronic structure gives period and group directly. Group 0 is unreactive because the outer shell is full; Group 1 reactivity increases down the group and Group 7 reactivity decreases, both explained by distance from the nucleus and shielding. A more reactive halogen displaces a less reactive one from solution.