Atomic structure and the periodic table.

HookThe swap Mendeleev couldn't explain

ModelAtoms, elements, compounds — and separating mixtures

Start with the building blocks. An atom is the smallest particle of an element that can exist. An element contains only one type of atom — there are about 100, each with a one- or two-letter symbol on the periodic table. A compound contains two or more elements chemically combined in fixed proportions — water is always two hydrogen to one oxygen — and a compound can only be separated back into its elements by chemical reactions or electrolysis, because forming it makes and breaks bonds. A mixture is different: two or more substances not chemically joined, so each keeps its own properties and can be pulled apart by physical means that need no reaction.

Those physical methods are worth knowing by trigger. Filtration separates an insoluble solid from a liquid, such as sand from water. Crystallisation recovers a soluble solid from its solution by evaporating the solvent until crystals form, such as copper sulfate from solution. Simple distillation separates a solvent from a solution by boiling it off and condensing it, giving pure water from salt water. Fractional distillation separates two or more liquids with different boiling points, such as ethanol from water or the fractions of crude oil. Chromatography separates dissolved substances by how differently they travel through a medium. Each is a physical change — no new substance is made — which is itself the giveaway that you are dealing with a mixture, not a compound.

CaseThe model of the atom — rebuilt as the evidence arrived

The atom is the best example in the course of a scientific model being rebuilt as evidence accumulates. John Dalton, around 1803, pictured atoms as tiny solid spheres that could not be divided. Then in 1897 J. J. Thomson discovered the electron — a particle far smaller than an atom, with a negative charge — which meant atoms were divisible after all. His 'plum pudding' model imagined a ball of positive charge with negative electrons dotted through it.

The decisive evidence came around 1909, when Ernest Rutherford's team fired positively charged alpha particles at a very thin gold foil. Most passed straight through, but a few were deflected and a very few bounced back sharply. A spread-out pudding of charge could not do that; the result only made sense if the positive charge and nearly all the mass were concentrated in a tiny central nucleus, with the atom mostly empty space. This nuclear model was refined by Niels Bohr, who showed the electrons orbit at fixed distances in energy levels rather than anywhere — a change that explained why atoms are stable and matched the light atoms emit. Later work identified positive protons in the nucleus, and James Chadwick's experiments around 1932 provided evidence for the neutron. This whole story is common content shared with the physics course, and examiners like asking you to match a piece of evidence to the model it forced.

ModelSubatomic particles, mass number and isotopes

An atom has a nucleus of protons and neutrons, surrounded by electrons. Their relative charges are the numbers to memorise: a proton is +1, an electron is −1, and a neutron is 0. Because atoms have equal numbers of protons and electrons, the charges cancel and an atom overall is neutral. The relative masses matter too: a proton and a neutron each count as 1, while an electron is only about 1/1835 of that — so small that electrons barely affect an atom's mass. Almost all the mass sits in the nucleus.

Scale is startling. A typical atom has a radius of about 1 × 10⁻¹⁰ m (0.1 nanometres), while its nucleus is around 1 × 10⁻¹⁴ m — roughly ten thousand times smaller, so an atom is overwhelmingly empty space. Two counts define an atom: the atomic number is the number of protons and it defines which element you have, while the mass number is the total of protons plus neutrons. Isotopes are atoms of the same element — the same number of protons — with different numbers of neutrons, so they share an atomic number but differ in mass number. Carbon-12 and carbon-14 are the same element and behave identically in reactions; carbon-14 just carries two extra neutrons.

DataRelative atomic mass — a weighted average of the isotopes

Because most elements exist as a mixture of isotopes, the mass printed on the periodic table is not the mass of a single atom but a weighted average — the relative atomic mass (Ar). Weighted means each isotope counts in proportion to how common it is. The method is always the same: multiply each isotope's mass number by its percentage abundance, add these together, and divide by 100 (the total abundance). If you are instead given fractional abundances that add up to 1, you divide by 1 — that is, you do not divide at all.

This is why relative atomic masses are so often not whole numbers, even though every individual atom has a whole-number mass. Chlorine's familiar 35.5 is not the mass of any real chlorine atom; it is the average of two isotopes, weighted by how much of each exists in nature.

ModelElectronic structure — why the table is arranged as it is

Electrons are not scattered randomly; they occupy energy levels, or shells, around the nucleus, filling the lowest (innermost) level first before starting the next. For the first twenty elements the capacities are two electrons in the first shell, then eight, then eight. Sodium, with 11 electrons, is written 2,8,1: two in the first shell, eight in the second and one left over in the third. This arrangement is the element's electronic structure, and it can be drawn as rings of dots or written as numbers separated by commas.

The payoff is that electronic structure maps straight onto the periodic table, which is why the table works at all. The number of electrons in the outer shell equals the group number for the main groups — every Group 1 element has one outer electron, every Group 7 element has seven. The number of occupied shells equals the period, or row, number. Because chemical reactions involve the outer electrons, elements with the same number of outer electrons react in similar ways — that shared outer count is the reason a whole group behaves as a family. Get an element's electronic structure right and you can predict its group, its likely ion and much of its chemistry before you know anything else about it.

ModelThe periodic table, its development, and metals versus non-metals

The modern periodic table arranges all the elements in order of increasing atomic number. The vertical columns are groups, whose elements share the same number of outer electrons and therefore similar chemical properties; the horizontal rows are periods, and moving across one adds electrons to the same outer shell. Getting to this arrangement took a century of wrong turns. Early chemists ordered elements by atomic mass and grouped them loosely — John Newlands' 'law of octaves' lined them up in eights but broke down because he assumed every element had already been found and left no room for gaps. Mendeleev's leap, described above, was to leave gaps for undiscovered elements and to order by properties where mass misled him. He then predicted the properties of the missing elements — his 'eka-silicon' turned out to be germanium, discovered later with almost exactly the density and behaviour he had forecast. Only when atomic number replaced atomic mass as the ordering rule did the last anomalies disappear.

The table also splits into metals and non-metals. Metals sit on the left and towards the bottom; they have few outer electrons and react by losing them to form positive ions. Non-metals sit on the upper right; they tend to gain or share electrons, forming negative ions or covalent bonds. The dividing staircase runs down the right-hand side, and elements near it can show a mix of behaviours.

MechanismGroup trends — noble gases, alkali metals and halogens

Three groups carry the trends the exam tests. Group 0, the noble gases, have full outer shells — helium with two, the rest with eight — which makes them stable and very unreactive, so they exist as single atoms rather than bonding. Their boiling points increase down the group as the atoms get bigger and the forces between them strengthen. Group 1, the alkali metals, each have one outer electron, which they lose easily to form a 1+ ion. They react with water to produce a metal hydroxide and hydrogen gas, fizzing and skating across the surface, and — the key trend — they get more reactive down the group. That is because the outer electron sits further from the nucleus, with more inner shells shielding it, so it is held less tightly and lost more readily.

Group 7, the halogens, are the mirror image. Each has seven outer electrons and needs to gain just one, so they exist as diatomic molecules and form 1− ions. Their reactivity decreases down the group: as the atoms get larger, the outer shell is further from the nucleus, so it is harder to attract and capture the extra electron. Their melting and boiling points, meanwhile, increase down the group. A more reactive halogen will displace a less reactive one from a solution of its salt — add chlorine to potassium bromide and the chlorine, being more reactive, takes the bromine's place, turning the solution orange as bromine is released.

VocabularyKey terms the mark scheme pays for

TrapsMisconceptions that cost marks

ExamWhat examiners want