Atoms, for the exam.

Part 1The building blocks

All substances are made of atoms — the smallest part of an element that can exist. An atom is tiny: its radius is about 0.1 nm (1 × 10⁻¹⁰ m), and the nucleus is around 1/10 000 of that.

An element is a substance made of only one type of atom — there are about 100, each shown by a symbol on the periodic table. A compound is two or more different elements chemically bonded together, in fixed proportions; you can only split a compound up by a chemical reaction. A mixture is two or more substances not chemically joined, so they keep their own properties and can be separated by physical methods.

Part 2Separating mixtures

Because the parts of a mixture aren't bonded, you can separate them by physical methods that exploit a difference in their properties.

Part 1From solid spheres to a nucleus

Dalton (early 1800s) pictured atoms as tiny, solid, indivisible spheres — different for each element. Then J. J. Thomson discovered the electron (a tiny negative particle), which meant atoms aren't indivisible. His plum pudding model showed the atom as a ball of positive charge with electrons dotted through it.

Next, Rutherford's team fired positive alpha particles at thin gold foil. Most passed straight through, but a few bounced back. This alpha scattering experiment disproved the plum pudding model and led to the nuclear model: the mass and positive charge are concentrated in a tiny central nucleus, with most of the atom being empty space.

Part 2Bohr and Chadwick finish the picture

Niels Bohr refined the model: electrons orbit the nucleus at fixed distances in energy levels (shells). His calculations agreed with experiment. Later work showed the positive charge of the nucleus is made of smaller particles — protons. Finally, James Chadwick provided evidence for the neutron, a particle in the nucleus with mass but no charge.

Part 1Inside the atom

Atoms contain three sub-atomic particles. Protons and neutrons sit in the nucleus; electrons orbit it in shells. Atoms have no overall charge because the number of protons (positive) equals the number of electrons (negative).

Two numbers describe an atom. The atomic number (bottom) is the number of protons — it defines the element. The mass number (top) is the total number of protons + neutrons. So:

neutrons = mass number − atomic number

Part 2Isotopes & relative atomic mass

Isotopes are atoms of the same element (same number of protons) with different numbers of neutrons — so they have the same atomic number but different mass numbers. Because chemistry depends on electrons, isotopes of an element react the same way.

The relative atomic mass (A_r) of an element is the average mass of its atoms, taking into account the abundance (how common) each isotope is.

Part 1Filling the shells

Electrons occupy the lowest available energy level (shell) first, filling inner shells before outer ones. The shells hold a maximum number of electrons:

First shell: up to 2. Second shell: up to 8. Third shell: up to 8 (for the elements you need). So the pattern is 2, 8, 8.

We write the electronic structure as the number in each shell, e.g. sodium (11 electrons) is 2, 8, 1.

Part 2The outer shell decides behaviour

The number of electrons in the outer shell equals the group number in the periodic table (for the main groups). It controls how an element reacts. This is why elements in the same group behave alike — they have the same number of outer-shell electrons.

Part 1How it's organised

The modern periodic table arranges elements in order of atomic number (number of protons). Elements with similar properties line up in vertical groups; horizontal rows are periods. Elements in the same group have the same number of outer-shell electrons, which is why they react in similar ways.

Part 2Mendeleev, and metals vs non-metals

Mendeleev (1869) arranged the known elements mainly by atomic mass, but he did two clever things: he left gaps for undiscovered elements and even predicted their properties, and he swapped the order of a few elements so they fell into groups with similar properties. When the missing elements were later found and matched his predictions, his table was accepted. We now know the gaps and swaps make sense because the table really follows atomic number.

The table splits into metals (on the left and centre) and non-metals (top right). Metals form positive ions by losing electrons; non-metals tend not to. Most elements are metals.

Part 1Unreactive by design

The noble gases in Group 0 (helium, neon, argon and so on) are very unreactive. The reason is their electron arrangement: they have a full outer shell of electrons (helium has 2; the rest have 8). A full outer shell is stable, so they don't easily gain, lose or share electrons. This is why they're often called inert.

Part 2The boiling-point trend

As you go down Group 0, the boiling points increase. The atoms get bigger and heavier, so the forces between the atoms (intermolecular forces) get stronger and need more energy to overcome. That's why helium (top) is the lowest-boiling and the gases lower down boil at higher temperatures.

Part 1One outer electron

The alkali metals (lithium, sodium, potassium…) are in Group 1. They each have one electron in their outer shell, which they lose when they react to form a +1 ion. They are soft (you can cut them with a knife) and react readily — so they're stored under oil to keep air and water away.

With water, they react to make a metal hydroxide (an alkali — hence the name) and hydrogen gas. The metal fizzes and floats, moving across the surface. With oxygen they tarnish quickly, forming a metal oxide.

sodium + water → sodium hydroxide + hydrogen

Part 2Reactivity increases down the group

Going down Group 1, the metals get more reactive. Lithium fizzes gently; sodium reacts faster and melts into a ball; potassium reacts so vigorously the hydrogen ignites with a lilac flame. This is the opposite of the trend you might expect, and there's a clear reason behind it.

Part 1Seven outer electrons

The halogens (fluorine, chlorine, bromine, iodine) are in Group 7. They each have seven electrons in the outer shell, so they gain one electron when they react, forming a −1 ion. They exist as molecules of two atoms (e.g. Cl₂, Br₂).

Going down the group the halogens get less reactive, and their melting and boiling points increase — so at room temperature chlorine is a gas, bromine a liquid and iodine a solid.

Part 2Displacement reactions

A more reactive halogen displaces a less reactive halogen from a solution of its salt. For example, chlorine (more reactive) displaces bromine from potassium bromide solution, turning it orange as bromine is released:

chlorine + potassium bromide → potassium chloride + bromine

Iodine could not displace chlorine or bromine, because iodine is the least reactive of the three.