Most of how an element behaves comes from the arrangement of its electrons, and the periodic table is organised so that elements with similar arrangements sit together.
An element contains one kind of atom; a compound contains two or more elements chemically bonded in fixed proportions and can only be broken down into simpler substances by chemical processes such as electrolysis.
A mixture contains substances that are not chemically bonded, so each retains its chemical identity and physical methods will separate them: filtration, crystallisation, simple and fractional distillation, and chromatography.
Protons have a relative charge of +1 and a relative mass of 1; neutrons are neutral with a mass of 1; electrons are −1 with a mass close to zero.
An atom has equal numbers of protons and electrons, so it is neutral overall. The number of protons, the atomic number, is what makes an atom that element.
An atom has a radius of about \(1\times10^{-10}\) m. The nucleus is tiny by comparison, with a radius less than \(\frac{1}{10000}\) of the atom's, yet it holds almost all the mass.
So an atom is overwhelmingly empty space, with a very dense centre.
Isotopes are atoms of the same element with different numbers of neutrons: the same atomic number but a different mass number.
Relative atomic mass is the average of the relative isotopic masses of an element, weighted by how common each isotope is. That is why chlorine's is 35.5 rather than a whole number: it is roughly three parts chlorine-35 to one part chlorine-37.
To calculate it, multiply each isotope’s relative isotopic mass, usually approximated by its mass number, by its percentage abundance, add the results, and divide by 100.
Electrons occupy shells around the nucleus, filling the lowest first. The first shell holds up to 2 electrons, the second up to 8 and the third up to 8 at this level.
Sodium, with 11 electrons, is written 2,8,1. The number of electrons in the outer shell strongly influences an element's chemical properties. For Groups 1 to 7 it equals the group number, and Group 0 has a full outer shell.
Atoms were first modelled as tiny indivisible spheres. Discovering the electron gave the plum pudding model, a ball of positive charge with electrons dotted through it.
The alpha scattering experiment overturned it: most alpha particles went straight through gold foil, and a few bounced back, showing a small dense positive nucleus in mostly empty space. Bohr then proposed that electrons occupy particular energy levels, or shells; the proton was identified later, and the neutron about twenty years after that.
This is an example of a scientific model changing because new experimental evidence demanded it.
Early tables ordered elements by atomic weight, which put some elements in groups that did not match their properties. Mendeleev left gaps for undiscovered elements and moved a few out of strict weight order, and when the missing elements were found with the predicted properties his table was accepted.
The modern table is ordered by atomic number, which resolved the anomalies. Understanding isotopes later explained why ordering by atomic mass had sometimes been misleading. Periods are rows. Groups are columns, and for the main groups the elements in a group have the same number of outer-shell electrons and so react similarly.
Metals are on the left and centre. They lose electrons to form positive ions, conduct heat and electricity, and are usually malleable with high melting points.
Non-metals are on the right. They gain or share electrons, are usually poor conductors, and are typically brittle when solid.
The noble gases are unreactive because they have a full outer shell of eight electrons, or two for helium, so they have very little tendency to gain, lose or share electrons.
Boiling point increases down the group because larger atoms experience stronger intermolecular forces.
The alkali metals have one outer electron, which they lose readily to form 1+ ions. They are soft, have low densities, and react vigorously with water to give a metal hydroxide and hydrogen.
Reactivity increases down the group: the outer electron is further from the nucleus and more shielded, so it is lost more easily. Lithium fizzes, sodium melts into a ball, potassium ignites.
Chlorine, bromine and iodine exist as diatomic molecules, and the halogens have seven outer electrons, so they gain one to form 1− ions. Melting point and boiling point increase down the group: chlorine is a green gas, bromine a brown liquid, iodine a grey solid.
Reactivity decreases down the group, the opposite of Group 1, because the incoming electron is further from the nucleus and less strongly attracted. A more reactive halogen therefore displaces a less reactive one from its salt.
Group 1 elements lose one electron, and their reactivity increases down the group. Group 7 elements gain one electron, and their reactivity decreases down the group.
The transition metals in the central block are much less reactive than Group 1 and generally have higher melting points and densities. Mercury is an exception, being liquid at room temperature.
They also have properties Group 1 metals lack: they form ions with different charges, their compounds are often coloured, and many are useful as catalysts, such as iron in the Haber process and nickel in hydrogenation.
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