AS Chemistry - Periodicity

The periodic table, trends in properties of elements in period 3, ionisation energies

HideShow resource information

The periodic table

Metals and non-metals are separated by a stepped black line, metals on the left and non-metals on the right. The elements that touch the right of the line are calles 'metalloids' because although they are non-metals, they show metallic properties.

The periodic table is split into blocks called s-, p-, d- and f-. They relate to the orbitals in which elements' electrons are found. For example, all the elements that have their highest energy electrons in the s-orbitals are found in the s-block.

Elements in the same group have similar chemical properties as they have the same number of electrons in the outer shell. In the s-block, elements tend to get more reactive down the group, whereas non-metals (on the right) get more reactive up the group. Transition metals are mostly unreactive, and so are very useful.

Periods are the horizontal rows of the periodic table. They have trends in physical and chemical properties.

1 of 3

Trends in Period 3 elements (1)

Na, Mg and Al are metals with giant structures. They form ionic compounds.

Si is a metalliod. It can make four covalent bonds.

P, S and Cl are non-metals and either form ionic or covalent compounds.

Ar is a noble gas so is unreactive.

Melting points: Giant structures have very high melting points, while molecular or atomic structures have lower ones.

Increases from Na to Al due to metallic bonding - charge of ion increases, so number of delocalised electrons increases, creating a stronger bond.

Silicon has a giant structure and therefore a high melting point.

M.p.s of non-metals depends on van der Waals forces, so molecules with more electrons and which can fit together more closely will have higher melting points.

2 of 3

Trends in Period 3 elements (2)

Atomic radii: decreases across the period, due to increased nuclear charge, which pulls the electrons closer to the nucleus.

Ionisation energy: increases across the period, due to increased nuclear charge, which makes it harder to remove an electron.

There is a drop in IEs between groups 2 and 3, because the outer electron is now in a higher energy level than before (p- instead of s-), so it is easier to remove.

There is a drop in IEs between groups 5 and 6, because there are now two electrons in an orbital rather than one, so the repulsion between these electrons makes it easier to remove an electron.

3 of 3


No comments have yet been made

Similar Chemistry resources:

See all Chemistry resources »See all The Periodic Table resources »