3.2.1 Periodicity

The periodic table The periodic table is a list of all the elements in order of increasing atomic number. You can predict the properties of an element from its position in the table. You can use it to explain the similarities of certain elements and the trends in their properties, in terms of electronic arrangements.

A group is a vertical column of elements. Elements in the same group have similar properties. Elements in the same group have the same number of electrons in the outer main energy level. Horizontal rows of elements in the periodic table are called periods. There are trends in physical and chemical properties as you go across a period.

Elements are classified as s, p or d block, according to which orbitals the highest energy electrons are in. For example: Na: s block 1s 2 2s 2 2p 6 3s 1 C: p block 1s 2 2s 2 2p 2 Fe: d block 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6

Trends in the properties of elements of Period 3 These are explained by the electron arrangements of the elements: The elements in Group 1, 2 and 3 (sodium, magnesium and aluminium) are metals with giant structures. They lose their outer electrons to form ionic compounds.

Silicon in Group 4 has 4 electrons in its outer shell with which it forms four covalent bonds. The element has some metallic properties and is classed as a semi-metal. The elements in Groups 5, 6 and 7 (phosphorus, sulphur and chlorine) are non- metals. They either accept electrons to form ionic compounds or share their outer electrons to form covalent compounds. Argon in Group 0 is a noble gas, it has a full outer shell and is unreactive.

The variation in melting is linked to the structure of the elements. Giant structures tend to have high melting points. Molecular or atomic structures tend to have low melting and boiling points.

Sodium, magnesium and aluminium are metals with giant metallic structure. The melting points increase from sodium to aluminium because the strength of the metallic bonding increases. As you go from left to right the charge on the metal ion increases so more electrons join the “sea” of delocalised electrons that holds the giant metallic lattice together.

Silicon has a high melting point because it has giant covalent (macromolecular) structure: strong covalent bonds link all the atoms in three dimensions and a large amount of energy is needed to break these bonds.

Phosphorus (P 4 ), sulphur (S 8 ) and chlorine (Cl 2 ) all have simple molecular structures. Their melting points are determined by the strength of the Van der Waals’ forces between molecules which, in turn, is determined by the number of electrons and how closely the molecules can pack together. Therefore the melting points of these elements are in the order S 8, P 4, Cl 2 (largest and most polarisable to smallest and least polarisable). These elements have low melting points because the van der Waals’ forces are weak and easily broken.

Atomic radii decrease across a period because the nuclear charge (proton number) increases but there is no change in the number of electron shells so there is no change in shielding. Therefore the outer electrons are pulled closer to the nucleus and the size of the atom decreases.

The first ionisation energy is the energy required to convert one mole of isolated gaseous atoms into one mole of singly positively charged gaseous ions.

First ionisation energies generally increase across a period because the nuclear charge (proton number) increases but there is no change in the number of electron shells so there is no change in shielding. Therefore the outer electrons are attracted increasingly strongly to the nucleus and it is increasingly difficult to remove an electron.

There is a drop in ionisation energy from magnesium to aluminium. This is because aluminium (1s 2 2s 2 2p 6 3s 2 3p 1 ) loses a 3p electron whereas magnesium (1s 2 2s 2 2p 6 3s 2 ) loses a 3s electron. The 3p electron is higher in energy than the 3s electron and so it takes less energy to remove it.

There is a drop in ionisation energy from phosphorus to sulfur. This is because sulfur (1s 2 2s 2 2p 6 3s 2 3p 4 ) has two electrons in one p orbital, whereas in phosphorus (1s 2 2s 2 2p 6 3s 2 3p 3 ) all the electrons are unpaired in 3 separate p orbitals. The paired 3p electrons repel each other and so are easier to remove than unpaired 3p electrons.