Periodicity

List the characteristic properties of transition elements

Explain why Sc and Zn are not considered to be transition elements

Explain the existence of variable oxidation number in ions of transition elements

Describe and explain the formation of complexes of d-block elements

Define the term ligand

Explain why some complexes of d-block elements are coloured

State examples of the catalytic action of transition elements and their compounds

Outline the econimic significance of catalysts in the Contact and Haber processes

Elements are placed in order of increasing atomic no. (Z)

The rows are the periods and the columns are the groups

The period refers to the number of energy levels occupied and the group, to the number of electrons in the valence shell

The number of electrons in the valence shell gives an element its group number

Explain the physical states (under standard conditions) and electrical conductivity (in molten state) of the chlorides and oxides of the elements in Period 3 in terms of their bonding and structure

Describe the reactions of chlorine and the chlorides referred to in 13.1.1 with water

The fiirst ionization energy of an element is the energy required to remove one mole of electrons fron on mole of gaseous atoms.electronegativity

Electronegativity is the ability of an atom to attract electrons in a covalent bond

The effective nuclear charge felt by an atom's valence electrons increases with the group number of the element. it increases across a period but stays approximately the same down the group

Sheilding is where an electron is sheilded from the full attrcation of the nucleus

Li --> Cs (down the alkali metals)

Atomic radius increases due to increased electron shielding.
Ionic radius increases due to increased electron shielding.

Ionisation energy decreases due to increased electron shielding.

Melting/boiling point decreases due to increased electron shielding --> decreased forces.

Electronegativity decreases due to increased shielding --> decreased attraction for outer electrons.

F --> I (Down the halogens):

Atomic radius increases due to increased electron shielding.

Ionic radius increases due to increased electron shielding.

Ionisation energy decreases due to increased electron shielding.

Melting/boiling point increases due to increased number of electrons->increased london dispersion forces.

Electronegativity decreases due to increased shielding -> decreased attraction for outer electrons.

Na --> Ar (across period 3):

Atomic radius decreases due to increased nuclear charge --> greater attraction for electrons.

Ionic radius decreases Na --> Al (due to increased nuclear charge) jumps Al --> Si (due to reversal of ionisation direction...increased electron-electron repulsion) decreases Si --> Ar (due to increased nuclear charge).

Ionisation energy increases due to increased nuclear charge.

Electronegativity increases due to increased nuclear charge --> greater attraction for electrons.

Compare the relative electronegativity values of two or more elements based on their positions in the Periodic Table

Largely determined by the number of electrons in the valence shell of an element

Discuss the similarities and differences in the chemical properties of elements in the same group

Reactions of elements in the same group are similar because they have identical outer shells

Look at link 3.3.1

Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across Period 3

Elements on the left are metallic, right are non-metals, Al is a metalloid (semi-metal).

Look at link 3.3.2