Topic 3.3: Chemical properties

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

Group 0: Noble Gases

• Colourless
• Monatomic
• Unreactive

Their lack of reactivity can be explained by the inability of their atoms to lose or gain electrons.

Group 1: Alkali Metals

• Good conductors of electricity
• Low Density
• Shiny surface 
• Reactive metals
• Forms ionic compounds with non-metals

They form single charged ions M+. Their low ionization energies give an indication of the ease with which the outer electron is lost.

Reaction with Water

The alkali metals react with water to produce hydrogen and the metal hydroxide.

The reaction becomes more vigorous as we descend the group.


Group 7: Halogens

• Diatomic molecules
• Coloured
• Gradual change from gases (Fluorine and Chlorine) to Liquid (Bromine) to Solids (Iodine and Astatine)
• Reactive non-metals
• Reactivity decrease down the group
• Form ionic compounds with metals or covalent compounds with other non-metals

Reaction with Group 1 metals

Halogens react with Group 1 metals to form ionic halides. The halogen atoms gains one electron from the Group 1 elements to form a halide ion. The electrostatic force of attraction between the alkali metal ions and halides bonds the ions together.

Once the transfer is complete, the ions are pulled together by the mutual attraction of their opposite charges.

This would be the equation

Displacement reactions

The relative reactivity of the elements can also be seen by placing them in direct competition for an extra electron.

A chlorine nucleus has a stronger attraction for an electron than a bromine nucleus because of its smaller atomic radius and so takes the electron from the bromide ion. The chlorine has gained an electron and so forms the chloride ion. The bromide ion loses an electron to form bromine.

Reaction with Silver

The halogens form insoluble salts with silver. Adding a solution containing the halide to a solution containing silver ions produces a precipitate which is useful in identifying the halide

Silver Chloride = White

Silver Bromide = Cream

Silver Iodide = Yellow
Silver Nitrate = Colourless

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

The acid-base properties of the oxides are closely linked to their bonding. Metallic elements, which form ionic oxides, are basic; non-metal oxides, which are covalent are acidic. Aluminium oxide, which can be considered as an ionic oxide with some covalent character, shows amphoteric properties reacting with both acids and bases.

Sodium and Magnesium Oxide are bases

Na2O + HCl →2NaCl + H2O
MgO + H2SO4 → MgSO4 + H2O



Aluminium oxide is amphoteric (alkali and acid) so it reacts with both bases and acids


Al2O3 + 3H2SO4 → Al2(SO4)3 + 3H2O
Al2O3 + 2NaOH + 3H2O → 2NaAl(OH)4



Phosphorous Oxide, Sulfur Oxide, Dichlorine Oxide reacts with water to form Acid Solutions

P4O10 + 6H2O → 4H3PO4
P4O6 +6H2O → 4 H3PO3

SO3 + H2O → H2SO4
SO2 + H2O → H2SO3

Cl2O7 + H2O → 2HClO4
Cl2O + H2O → 2HClO


Silcon dioxide doesn't react with water, but reacts with concentrated alkalis to form silicates

SiO2 + 2OH → SiO3 + H2O