The
general properties of the main group elements of the Periodic Table
Group
1 and 2 elements are known as the s-block elements. The s-block metals are characterised by their reactivity, low
density and their softness.
The
metals in group 1 are collectively known as the alkali metals.
There are six of them: lithium, sodium, potassium, rubidium, caesium and
francium. They are the most
reactive of all the s-block metals and are usually stored under oil due to their
reactive nature. These elements are
too reactive to find uncombined and are normally extracted from their compounds
by electrolysis. The electrolysis
of aqueous solutions produces the hydroxide.
These six elements are all remarkably similar in terms of both their
chemical and physical properties. They
are all white , soft, lustrous but
rapidly tarnish in air. Group 1
elements are good conductors of both heat and electricity. Good conductors of both electricity and heat.
Good reducing agents and have low electronegativity values due to being
very electropositive therefore tend to lose their electrons relatively easily.
Their valency is the same as their group number.
Group 1 elements all form the M2O under normal conditions.
Sodium peroxide(Na2O2). Metals
below sodium all form hyperoxides (or superoxides) when heated in oxygen.
Group 2: Alkaline Earth
Metals
Known
as the alkaline earth metals. They
have similar properties to those of group 1.
They are good reducing agents. Due
to an increase in nuclear charge they have low electronegativity values higher
than group 1 metals in the same period. They show an oxidation number of +2 in their compounds.
Form positive ions with a 2+ charge.
The Alkaline Earth Metals all form the oxide MO when heated in oxygen.
Barium however forms the peroxide(BaO2) when heated to between 500-800°C.
This thermally decomposes above 800°C to the normal oxide.
If
the s-block elements are heated with chlorine they always produce the chloride
MCl and MCl2 accordingly.
With
the exceptions of Be and Mg, the s-block elements react with cold water to
produce hydrogen and metal hydroxide. Group
1 oxides being basic react with water to form the metal hydroxide:
M20(S)
+ H2O(l)
= 2MOH(aq)
Group
2 oxides (excepted beryllium) react with water rapidly, again forming the metal
hydroxide:
MO(s) +
H2O(l) =
M(OH)2(aq)
Lithium
and beryllium often exhibit non-characteristic properties due to their compounds
having largely covalent character.
Group 3
Group
3 metals are softer and have lower melting points than the transitional metals.
They are less reactive than group 1 and 2 metals and show two valencys
except aluminium. Group 3 chlorides
are easily hydrolysed as they have a lacking outer shells and so can act as a
Lewis acids accepting lone pairs of electrons from water into the vacant
orbital.
Group
3 element Boron forms Boric (III) acid B(OH)3 which is usually written as H3BO3
because of its weakly acidic nature. Boron
oxide (B2O3) is also weakly acidic, reacting with water to form boric (III)
acid.
Aluminium
hydroxide Al(OH)3 is amphoteric:
Al(OH)3
(S) +
3HNO3 (aq) = Al(NO3)3
(aq) +
3H2O(l)
Al(OH)3
(s) +
NaOH (aq) = NaAl(OH)4
(aq)
Aluminium
oxide is also amphoteric. It is
insoluble in water.
Al2O3
(s) +
3H3SO4 (aq) =
Al2(SO4)3 (aq) + 3H2O (l)
Al2O3
(s) +
2NaOH (aq)
= 2NaAl(OH)4 (aq) + 3H2O (l)
Group
4
This
is a group of elements which shows a trend from non-metals at the top to metals
at the bottom. As group 4 is
descended the increase in metallic character is particularly noticeable.
This is because with increasing atomic radius and thus a decreasing
attraction for the outer electrons. This
cause increasing delocalisation, which is characteristic of metals.
The characteristic oxidation states of the group are +4 and +2.
The +2 oxidation state becomes increasingly more stable as the group is
descended, Lead and Tin highlight this.
The
element Carbon forms two oxides, Carbon monoxide and Carbon dioxide.
Carbon dioxide is weakly acid and forms carbonic acid upon addition to
water.
CO2 (aq) + H2O (l) = H2CO3
(aq)
Silicon
dioxide forms a giant molecular compound which is weakly acidic:
SiO2
(s) + 2NaOH (aq) = Na2SiO3 (aq) +
H2O (l)
Lead
forms three oxides Pb, PbO2 and Pb3O4. Each
oxide exhibits amphoteric nature:
PbO + 2H+ = Pb
2+ +
H2O
PbO + OH- +
H2O =
Pb(OH)3 -
PbO2 + 4H+(conc)
= Pb 4+
+ 2H2O
PbO2 + 2OH- = PbO3
2- + H2O
PbO4
behaves as a mixture of PbO and PbO2.
Group 7: Halogens
Group7
is known as the halogens (salt formers) as they readily combine with metals to
form salts. They are the most
reactive of non-metals. All
halogens exist as diatomic molecules. These
molecules exist in gaseous, liquid and solid states.
The hydrogen halides are soluble in water and all, with the exception of
HF, form strong acids. HF forms a
weak acid as the strength of the HF bond only allows partial dissociation.
The halogens are all coloured and the depth of colour increases with
increasing atomic number. They react with metals to form ionic compounds e.g. NaCl.
With non-metals they tend to form covalent compounds.
In these the halogen is linked by a single covalent bond to the other
element. Halogen disproportionation
is the simultaneous oxidation and reduction of the same species in solution.
Chlorine has an oxidation state 0 and disproportionates when dissolved in
water:
Cl2 (g) + H2O (l) = HClO (aq) +
HCl.
(Cl
oxidation (Cl oxidation
number +1)
number -1)
When
chlorine is dissolved in alkali the chlorate (I) ion is formed:
Cl2
(g) + 2OH- (aq) = ClO- (aq) + Cl- (aq)
+ H2O (l)
The
chlorate (I) ion disproportionates slowly upon standing:
3ClO- (aq) = 2Cl-
(aq) +
ClO3 (aq).
Group
7 elements are strong oxidising agents. The
oxidising power of the halogens decreases as the group is descended.
During the extraction of bromine from the sea water chlorine oxidises the
bromide ions to bromine and is itself reduced to chloride ions:
Cl2 (g) + 2Br- (aq)
= 2Cl- (aq) +
Br2 (g).
Group 0: Noble Gases
Known as the noble gases. They were called the ‘inert gases’ until it was discovered that krypton and xenon can form compounds. They have a full outer shell and therefore are stable and very unreactive. The noble gases exist as single atoms in the gas phase at room temperature. These symmetrical non-polar atoms have no permanent dipole and do not form normal bonds. If the temperature is low enough they will condense to liquids and form solids, this suggests the presents of intermolecular forces.
Anne Marie Mc Ferran: Loreto College, Coleraine.