Download Variation in Properties of Group II Compounds

Yip Wai Sum 7B (12)
97-AL Q.9
Each group of elements embodied in the periodic table has their own unique
properties. As for group II elements, they are classified as one of the s-block elements,
also named as alkaline earth metals. In this essay, the variation in properties of group
II elements and their compounds are illustrated.
Variation in properties of group II elements and their compounds include both
physical and chemical properties.
Variation in Physical Properties of the Elements
Variation in Atomic and Ionic radii
There is a general increase in atomic and ionic radii on descending group II. The
increase in both radii is due to the addition of one more electron shell on going down
the group.
However the ionic radius is always smaller than the atomic radius. Group II
atoms tend to lose their outermost s-electrons to form ions. As a result, the cation has
one electron shell less than the atom.
Variation in Ionization Enthalpies
The alkaline earth metals have generally low first and second ionization
enthalpies. The two s-electrons are relatively easy to remove, as they are well
screened from the nucleus by inner shells electrons.
Down group II, the first and second ionization enthalpies decrease. As the atomic
radii increase down the group, the outermost shell electrons are far away from the
nucleus, and the effective nuclear charge decreases.
Variation in Electronegativity
The electronegativity of an atom provides a numerical measurement of the power
of that atom to attract electrons. Generally, the electronegativity decreases down
group II. As the atomic radius increases down the group, the effective nuclear charge
decreases, the atoms have less electron-attracting power, and therefore decrease in
electronegativity.
Variation in Melting Point
There is a general decrease in melting point down group II elements. Melting
point is interpreted in terms of the strength of metallic bonds. The metallic bond
strength depends on ionic radius, number of electrons in the outermost shell and
packing of the metallic crystal.
As mentioned, the ionic radius increases down group II. As the number of
outermost shell electrons is always the same — 2, there is a decrease in metallic bond
strength on moving down the group.
Chemical Properties of the Elements
All the metals in group II are high in the reactivity series. They are strong
reducing agents. The reducing power increases down the group, as the atomic size
increases, it becomes easier to remove their outermost electron and the I.E. decreases
accordingly.
Reaction with Oxygen
Group II elements generally form only the normal oxidesO2- because the size of
the M2+ is small and the charge is high. The polarizing power of the metal ions
decreases down the group. They form more stable oxides with O2-.
(Can peroxides and superoxides form? Give appropriate equations for this part.)
Reaction with Chlorine
Reactivity increases down the group. As the size of atom increases, the I.E. of the
atom decreases. The ΔH formation of chloride becomes more exothermic.
As a consequence of high I.E. of beryllium, its chloride is essentially covalent,
with comparatively low melting point, non-conducting and the chloride can dissolve
in organic solvents. The lower members in group II form essentially ionic chlorides,
for example MgCl2, CaCl2, SrCl2, BaCl2, with high melting points, conducting
electricity and soluble in water but not in organic solvents.
Reaction with hydrogen
All group II elements except beryllium react directly with hydrogen at
temperatures between 300 and 700℃ to form hydrides,
e.g. Mg(s) + H2 (g) —> MgH2(s)
The reactivity increases with increasing reducing power down the group.
Reaction with water
Group II elements reduce water to produce hydroxide and hydrogen, except
beryllium.
Magnesium reacts slowly in cold water but rapidly in steam. The rest react with
increasing rapidity on descending the group, to form hydroxide and hydrogen.
e.g. Ca(s) + 2H2O(l) —> Ca(OH)2(aq) + H2(g)
Variation in Properties of Group II Compounds
Reaction of Oxides
The oxides of group II elements react readily with water to form alkaline
solutions. The reactivity of oxides towards water increases down the group. For
example, BeO has a high degree of covalent characters, is inert and almost insoluble
in water and acid; MgO is almost inert towards water but dissolves in acids to give
salts. Other group II oxides dissolve in water with increasing ease down the group.
The resulting solutions are slightly alkaline due to the reaction between oxides and
water.
e.g. CaO(s) + H2O(l) —> Ca2+(aq) + 2OH-(aq)
Basicity of group II peroxides also increase down the group so that barium
peroxide reacts readily with ice-cold water to form hydrogen peroxide.
e.g. BaO2(s) + 2H2O(l) —> Ba(OH)2(aq) + H2O2(l)
Reaction of Hydrides
Relative Thermal Stability of Carbonates and Hydroxides
The thermal stability of carbonates and hydroxides of group II metals increases
down the group. Down the group, as the size of the cation increases, the polarizing
power of the cation decreases. The compound with large anion becomes more stable.
(Illustrate with equations the thermal decomposition of carbonates and hydroxides.)
Relative solubility of sulphates and hydroxides
Solubility of group II sulphates decreases down the group. The cations are much
smaller than anions. Down the group, the change in size of the cations does not cause
a significant change in the ΔH lattice, but ΔH hydration becomes less negative. As a
result, ΔH solution becomes less negative and the solubility decreases down the
group.
However the solubility of group II hydroxides increases on descending the group.
The cations and anions are of similar sizes. Down the group, less energy is required to
break the lattice as the size of cations increases, but the change in ΔH hydration is
comparatively small. So, the ΔH solution becomes more negative and the solubility
increases down the group.
In conclusion, the variations in both physical and chemical properties of group II
elements and its compounds generally follow a trend.