All the simple salts of the group 1 elements dissolve in water, producing ions, and
concequently the solutions conduct electricity. Since the Li+1 is small, it is expec-
ted that the solutions of lithium salts would conduct electricity better than the rest.
However, the trend is the opposite; Cs+ > Rb+ > K+ > Na+ > Li+ . The reason is that
the ions are hydrated in solution. Since Li+ is very small, it has higher attraction to
the water molecules and hence more water molecules adhere to its ions making it
highly hydrated. This makes the effective ionic radius large and hence it moves very
slowly . The contrast happens for the Cs+.
For a salt to dissolve, i.e NaCl, it must break into its constituents:
NaCl ➔ Na+ + Cl- (lattice energy )
If a salt is solluble, then its latice energy is smaller than the energy of hydration. The
solubility of most group 1 salts in water decreases on decending the group. This is
because, the lattice energy only changes very slightly but the free energy of hydration
changes drastically. For example, the differences in the lattice energy between NaCl
and KCl is 67 Kj/Mol, yet the difference in free Energy of hydration is 76 Kj/Mol.
Thus KCl is less soluble than NaCl.
N.B. The fluorides and carbonates have a trend that increases as one decends the
group.
difference between lithium and the other group 1 elements
Lithium chemistry is closer to group 2 elements (more so Magnesium) than it is to the
rest of group 1. In comparision, Li has higher melting point, it is harder, reacts least
readily with oxygen, its hydroxide is less basic and many of its salts are less stable,
and are more heavily hydrated than the rest of group 1 elements.
In general, you will realise that fisrt elements (Li, Be, B, C, N, O, and F) in almost
all groups differ from the rest. This is partly because the first element is always much
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smaller and is more likely to form covalent compounds (Fujan’s rule).
The similarity between the Li of group 1 and Mg of group 2 is called diagonal re-
lationship.
Such a relationship also exists between other other pairs of elements; Be and Al,
and B and Si. This relationship a rises partly because of the similarities if ionic sizes
between the pairs and majorly is due to similarities in their electronegativities.
Alkali Earth metals
The members of alkaline earth metals group are: Be, Mg, Ca, Sr, and Ba. The alk-
aline earth family is the second most reactive group. They are called alkaline earth
metals because they form “alkaline” solutions (hydroxides) when thay react with
water. This term “alkaline” means that the solution has a pH greater than seven and
is basic. Thus the alkaline earth metals form very basic solutions and are excellent
reducing agents.
The alkaline earths have two electrons in the outer shell. They have smaller atomic
radii than the alkali metals. The two valence electrons are not tightly bound to the
nucleus, so the alkaline earths readily lose the electrons to form divalent cations.
Alkaline earths have low electron affinities and low electronegativities. As with the
alkali metals, the properties depend on the ease with which electrons are lost.
Summary of Common Properties
Physical Properties:
- low densities, but higher than densities of the comparable alkali metal.
- Genarally, alkaline earth metals are silvery white metals with high melting
and boiling points.
- Are ofter stronger than most metals.
- All are ductile, lustrous, and malleable metals.
The metals of Group 2 are harder and denser than sodium and potassium, and have
higher melting points. These properties are due largely to the presence of two va-
lence electrons on each atom, which leads to stronger metallic bonding than occurs
in Group 1.
Three of these elements give characteristic colours when heated in a flame:
Mg brilliant white
Ca brick-red
Sr crimson
Ba apple green
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Atomic and ionic radii increase smoothly down the Group. The ionic radii are all
much smaller than the corresponding atomic radii. This is because the atom contains
two electrons in an s level relatively far from the nucleus, and it is these electrons
which are removed to form the ion. Remaining electrons are thus in levels closer to
the nucleus, and in addition the increased effective nuclear charge attracts the electrons
towards the nucleus and decreases the size of the ion.
Chemical Properties
The chemical properties of Group 2 elements are dominated by the strong reducing
power of the metals. The elements become increasingly electropositive on descending
the Group.
Once started, the reactions with oxygen and chlorine are vigorous:
2Mg(s) + O (g) ➔ 2MgO(s)
2
Ca(s) + Cl (g) ➔ CaCl (s)
2
2
All the metals except beryllium form oxides in air at room temperature which dulls
the surface of the metal. Barium is so reactive it is stored under oil.
All the metals except beryllium reduce water and dilute acids to hydrogen:
Mg(s) + 2H+(aq) ➔ Mg(aq) + H (g)
2
Magnesium reacts only slowly with water unless the water is boiling, but calcium
reacts rapidly even at room temperature, and forms a cloudy white suspension of
sparingly soluble calcium hydroxide.
Calcium, strontium and barium can reduce hydrogen gas when heated, forming the
hydride:
Ca(s) + H (g) ➔ CaH (s)
2
2
The hot metals are also sufficiently strong reducing agents to reduce nitrogen gas
and form nitrides:
3Mg(s) + N (g) ➔ Mg N (s)
2
3
2
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Magnesium can reduce, and burn in, carbon dioxide:
2Mg(s) + CO (g) à 2MgO(s) + C(s)
2
This means that magnesium fires cannot be extinguished using carbon dioxide fire
extinguishers
Oxides
The oxides of alkaline earth metals have the general formula MO and are basic.
They are normally prepared by heating the hydroxide or carbonate to release carbon
dioxide gas. They have high lattice enthalpies and melting points. Peroxides, MO ,
2
are known for all these elements except beryllium, as the Be2+ cation is too small to
accommodate the peroxide anion.
Hydroxides
Calcium, strontium and barium oxides react with water to form hydroxides:
CaO(s) + H O(l) ➔ Ca(OH) (s)
2
2
Calcium hydroxide is known as slaked lime. It is sparingly soluble in water and the
resulting mildly alkaline solution is known as lime water which is used to test for
the acidic gas carbon dioxide.
Halides
The Group 2 halides are normally found in the hydrated form. They are all ionic
except beryllium chloride. Anhydrous calcium chloride has such a strong affinity for
water it is used as a drying agent.
Oxidation States and lonisation Energies
In all their compounds these metals have an oxidation number of +2 and, with few
exceptions, their compounds are ionic. The reason for this can be seen by examination
of the electron configuration, which always has two electrons in an outer quantum
level. These electrons are relatively easy to remove, but removing the third electron
is much more difficult, as it is close to the nucleus and in a filled quantum shell.
This results in the formation of M2+. The ionisation energies reflect this electron
arrangement. The first two ionisation energies are relatively low, and the third very
much higher.
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Hydration energy
The hydration energy of the group II ions are four to five times greater than those of
group I. This is largely due to their smaller sizes and increased charge. The hydration
energy does decrease down the group as the size of the ions increase. In the case of
Be, a further factor is the very strong complex [Be(H O) ]2+ that is formed. Thus the
2
4
crystalline compounds of group II contain more water of crystallization than the crys-
talline compounds of group 1. Thus NaCl and KCl are unhydrous but MgCl .6H O,
2
2
CaCl .6H O, and BeCl .2H O have waters of crystallization. Note that the numbers
2
2
2
2
of the waters decrease as the ion becomes larger.
Solubility and Lattice energies
The solubility of salts decrease with increasing atomic weight, though the usual
trend is reversed in this group with regards to fluorides and hydroxides. Generally
the latice energy decrease as the size of the ion increases. The hydration energy also
decreases as the ion becomes larger. A decrease in lattice energy favours increased
dissolution, but a decrease in hydration energy favours decreased dissolution. With
most compounds, hydration energy changes more rapidly than lattice energy hence
the compounds become less soluble as the metal gets larger. However, for Flourides
and hydroxides, their lattice energies change more rapidly than hydration energy, and
so their solubility increases on descending the group. This explains why Ca and Mg
ions are the once causing water hardness.
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Block elements
Group 13
The Group 13 consists of five elements: boron (B), aluminum (Al), gallium (Ga),
indium (In), and thallium (TI) (Figure 1.1). These elements are characterized by
having three electrons in their outer energy levels, but reflect a wide range in the oc-
currence and behavior. Among them, only B is metalloid and plays a significant role
in plants. Aluminum, being one of the basic constituents of the lithosphere, reveals
amphoteric properties.
The elements of Group 13 are:
Name
symbol
electron configuration
boron
B
[He]2s22p1
aluminium
Al
[Ne]3s23p1
gallium
Ga
[Ar]3d104s2 4p1
indium
In
[Kr]4d105s2 5p1
thallium
Tl
[Xe]4f14 5d106s2 6p1
Physical Properties
1. Boron is a non-metallic grey powder.
2. All the other members of Group 13 are soft, silvery metals.
3. Thallium develops a bluish tinge on oxidation.
4. The densities of all the Group 13 elements are higher than those of Group 2
elements.
5. The melting points of all the elements are high, but the melting point of boron
is much higher than that of beryllium in Group 2, whereas the melting point
of aluminium is similar to that of magnesium in Group 2.
6. Their ionic radii are much smaller than the atomic radii. This is because the
atom contains three electrons in valence shell and when ionized the remaining
electrons are in levels closer to the nucleus. In addition, the increased effective
nuclear charge attracts the electrons towards the nucleus and decreases the
size of the ion.
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Chemical Properties
General Reactivity trend
The general trend down Group 13 is from non-metallic to metallic character. Boron
is a non-metal with a covalent network structure. The other elements are conside-
rably larger than boron and consequently are more ionic and metallic in character.
Aluminium is at borderline between ionic and covalent character in its compounds.
The remaining Group 13 members are generally considered to be metals, although
some compounds exhibit covalent characteristics.
The chemical properties of Group 13 elements reflect the increasingly metallic cha-
racter of descending members of the Group. Here only boron and aluminium will
be considered.
Boron is chemically unreactive except at high temperatures. Aluminium is a highly
reactive metal which is readily oxidised in air. This oxide coating is resistant to acids
but is moderately soluble in alkalis. Aluminium can therefore reduce strong alkalis,
a product being the tetrahydroxoaluminate ion, Al(OH) -. Aluminium also reacts
4
violently with iron (Ill) oxide to produce iron in the Thermit process:
2Al(s) + Fe O (s) ➔ 2Fe(l) + Al O (s)
2
3
2
3
Oxides
Boron oxide, B O , is an acidic oxide and an insoluble white solid with a very high
2
3
boiling point (over 2000K) because of its extended covalently-bonded network
structure. Aluminium oxide, Al O , is amphoteric.
2
3
Halides
The most important halide of boron is boron trifluoride, which is a gas. Aluminium
chloride, AlCl , is a volatile solid which sublimes at 458K. The vapour formed on
3
sublimation consists of an equilibrium mixture of monomers (AlCl ) and dimers
3
(Al Cl ). It is used to prepare the powerful and versatile reducing agent lithium te-
2
6
trahydridoaluminate, LiAlH .
4
Both boron chloride and aluminium chloride act as Lewis acids to a wide range of
electron-pair donors, and this has led to their widespread use as catalysts. Aluminium
chloride is used in the important Friedel-Crafts reaction.
Hydrides
Boron forms an extensive and interesting series of hydrides, the boranes. The simplest
of these is not BH as expected, but its dimer B H .
3
2
6
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Oxidation States and Ionisation Energies
Boron and aluminium occur only with oxidation number +3 in their compounds,
and with a few exceptions their compounds are best described as ionic. The electron
configuration shows three electrons outside a noble gas configuration, two in an s
shell and one in a p shell. The outermost p electron is easy to remove as it is furthest
from the nucleus and well shielded from the effective nuclear charge. The next two s
electrons are also relatively easy to remove. Removal of any further electrons disturbs
a filled quantum shell so is difficult. This is reflected in the ionization energies. The
first three ionization energies are low, and the fourth very much higher.
Inert pair effect
The term inert pair effect is often used in relation to the increasing stability of oxidation states that are 2 less than the group valency for the heavier elements of
groups 13, 14, 15 and 16. As an example in group 13 the +1 oxidation state of Tl is
the most stable and TlIII compounds comparatively rare. The stability increases in
the following sequence:
AlI < GaI < InI < TlI.
The situation in groups 14, 15 and 16 is that the stability trend is similar going down
the group, but for the heaviest members, e.g. lead, bismuth and polonium both oxi-
dation states are known.
The lower oxidation state in each of the elements in question has 2 valence electrons
in s orbitals. On the face of it a simple explanation could be that the valence electrons
in an s orbital are more tightly bound are of higher energy than electrons in p orbitals
and therefore less likely to be involved in bonding. Unfortunately this explanation
does not stand up.
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Group 14
The members of this group are C, Si, Ge, Sn, and Pb.
Physical properties
1. Going across a period, ionic radii decrease, however reaching group 14 (IVA),
anions decrease in size as you head towards the noble gases.
2. The atomic radii increases as we decend the group. The difference is size
between the Si and Ge is less than expected because Ge has a full 3d shell
which shields the nuclear charge rather ineffectively.
3. They show gradation downward. Carbon is non-metallic to Lead (Pb) whose
oxides are amphoteric but the element is mainly metallic.
4. ‘Diagonal line’ in the p-block that separates the metals from non-metals pas-
ses through Si and Ge. Indicating that Si is more non-metallic whereas Ge is
mainly metallic in behaviour. Thus both are considered metalliods.
5. The first four ionization energies suggest that the elements form +4 oxidation
state
6. They exhibit +4 oxidation state though +2 becomes stable as one decends the
group. In this case Pb resembles its metallic neighbours.
Chemical Properties
These elements are relatively unreactive, but the reactivity increases down the group.
The M2+ oxidation state becomes more stable with lower members. This unreactivity
is partly contribute by the surface coating of oxides or due to high overpotential of
reducing the H+ to H on the elements surface.
2
Reactions with water
C, Si, and Ge are unaffected by water. Sn reacts with steam to give SnO and H . Pb
2
2
is unaffected by water probably because of the oxide coating.
Reactions with acids
C, Si, and Ge are unaffected by dilute acids. Sn dissolves in dilute HNO , forming
3
Sn(NO ) . Pb dissolves slowly in dilute HCl, forming the sparingly soluble PbCl ,
3 2
2
though very readily in HNO , forming Pb(NO ) and oxides of nitrogen. Pb does not
3
3 2
dissolve in dilute H SO because of the surface coating of PbSO is formed.
2
4
4
Diamond (C) is unaffected by concentrated acids, but graphite reacts with both
HNO and HF/HNO . Si, too, is oxidized by HF/HNO . Ge dissolves slowly in hot
3
3
3
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concentrated H SO and in HNO . Sn dissolves in seversl concentrated acids. Pb does
2
4
3
not dissolve in concentrated HCl because during the reactiion, a surface coating of
PbCl is formed.
2
Inert pair effect
Here too, the inert pair effect shows itself increasingly with the lower members of the
group. There is a decrease in the stability of the (+4) oxidation state and an increase
in the (+2) state on decending the group. Ge(+2) is a strong reducing agent whereas
Ge(+4) is stable. Pb(+2) is ionic, stable and more common than Pb(+4), which is
oxidizing. The lower valencies are more ionic because of the radius of M2+ is greater
than that of M4+ and according to Fujans’ rule, the smaller the ion the greater is the
tendency for covalency.
Oxides and Oxoacids of Carbon and silicon
Unlike the later members of group 14, carbon forms stable, volatile monomeric oxides;
CO (in limited oxygen supply) and CO . Silicon is also stable as SiO .
2
2
(Exercise: In discussions with any module student list all the properties of CO and
CO that you may remember)
2
Carbon monoxide, with the chemical formula CO, is a colorless, odorless, and taste-
less gas. It is the product of the partial combustion of carbon-containing compounds.
Carbon monoxide has significant fuel value, burning in air with a characteristic blue
flame, producing carbon dioxide. Despite its serious toxicity, CO plays a highly useful
role in modern technology, being a precursor to a myriad of products.
O + 2 C → 2 CO
2
CO also is a byproduct of the reduction of metal oxide ores with carbon, shown in a
simplified form as follows:
MO + C → M + CO
A large variety of chemical reactions yield carbon dioxide, such as the reaction between
most acids and most metal carbonates. For example, the reaction between sulfuric
acid and calcium carbonate (limestone or chalk) is depicted below:
H SO + CaCO → CaSO + H CO
2
4
3
4
2
3
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The H CO then decomposes to water and CO . Foaming or bubbling, or both ac-
2
3
2
company such reactions. In industry such reactions are widespread because they
can be used to neutralize waste acid streams. The production of quicklime (CaO) a
chemical that has widespread use, from limestone by heating at about 850 °C also
produces CO :
2
CaCO → CaO + CO
3
2
The combustion of all carbon containing fuels, such as methane (natural gas), petro-
leum distillates (gasoline, diesel, kerosene, propane), but also of coal and wood, will
yield carbon dioxide and, in most cases, water. As an example the chemical reaction
between methane and oxygen is given below.
CH + 2 O → CO + 2 H O
4
2
2
2
Iron is reduced from its oxides with coke in a blast furnace, producing pig iron and
carbon dioxide:
2 Fe O + 3C → 4 Fe + 3 CO
2
3
2
Silicon Dioxide
The chemical compound silicon dioxide, also known as silica or silox (from the La-
tin «silex»), is the oxide of silicon, chemical formula SiO and has been known for
2
its hardness since the 9th century. Silica is most commonly found in nature as sand
or quartz. It is a principal component of most types of glass and substances such as
concrete.
Silicon dioxide is formed when silicon is exposed to oxygen (or air). A very thin layer
(approximately 1 nm or 10 Å) of so-called ‘native oxide’ is formed on the surface
when silicon is exposed to air under ambient conditions. Higher temperatures and
alternate environments are used to grow well-controlled layers of silicon dioxide on
silicon.
Thermal oxidation of silicon is easily achieved by heating the substrate to temperatures
typically in the range of 900-1200 degrees C. The atmosphere in the furnace where
oxidation takes place can either contain pure oxygen or water vapor. Both of these
molecules diffuse