The Main Group Elements (H&S pages 236-502).
The main group elements discussed here are groups 1 to 8 shown below:
1
2
1b
2b
3
4
5
6
7
H
8
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
……
Cu
Zn
Ga
Ge
As
Se
Br
Kr
Rb
Sr
……
Ag
Cd
In
Sn
Sb
Te
I
Xe
Cs
Ba
……
Au
Hg
Tl
Pb
Bi
The properties of these elements can be understood in terms of a few simple concepts.
These are:
1.
2.
3.
4.
The charge on the Lewis acid
The ionic radius of the Lewis acid
Electronegativity
The role of the inert pair in the heavy post-transition elements such as Hg, Tl, Pb,
and Bi.
5. Hard and soft acids and bases
6. Relativistic effects, that are largely responsible for effects 3 to 5.
The elements on the left hand side of the periodic table form cations that have largely
ionic bonding. Relativistic effects are not important here, and these elements are
classified as hard in Pearson’s HSAB classification. Recall that the HSAB principle states
that ‘Hard acids prefer hard bases, and soft acids prefer soft bases’.
Bases have donor atoms that occur on the right hand side of the periodic table. Such
bases (ligands) have unshared pairs of donor atoms that they can donate to Lewis acids.
The donor atoms produce hard or soft bases in the periodic table as shown:
SOFT ←
C
↓
SOFT ←
N
O
→ HARD
F
↑
P
S
Cl
As
Se
Br
Te
I
↓
SOFT
The acids we are considering here are classified into hard and soft as shown below:
1
2
……
1b
2b
3
4
5
H
Li
Be
B
blue
=
Na
Mg
Al
Si
black =
intermediate
K
Ca
……
Cu
Zn
Ga
Ge
red
soft
Rb
Sr
……
Ag
Cd
In
Sn
Sb
Cs
Ba
……
Au
Hg
Tl
Pb
Bi
=
hard
The chemistry of hard acids is dominated by considerations of size and charge. It is
generally true that the smaller the size of the metal ion, and the higher its positive charge,
the higher is the positive charge density of that Lewis acid. The higher the charge density,
the greater is the ability of the Lewis acid to attract the negative charge. Thus, for metal
ions of the same charge and differing size down a group we have:
Metal ion:
Be(II)
Mg(II)
Ca(II)
Sr(II)
Ba(II)
Ionic radius (Å):
0.27
0.74
1.00
1.18
1.36
Log K1(OH-):
8.4
2.6
1.1
0.9
0.7
Log K1(F-):
4.82
1.82
1.1
0.8
0.7
For metal ions of the same size but differing charge we have:
Metal ion:
Li(I)
Mg(II)
In(III)
Zr(IV)
Ionic radius (Å):
0.7
0.74
0.80
0.84
Log K1(OH-):
<0
2.6
10.0
14.6
Log K1(F-):
<0
1.82
4.6
9.8
What we see is that for hard metal ions the smaller they are, and the higher their cationic
charge, the stronger Lewis acids they are.
The closer an element is to gold in the periodic table, the softer it is. For soft metal ions,
their affinity for ligands is governed by their electronegativity. This can completely override the effects of size and charge. Thus, we see that the affinity of Hg(II) and Ag(I) for
soft Lewis bases is enormous, in spite of their large size and fairly low charge. Thus,
compared to the similarly sized hard ions Ca(II) and Na(I) we have:
Metal ion:
Ca(II)
Hg(II)
Na(I)
Ag(I)
Ionic radius (Å):
1.00
1.03
1.00
1.05
Log K1(OH-):
1.1
10.6
<0
3.32
Log K1(F-):
1.1
1.5
<0
0.37
Log K1(I-):
<0
13.5
<0
6.58
The inert pair effect causes the heavy post-transition elements (Tl, Pb, Bi) to have as their
most stable oxidation states those that are two less than the group oxidation state. It is due
to the high electronegativity of these elements that a pair of electrons is retained. The
inert pair occurs as follows:
Group
Oxidation
State:
2
3
4
5
Zn
Ga
Ge
As
Se
Br
Kr
Cd
In
I
Xe
Tl
I
Sb
(III)
Bi
III
Te
Hg
(O)
Sn
(II)
Pb
II
← stable oxidation states
Ordinarily, one would expect elements to have as their most stable oxidation state the
group oxidation state. Thus, for Ga and In the trivalent state is the most stable state, and
the monovalent state is found only in a few unstable solid state compounds such as GaCl
and InCl, as well as AlCl. However, for Tl the monovalent state is by far the more stable
oxidation state, and Tl(III)I3 is, for example, an unknown compound. Bi(V) is known in
only one or two compounds of doubtful validity. The resistance of Hg metal to oxidation,
and its existence as a liquid at room temperature, can be viewed as a manifestation of the
inert pair causing it to hold on to its electrons and remain as the metal.
Coordination Number.
Coordination number (C.N.) is determined largely by metal ion size, and also to some
extent by metal ion charge. Larger metal ions tend to have higher coordination numbers.
Thus, if we look at the group 2 metal ions we see that the preferred C.N.’s are as follows:
Metal ion:
Be
Mg
Ca
Sr
Ba
Ionic radius (Å):
0.27
0.74
1.00
1.18
1.36
Coordination number:
4
6
6/7
8
9
Group 1, Hydrogen (H&S chapter 9).
Hydrogen is the simplest element, consisting of a single proton and electron. It has a
reasonably high electronegativity of 2.1, which means that it forms covalent bonds with
carbon which has a similar electronegativity of 2.5. Once it has ionized to form a proton,
it has no remaining electrons, and, in theory has an ionic radius of zero. In fact, it is never
a bare proton, and always retains some electron density, but still has a very small size
where it is cationic. This gives it a very high charge density, and the proton is a very
strong Lewis acid. This can be seen in its affinity for some Lewis bases:
H+(aq)
+
OH-(aq)
=
H2O(l)
log K = 14.0
H+(aq)
+
F-(aq)
=
HF(aq)
log K = 3.2
H+(aq)
+
NH3(aq)
=
NH4+(aq)
log K = 9.22
It is, however, fairly hard in the HSAB sense, so that we find that it has high affinity for a
hard ligand such as F-, but virtually no affinity for soft halide ions such as Cl-, Br-, or I-.
The small proton in aqueous solution forms the linear two-coordinate [H2O-H-OH2]+ ion
in solution, shown below. This low coordination number is expected from the small size
of the proton.
Figure 1. The hydronium ion.
The electronegativity of H is high enough that it can form a negative hydride anion, H-. It
is less electronegative even than iodine however, and so the H- anion is not very stable.
Thus, H- reacts with water to give H2 and NaOH:
NaH (aq)
+
H2 O
=
NaOH (aq)
+
H2(g)
The saline (salt-like) hydrides are formed by the group 1 and 2 metal ions, e.g. NaH or
CaH2. Here the hydride ion resembles an F- anion in salts such as NaF, and has about the
same ionic radius as F-. Many other more electronegative metals form more covalent
hydrides, such as [Al2H6] or transition metals ions that have covalent bonds to H, as will
be discussed later.
One of the most important properties of the proton is its ability to form H-bonds when
attached to more electronegative donor atoms such as F, O, or N. The H-bond holds water
together, and is the key to molecular structure in life, helping to hold proteins together,
and controlling molecular recognition.
Figure 2. Hydrogen bonding of four water
molecules around a central water molecule.
The Alkali Metal Ions (Group 1). Li, Na, K, Rb, Cs (H&S Chapter 10).
The alkali metals are very reactive, and react violently with water to give the metal
hydroxide and H2 gas. The standard reduction potentials are very negative in accord with
this:
Li+(aq)
Li(s)
+
+
e
=
Li(s)
H2O
=
Li+(aq) +
Eo
=
OH-(aq) +
-3.04 V
H2(g)
Because of their low charge and large size, the ability of the group1 metal ions to form
complexes in solution is limited. Thus, the metal hydroxides are completely ionized to
give metal cations and hydroxide ions. They are therefore strong bases. The coordination
numbers increase with increasing metal ion size:
Metal ion:
Li+
Na+
K+
Rb+
Cs+
Ionic radius (Å):
Coord. No.:
0.76
4-6
1.02
6-7
1.38
6-8
1.52
8-9
1.67
8-9
A four coordinate complex of Li+ is seen below with four THF (tetrahydrofuran)
molecules attached to the Li:
Figure 3. Structure of the Li(I) complex with four
THF (tetrahydrofuran) molecules coordinated to the Li.
The low electronegativity of the alkali metals means that they are very hard in the HSAB
sense, and their chemistry is largely that of being bound to the hard oxygen donor atoms,
as seen for [Li(THF)4]+ above. The most important aspect of their chemistry is their
ability to bind to crown ethers and cryptands. The crown ethers were discovered in 1967
by Charles Pedersen when he was working at DuPont. These are cyclic polyethers called
macrocycles (‘large cycles’). Some examples of crown ethers and cryptands are shown
below:
O
O
O
O
O
O
O
O
O
O
O
18-crown-6
15-crown-5
O
N O
O
O
O
O N
O
cryptand-222
N
O
O
O
12-crown-5
O
N
O
O
O
O
cryptand-221
Figure 4. Cryptands and crown ethers.
The important aspect of the crown ethers was that these complexed alkali metal cations in
solution. Up until that time it was considered that the alkali metal ions had very little
ability to form complexes in aqueous solution. This was important, because ion channels
in cell membranes allowed K+ and Na+ to pass through selectively, and the properties of
the crown ethers suggested how this might be achieved. The striking feature of crown
ethers was their ability to complex alkali metal ions selectively on the basis of their size.
Thus, the log K1 values for 18-crown-6 with alkali metal ions varies in aqueous solution
as shown below. The diagram shows that 18-crown-6 has a definite preference for the K+
ion. This can be understood by looking at space-filling drawing of 18-crown-6, and how
the K+ cation can fit into the cavity in the ligand.
Figure 5. Variation in log K1 for 18-crown-6 complexes
as a function of metal ion radius for alkali metal ions.
Figure 6. The D3d conformer of the free
18-crown-6 ligand, and its complex with
K+, showing how well the K+ cation fits
into the cavity of the ligand.
The alkali Earth Metals (group 2). (H&S Chapter 11)
The alkali earth metal ions resemble the alkali metal ions in having a low
electronegativity, and being very hard in the HSAB classification. The big difference,
though, is their charge, which makes them stronger Lewis acids. The effect of charge for
log K1 for hard metal ions with EDTA, all having an ionic radius of about 1.0 Å, makes
this point:
Metal ion:
Ionic radius (Å):
Log K1 (EDTA):
Na+
Ca2+
La3+
Th4+
1.02
1.86
1.00
10.65
1.03
15.36
0.94
23.2
We thus find that the metal ions in Group 2 are much better at complexing with ligands
than are those in Group 1. Being hard, complexing of Group 2 cations is confined largely
to oxygen donors, and to nitrogens, more so where the nitrogen donors are part of a
ligand that also has some oxygen donors, such as in EDTA.
O
O
-
O
O
N
OO-
N
O
O
EDTA
With Group 2, we come to the small Be2+ ion, which is a strong Lewis acid because of its
small size. Thus, we have affinity for F- and OH- as follows:
Metal ion:
Be2+
Mg2+
Ca2+
Sr2+
Ba2+
Ionic radius (Å):
Coord. No.:
0.27
4
0.74
6
1.00
7
1.18
8
1.36
9
Log K1(OH-):
8.4
2.6
1.1
0.9
0.7
Log K1(F-):
4.82
1.82
1.1
0.8
0.7
Log K1(EDTA):
9.7
8.79
10.65
8.72
7.88
What one is seeing with log K1 for EDTA is the effect of coordination number. Be(II)
and Mg(II) are stronger Lewis acids than Ca2+ with unidentate ligands such as OH- or F,
but with hexadentate EDTA the large Ca2+ is better able to accommodate the large EDTA
ligand. The structure of 8-coordinate Ca(II) in [Ca(EDTA)(H2O)2]2- is seen below:
Figure 7. Structure of the EDTA complex
of Ca2+ (CSD code: MUJFEC).
The alkali earth metal ions Ca2+, and particularly Sr2+, and Ba2+ are large enough to fit
well into the cavities of crown ethers and cryptands, and actually form more stable
complexes than large alkali metal ions. Thus, we can compare log K1 values with some
crown ethers and cryptands for Ba2+ and K+, which are almost identical in size:
Ligand:
18-crown-6
15-crown-5
cryptand-222
Log K1(K+):
2.05
0.75
5.5
Log K1(Ba2+):
3.89
1.71
9.6
Thus, even with these ligands, the charge on the metal ion has an effect on complex
stability.
Group 3. B, Al, Ga, In, Tl. (H&S Chapter 12).
In group 3 the electronegativity of the metals is getting a bit higher, and the heavier
metals Ga, In, and Tl are actually post-transition elements (they are close to Au), so have
much higher electronegativity and a very different chemistry from B and Al. They form
trivalent cations that form very strong complexes, and for the first time we have to
consider the inert pair effect. Thus, for Tl, the most stable oxidation state is not Tl(III)
but Tl(I). The Tl(I) ion has an ionic radius of 1.50 Å, and so resembles K+ and Rb+ to
some extent in its chemistry. It does have some tendency towards covalence (it is soft),
and so forms many complexes where it is bound to soft donors such as S. Below is seen
the complex of Tl(I) with the sulfur-donor macrocycle 9-ane-S3.
Figure 8. Structure of the Tl(I) complex
with the S-donor macrocycle 9-ane-S3.
The group 1 metal ions such as K+ and Rb+ are too hard to form such complexes with Sdonor ligands. Tl(III) is very easily reduced to Tl(I), as can be seen from the following
potential:
Tl3+(aq)
+
2e
=
Tl+ (aq)
Eo = +1.82 V.
The Tl(III) ion is stabilized by complexation with ligands, and is an extremely powerful
Lewis acid. This can be seen by examining the log K1(OH-) values for the group 3 M3+
ions:
Metal ion:
Al3+
Ga3+
In3+
Tl3+
Log K1(OH-):
8.45
10.6
9.47
12.8
Log K1(F-):
6.42
4.47
3.74
2.6
Log K1 (Cl-):
-1.0
0.01
2.32
6.72
HARD ←
→ SOFT
Tl(III) is a very soft metal ion in the HSAB classification. This is seen in the high affinity
for the soft chloride ligand. At the other end of the group, Al3+ is very hard, and has very
little affinity for Cl-. However, with the hard F- ion, the order of preference is reversed.
Boron is very different in its chemistry from the other members of the group. While they
all have preferred coordination numbers of 6, with occasional higher coordination
numbers of 7 or 8, boron always has a coordination number of four or less. Thus, B(III)
in aqueous solution exists as B(OH)3(aq) at lower pH, and is too acidic to ever be
protonated to yield a B3+(aq) ion. At higher pH (9.1) a water coordinated to B(OH)3(aq)
ionizes to yield the borate anion:
B(OH)3.OH2(aq)
“boric acid”
=
[B(OH)4]- (aq)
borate anion
+
H+(aq)
pKa = 9.1
This behavior is readily understood in terms of the small size (ionic radius = 0.11 Å) and
high charge on B(III).
B(III) forms compounds of considerable covalency, and forms a reasonably stable
hydride, as in Li[BH4], lithium borohydride. Here we have a Td [BH4]- anion, which is
used in chemistry as a mild reductant. The chemistry of the boranes, those compounds
involving boron and hydrogen, is enormous. The structure of [B2H6] is shown below.
Figure 9. B2H6, showing the bridging H-atoms,
which donate electron density to the adjacent B atom.
Group 4. C, Si, Ge, Sn, Pb. (H&S Chapter 14).
Here the group valency is four. The electronegativity of the elements has risen quite high,
with the C atom having an electronegativity of 2.5. The C-C bond is quite strong
compared to the C-O bond. Thus, carbon forms stable organic compounds based on the
stability of the C-C. The small carbon atom has a coordination number of four in
compounds such as CH4 or CCl4. This drops to three in sp2 hybridized C atoms such as
are present in CH2=CH2, or two in sp hybridized C in acetylene (H-C≡C-H). The bonding
in compounds containing C, Si, and Ge, have quite covalent bonding, and forms
compounds of considerable stability. In the elements, these all have a diamond-like
structure with covalent C-C, Si-Si, or Ge-Ge bonding. None of these ions forms an M4+
cation in solution. Carbon forms the CO32- and HCO3- anions at higher pH, and at lower
pH (<6) breaks up to form CO2(g). Si and Ge form many compounds with a coordination
number of four, such as SiCl4 or GeCl4. They also readily expand their coordination
numbers to six, as in complexes such as [SiF6]2- and [GeF6]2-.
The high electronegativity of these elements leads to a strong inert pair in Sn and Pb.
For Sn both the Sn(IV) abd Sn(II) state are relatively stable. For Pb, the Pb(IV) state is of
rather low stability. Important Pb(IV) compounds are PbO2, which is important in the
lead/acid battery, and Pb(CH2CH3)4 (tetraethyl lead), which used to be added to gasoline
to prevent ‘knock’ (premature ignition on compression). The lead/acid battery works on
the cell:
PbO2(s) + 4 H+(aq)
+
Pb(s) =
2 Pb2+(aq)
+ 2 H2O
Eo =
+1.2 V
The Pb(II) and Sn(II) ions display a sterically active inert pair, which means that in
structures of the complexes of these cations, there is usually a gap in the coordination
geometry which is occupied by the lone pair. This resembles the structure of NH3 as
predicted by VSEPR, where the structure is derived from a tetrahedron, with one site
occupied by the lone pair. This is seen in the structures below of the [SnCl3]- and the
[Pb(C6H5)3]- anions:
Figure 10. Structure of a) [SnCl3]-, and b) [Pb(C6H5)3]-, showing
the positions occupied by the stereochemically active loan pairs.
The Pb2+ ion is about the same size as Sr2+, but because of its electronegativity, due to
relativistic effects, it forms much more stable complexes. Pb2+ forms the most stable
complexes of any metak ion with crown ethers. Some examples are:
Ligand:
OH-
EDTA
Log K1(Pb2+):
Log K1(Sr2+):
6.0
0.7
18.0
8.72
18-crown-6 cryptand-222
4.1
1.5
12.0
5.7
Group 5. N, P, As, Sb, Bi. (Chapter 15 H&S).
Here electronegativity increases further, with nitrogen having E. N. = 3.0. Elemental N
and P are not metallic, and As and Sb are somewhat metallic. Bi is a genuine metal. The
N2 molecule is held together by triple bonds of enormous strength. N, P, As, and Sb have
both trivalent and pentavalent stable oxidation states. In the trivalent state, following
VSEPR rules, these elements have a lone pair that enables them to act as ligands, as in
NH3, and P(CH3)3. These compounds are bases. NH3 is quite hard, while the less
electronegative P(CH3)3 and other phosphines such as P(C6H5)3 or P(OCH2CH3)3 are
soft. Examples of complexes where these compounds act as ligands are:
Figure 11. Complexes with ammonia and phosphines.
a) is [Ni(NH3)6]2+, and b) is trans-[Pt(P(CH3)3)2Cl2]
The pentavalent state is found in the oxides that form the acids HNO3, H3PO4, and
H3AsO4. These compounds are acidic because they have such high positive charge
density on the central atoms that they are able to drive all the protons off any oxide
groups attached to them, To summarise this aspect, the oxides of the first-row elements in
aqueous solution are classified as follows:
Group:
1
2
Oxide:
Li2O
BeO
In water:
LiOH
Oxide:
In water:
3
B2O3
4
5
CO2
N2O5
Be(OH)2 H[B(OH)4]
H2CO3
HNO3
Na2O
MgO
SiO2
P2O5
NaOH
Mg(OH)2 Al(OH)3
Al2O3
amphoteric/ weak
Strong
weak
acid/
[Si(OH)4] H3PO4
weak
6
7
-
-
SO3
Cl2O7
H2SO4
HClO4
strong/
weak strong
very
strong
Bases
base
amphoteric
acids
acid
acid
acid
We can color-code the oxides from basic to acidic as follows:
Group
1
2
3
Li2O
BeO
B2O3 CO2
N2O5
Na2O MgO Al2O3 SiO2
P2O5
K2O
CaO
4
5
6
7
SO3
Cl2O7
Ga2O3 GeO2 As2O5 SeO3 Br2O7
Rb2O SrO
In2O3 SnO2 Sb2O5 TeO3 I2O7
Cs2O BaO
Tl2O3 PbO
Strong base
Weak base
Bi2O3
amphoteric
weak acid
strong acid
Bismuth has no well-defined Bi(V) complexes. In its Bi(III) state it resembles the isoelectronic (has same electronic configuration) Pb(II) ion, and has an inert pair of
electrons that may be stereochemically active. Because of its trivalent state, Bi(III) forms
complexes of considerably higher stability than does Pb(II). Thus, log K1(OH-) for Bi(III)
is 12.8 as compared to 6.0 for Pb(II). For the EDTA complexes the log K1 values are 18.0
for Pb(II), and 27.0 for Bi(III).
Group 6. O, S, Se, Te. (Chapter 16 H&S).
Oxygen is the second most electronegative element, with E.N. = 3.5. We are now truly
into purely non-metallic elements down the whole group. The importance of oxygen and
sulfur to us here is primarily their importance as donor atoms for ligands,
Group 7. The Halides. (Chapter 17 H&S).
Here electronegativity reaches a maximum. Fluoride is the hardest ligand, while Cl-, Brand I- range from intermediate to soft. In their highest oxidation state, as in ClO4(perchlorate) the halogens form the strongest oxo-acids (e.g. HClO4) because of the high
formal charge of the central atom.