The “p” Block Elements
Lecture Notes 2
These are not polished notes. They are essentially transcripts of my lectures. I have
borrowed liberally from other books and the Internet. These are meant to help you
remember what was taught in the class. You are encouraged to read relevant sections of
the text book “Inorganic Chemistry” by Shiver, Atkins.
Physical Properties of the elements
(i) Melting points and enthalpies of atomization
There is no obvious trend, usually because the structures change as you go descend the
Group. But if the structures are the same then the melting points generally decrease down
the Group.
mp /ºC
C 3550
As 814
Si 1410
Sb 631
Ge 937
Bi 271
(II) Volatilities
As do down a Group for molecular compounds the boiling points tend to increase
because of increased atomic size and polarizabilities. Higher volatilities if molecular as
van der Waals forces are weak. Therefore larger dipole moments reduce volatilities due
to increased polarization, ie greater interaction with neighbouring molecules
VSEPR Theory
In the VSEPR model, regions of enhanced electron density take up positions as far apart
as possible, and the shape of the molecule is identified by referring to the locations of the
atoms in the resulting structure. The basic idea is that the enhanced electron density (lone
pair, bonded pair) will take up positions as far apart as possible so that the repulsions
between them are minimized.
The basic arrangement of the regions of electron density according to the VSEPR model:
No. of electron region
2
3
4
5
6
Arrangement
Linear
Trigonal Planar
Tetrahedral
Trigonal bipyramidal
Octahedral
A few generalizations on the basis of this theory are summarized below:
1. Lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
Questions:
A. Explain: CH4 (109.5˚), NH3 (107.3˚) and H2O (104.5˚)
B. Explain: CO2 (180˚), SO2 (119.5˚)
C. Explain: NO2+ (180˚), NO2 radical (134˚) and NO2- (115˚)
2. Repulsion of the bond pair decreases as the elctronegativity of the atom bonded to
the central atom increases
Explain: NH3 107.3˚, NF3 102˚
3. Lone pairs may sometimes be transferred from a filled shell of one atom to an
unfilled shell of another bonded atom.
Explain: PF3 (97.8˚), PH3 (93.8˚)
Exception to VSEPR: XeF6 is a distorted octahedron
MR3 (group 15) are pyramidal and lone pair sterically active.
Inversion energy (kJmol-1):
N(CH3)3
34
P(CH3)3
133
As(CH3)3
122
Sb(CH3)3
112
For the amines the process is very rapid, 10-11 per second. Only possible to characterize if
the rotation is slower. Thus AsMeEtPh has been resolved into optical isomers (inversion
energy 177).
(iii) Coordination Numbers
The radii of the atoms increases down a Group, therefore both the covalent and ionic
compounds tend to adopt higher coordination numbers.
e.g.
Fluorine
F-F
Chlorine
Cl-F
ClF3
ClF5
Bromine
BrF3
BrF5
Iodine
IF3
IF5
IF7
Similar effect seen for oxoanions:
Nitrogen
NO2 , NO3
Phosphorous
PO43Aresnic
AsO43Antimony
Sb(OH)6Heavier elements adopt higher coordination numbers by adopting polymeric structures.
S
O
S
O
O
Se
Se
O
Monomer
O
O
O
S
O
S
O
Chain
Ring
The tendency of a molecular compound to increase its coordination number also affects
its ability to function as a Lewis acid. (Lewis acid is an electron acceptor, a Lewis base
an electron donor) e.g. SiF4 is a strong Lewis acid as can increase its coordination
number to 5 or 6 to form compounds such as [SiF6]2-. CF4 is a not a Lewis acid. As you
go down the Group the Lewis acidity increases for the same coordination number as there
is more space to accommodate the incoming Lewis base.
SF6
SeF6 TeF6
–––––––––––––→
increasing Lewis acidity
TeF6 is a strong enough Lewis acid to form species such as [TeF7]- and [TeF8]2-.
The Lewis acid strength also gives a good indication of the ease of hydrolysis with water
acting as the Lewis base. SF6 is inert to water, but TeF6 is readily hydrolyzed. CF4 is
water stable, but SiCl4, GeCl4, and SnCl4 all fume (HCl) in moist air due to hydrolysis.
2nd row elements generally have a maximum coordination number of 4
3rd and 4th row elements favour a maximum of 6 and octahedral geometry, but lower
are still common
5th and 6th row can form higher coordination number compounds, e.g. [TeF7]-,
[TeF8]2- and IF7, [IF8]-.
(iv) Molecular Geometries
Know how to determine the molecular geometries for p-block compounds from VSEPR.
But why the following variation in bond angle?
H2O 105º
H2S 92º
H2Se 91º
H2Se 91º
This is one example of a collection of phenomena that are collectively known as “2nd row
anomalies”.
nd
The 2 row elements only have a 1s core, therefore the valence 2s and 2p orbitals are
very contracted. In addition the orbitals have similar maxima in their rdfs (radial
distribution functions) and are therefore able to hybridize effectively to form sp, sp2 and
sp3 hybrids for bond formation. (carbon chemistry). This introduces more s character into
the bonding. In subsequent rows the valence ns and np orbitals have more marked
differences in their rdfs, and therefore cannot hybridize so effectively. The result of this is
that the p-orbitals are predominantly involved in the bonding interactions, whereas the s
electron density is concentrated into the lone pair regions. Therefore, the X-E-X bond
angles want to adopt values best suited to the p-orbitals, ie 90º. The angle between the
lone pair and the X atoms also increases because of the increased s-character of the lone
pair.
Second row: both s and p orbitals are involved in both bonding and lone pairs
Third row and below: p-orbitals become concentrated in bonding, s-orbitals in the
lone pairs.
(v) Multiple Bonds
The 2s and 2p have similar radial distribution functions and can hybridize effectively.
The 2p orbitals of C to F can form effective multiple bonds. Therefore the compounds of
these elements posses some of the strongest bonds: (N2 is 945 kJmol-1 and CO is 1076
kJmol-1).
The overlap npπ- npπ for n>3 diminishes rapidly with n, therefore multiple bonding
becomes weaker as go down the group.
As a result of this it is not surprising that the 2nd row elements provide the largest number
of examples of small molecules with multiple bonds such as O2, N2, CO, CO2, NO etc.
C=C
Si=Si
Ge=Ge
Sn=Sn
602
310
270
190
N
P
As
Sb
N
P
As
Sb
C-C
Si-Si
Ge-Ge
Sn-Sn
356
226
188
151
N-N
P- P
As - As
Sb - Sb
945
493
380
293
O=O
S=S
Se=Se
Te=Te
513
430
290
218
167
209
180
142
O-O
S-S
Se-Se
Te-Te
144
226
172
149
The relative weakness of the N-N and O-O single bonds (as well as e-- e- repulsion
between the lone pairs on N and O) also leads to the occurrence of radical molecules with
unpaired electrons.
N2O4
2 NO2
Exists in equilibrium:
colorless
diamagnetic
brown, radical
paramegnetic
O
O
N
O
O
N
O
N
O
At 140 ºC dissociation is complete; at higher temp it dissociates into NO and O2
At low temperatures the dimer is formed, but the N-N bond is not sufficiently strong to
survive at higher temperatures. (Other examples include NO + NO2 to give N2O3 see
later)
Therefore :
2nd row multiple bonds (σ + π) prevalent Æ small molecules
3rd row single bonds become more prevalent (and higher) Æ polymers, clusters,
chains
This can be used to explain the allotropes observed for the 2nd and 3rd row elements
2nd row
CO2 (O=C=O)
N2 (N≡N triple bond)
O2 and O3 with O=O bonds
N2O3 (NO +NO2) with N=O
and N-N bonds
P
P
P
P
6 X P-P = 6 X 209 = 1254 kJmol-1
2X P P = 2 X 493 = 986 kJmol-1
3rd row
SiO2 (infinite glass like structure of
tetrahedral SiO4 units)
White (P4), red and black phosphorus, with
P-P single bonds
S8 and other rings/chains with S-S single
bonds
P2O3 comprised of P4O6
O
P
P
O
P
O
O
O
P
O
P
O
O
P
O
O
O
P
O
O
P O
O
O
P4O10
P4O6
Multiple bonded heavier elements difficult to isolate as they polymerize, e.g. Si2H4,
Si2H2. The multiple bonds can be stabilized by replacing them with organic groups that
are sterically demanding.
R
R
Si
Si
R
R
E+
X
R
R
Si
Si
R
R
E
X
Si
Si
Si
Si
The multiple bonds that are formed are weaker than the ones in 2nd row. This is seen by
the bond contraction data. Also, the geometry changes down the group as the molecules
become less planar.
R
R
R
X
X
R
H3C-CH3
H2C=CH2
R2HSi-SiHR2
R2Si=SiR2
R3Ge-GeR3
R2Ge=GeR2
R3Sn-SnR3
R2Sn=SnR2
R
θ
X
R
R
D(X-X) pm
154
133
235
214
245
235
282
276
X
R
R
X
X
R
R
% contraction
θ
13.6
R = mesityl
9.0
R = CH(SiMe3)2
4.0
R = CH(SiMe3)2
2.0
0
R
32
41
The three centered 2 electron bond of boranes. B2H6 has only 12 valance electrons .
However it shows structure as below.
H
H
H
B
H
B
H
H
H
H
H
B
H
B
H
H
(vi) Catenation
Ability of an element to form element-element bonds is called catenation. Clearly very
important and significant for C and life!
Orders of catenation:
Group 13 B > Al > In > Ga > Tl
Group 14 C > Si ≈ Ge > Sn ≈ Pb
Group 15 N < P > As > Sb > Bi (difference to Group 13 and 14 due to weak N-N bond)
Group 16 O < S > Se > Te > Po
For Group 17 other than elemental forms only I3−, I3+, I5−, I5+.
For 2nd row catenation only really important for B and C
For 3rd row Si, P & S
Less important for 4th, 5th and 6th row elements
Vertical Trends in Chemical Properties of Compounds
Acid Base Properties
Bronsted acidities
The acid strength of hydrides (EH) in aqueous solution is affected by the following
factors
Strength of E-H bond
Electronegativity of E
Energy of solvation of [EHn-1]−, small anions have more favourable solvation enthalpies.
Acid strength is often indicated by pKa values, where pKa = -log10Ka, and Ka is the
equilibrium constant for the formation of [EHn-1]− and H+ from EHn.
pKa values of EHn main group compounds
Group 14
Group 15
~58
39
~35
27
25
~19
20
~15
Group 16
14
7
4
3
Group 17
3
-7
-9
-10
Decrease in pKa (increase in acid strength) due to increase in the electronegativity of E
For oxo-acids the pKa is determined primarily by the number of oxo groups (E=O), as
additional oxo groups withdraw more electron density from central atom, therefore
encouraging greater dissociation.
pKa
HClO
4.53
HClO2
2.0
HClO3
-1.0
HClO4
-8
pKa
H2SO3
1.81
H2SeO3
2.46
H2TeO3
2.48
pKa
H2SO4
-3
H2SeO4
-3
Te(OH)6 {H2TeO4}
7.7
Oxides become less acidic down a Group
Oxides of PIII and AsIII are acidic
Oxide of SbIII is amphoteric
Oxide of BiIII is basic
Lewis Acidity and Basicity
Empty pz
Br B
Lone
pair
H
Br
Br
Br
Br
Br
B
N
N
H
H
H
H
H
Lewis acid
order of base strength
BBr3
NH3 > PH3 > AsH3 >> SbH3
Increasing s-orbital character in lone pair as go down group
due to changes in hybridization. For NH3 lone pair is sp
hybrid, but for lower members essentially s orbital. As the
s-orbital contracts down the Group, less well set up to
donate electron density to Lewis acid.
AlBr3
Me3N > Me3P > Me3As > Me2O > Me2S > Me2Se > Me2Te
(as above)
H+
PMe3 > PH3 > PF3
NMe3 > NH3 > NF3
Due to inductive effects as Me groups act as electron
donors, whilst the fluoride is electron withdrawing.
If keep the base constant and compare order of acid strength the following is observed.
Lewis base
order of acid strength
NMe3
SbF5 > AsF5 > PF5
Changes in Lewis acidity due to increased size of central
atom, therefore easier to accommodate incoming base.
NMe3
AlCl3 > AlBr3 > AlI3
Changes in Lewis acidity due to electronegativity of
halogen.
NMe3
BI3 > BBr3 BCl3 > BF3
The order for acid strength is reversed to that expected on
the basis of electronegativities, therefore something else
must have been going on.
We have seen earlier that the 2nd row elements are the best set-up for pπ- pπ bonding. In
the boron halides there is some/partial π character in the bonding as the full p-orbitals on
the halides can donate some electron density to the empty pz orbital on the central B
atom. The pπ- pπ interaction is greatest for B-F, and weakest for B-I bonds. The greatest
π bonding will occur for planar molecules. Therefore, BF3 will suffer the greatest loss of
π character, and hence has more to lose from acting as a Lewis acid. This will be
compounded as the same pz orbital is used for both π bonding and Lewis acid behaviour.
For Group 14
Order of Lewis acidity
SiF4 > SiCl4 > SiBr4 > SiI4
i.e. substituent electronegativity now important if no or reduced π bonding. Also more
space around SiF4 than SiI4, so combination of steric and electronic effects.
If the halogen is replaced by methyl or aryl substituents, then the Lewis acidity is much
reduced, and in some cases effectively removed. Alkyl and aryl groups can be considered
as electron donating, whilst halogen as electron withdrawing groups.
eg [SnCl6]2- is well known, whilst [SnMe6]2- is unknown, although both SnCl4 and SnMe4
are known.
This also has consequences for rate of hydrolysis as SnCl4 is readily hydrolysed, whereas
both SnMe4 and SnPh4 are unreactive to water under normal conditions.
Effect of Lewis Acids on structures
Lewis acidity can have a profound effect on the solid state structure if some compounds.
If a molecular compound MXn does not have a high Lewis acidity it does not attempt to
expand its coordination number by accepting electron pairs from a second molecule. Ie it
is coordinatively saturated.
The extent of coordinative unsaturation (at metal centre) is dependent on:
Positive charge
Availability of empty orbitals
Availability of space
The last two dominate.
If coordinative unsaturation is relieved by formation of one or two dative bonds then
resulting compound is either oligomeric or polymeric. If all the lone pairs on the anion
become involved in dative bond formation then resulting structure is an infinite polymer.
AlF3
Infinite solid
SiF4
Molecular
PF5
Molecular
Increase in coordinative unsaturation (compensates +ve charge increase)
CF4, SiF4, GeF4
Monomer, tetrahedral
SnF4, PbF4
Infinite polymeric structure, Octahedral
Increase in coordinative unsaturation, increase in radius
Some reactivity trend: not strictly vertical
Can compare the effect on the chemical properties of empty orbitals and lone pairs by
using a series of Group 13, 14 and 15 compounds.
A good example set are:
Me3In
Me4Sn
Me3Sb
All have similar M-C bond strengths and are thermodynamically unstable with respect to
both oxidation and hydrolysis as the oxides (In2O3, SnO2, and Sb2O3) and hydroxides
(In(OH)3, Sn(OH)4 and Sb(OH)3) are all thermodynamically more stable. But what about
the kinetics and activation energies?
Compound
Me3In
Me4Sn
Me3Sb
Reactivity to Air
Pyrophoric
Inert
Pyrophoric
Reactivity to water
Readily Hydrolysed
Inert
Inert
Hydrolysis requires an empty orbital on the central metal atom for the initial nucleophilic
attack (which is presumably rate determining)
Oxidation requires pre-coordination of O2 which since it has both filled and empty
frontier orbitals (i.e. the HOMO and LUMO are both the πu orbitals). Therefore,
oxidation can be achieved either if the molecule has a low lying empty acceptor orbital or
a filled lone pair orbital.
Me In
Me
Me
Me
Sn
Me
Me Me
Me
Sb
Me
Me
For InMe3 the empty p-orbital perpendicular to the molecular plane can function as an
effective Lewis acid site to both O2 and H2O, therefore highly reactive towards both air
and water.
In SnMe4 there is no empty p-orbital or lone pair, therefore inert to both water and air. In
the case of SnCl4 the Lewis acidity is greater due to the lower electron density on the Sn.
In SbMe3 there are no empty p-orbitals, therefore inert to nuclephilic attack by H2O, but
the lone pair is able to function as a Lewis base towards O2, therefore pyrophoric in air.