CH2201 – Main Group Chemistry
The Periodic Table
The periodic table serves several purposes, as a method of systemising the elements into periods
and groups with similar physical and chemical traits it allows the behaviours of a given element to be
predicted based on its neighbours behaviour in an analogous situation.
The main group can be considered to include the s-block (Groups I and II) and the p-block (Groups
XIII to XVIII) elements.
Trends in the Periodic Table
Atomic radius – atoms will be smaller the further along a row as the increase in effective nuclear
charge (due to addition of a proton) brings the electron cloud closer to the nucleus. Atoms will be
larger the further down a group as each row represents a new electron shell, increasing the shielding
of the outer electrons from the nucleus, leading to an expansion in size.
The relation between effective nuclear charge and the principle atomic number is;
E ∝Z∗
n2
Ionisation energy – the further across a group the higher the ionisation energy as these smaller
atoms are again more able to pull electron density towards themselves. Elements on the left of the
periodic table are more likely to lose an electron to achieve the octet and so are easy to ionise.
Ionisation energy decreases down the group as the shielding effect makes it easier to remove
electrons from the outer orbital of the larger elements.
There are however breaks in this trend, the ionisation of Group XIII metals are as follow;
There is a drop as expected from boron to aluminium; however there is then a small
rise between aluminium and gallium. This is due to the intervention of the first row
transition metals, in essence the effective nuclear charge is given a large series over
which to increase. A similar explanation is used for the increase in ionisation between
indium and thallium, this time due to the intervention of the lanthanides. This is
called the alternation effect.
The d-orbitals filled across the transition row provide very little shielding to outermost electrons and
so the increase in effective nuclear charge is the primary effect.
Electronegativity – the further along a row and the higher up in each group, the higher the
electronegativity, this is due to the atoms being smaller, thus being better able to withdraw electron
density towards them.
The same variation occurs as above though the trend is less smooth, there is a rise in
electronegativity between aluminium and gallium as the increased nuclear charge increases the
ability of the atom to withdraw electron density towards itself. The first element of the group is
always the most electronegative, this can lead to important chemical differences (i.e. the position of
M(g) M(g) +
B 798 Al 578 Ga 579 In 556 Tl 590
the delta positive/negative charges in the bond, in a C-H bond the C is δ- however in Si-H it is the H
that is δ-, the Si is δ+).
Electronegativity is measured by several different scales, these include;
Mulliken – proposed that the arithmetic mean of the first ionisation energy and activation energy
should be a measure of the tendency of an atom to withdraw an electron towards itself. It is termed
‘absolute electronegativity’ as it is not dependant on an arbitrary relative scale.
χ =I. E + E. A
540
The Mulliken electronegativity can therefore only be calculated for elements of which the first
ionisation energy is already known.
Pauling – the original proposition of electronegativity based upon the observation that the covalent
bond between A-B was stronger than expected given the strength of A-A bonds and B-B bonds. This
was explained as a result of the contribution of ionic canonical forms to the bonding and is
calculated thus;
χa − χb = 0.017 D A − B −1
2 (D A − A + D(B − B)
Since only differences in electronegativity are defined a reference state is required, this is given by
hydrogen which has a fixed electronegativity of 2.20 on the Pauling scale.
To calculate the Pauling electronegativity it is necessary to have information on the dissociation
energies of at least two types of covalent bond formed by that element.
Allred-Rochow – proposed that electronegativity should be related to the charge experienced by an
electron on the ‘surface’ of an atom, the higher the charge per unit area of atom surface the greater
its tendency to attract electrons. The effective nuclear charge can be calculated using Slater’s rules
while the surface area of an atom can be taken as the covalent radius squared, expressed in
Angstroms.
χ = 0.704 +0.359 x Z2
r2
Covalent radius – the main group elements of the p-block show less tendency to form Mn+ cations,
much preferring covalency. Small cations eg. B3+ or C4+ do not exist due to the very powerful electric
field generated by the ion, this polarises neighbouring anions perturbing the electron distribution
and thus forming electron-bond pairs. As the ionic radius increases the polarising power will
decrease.
To predict the behaviour of a bond Fajan’s rules can be utilised;
A small positive ion favours covalency
A large negative ion favours covalency
A large charge on either or both ions favours covalency
Polarisation and hence covalency is favoured if the positive ion does not have a noble gas
configuration
The variation of stability of the MI/II to MIII/IV states for group 13/14 is well known. The MI/II state
increases in stability upon descending the group, the MIII/IV state decreases in stability upon
descending the group (B→Tl, C→Pb).
Boron is normally found as BIII, aluminium is found as AlIII but AlI is known, for gallium and indium the
3+oxidation state is still predominant but the 1+ state is increasingly stable (and both GaI and InI are
powerful reducing agents) whereas thallium TlI is predominant and TlIII is a powerful oxidising agent.
Carbon is normally found as CIV, silicon is prevalently found as SiIV but SiII is known, germanium GeII is
more stable than SiII, tin SnII and SnIV are of comparable stability and for lead, PbII is predominant.
For the group III elements there is no special trend in terms of the sum of the second and third
ionisation energies (corresponding to removing the 2 s2 electrons) and so the term ‘inert pair effect’
is in fact a misleading one.
Instead consider the reactions in terms of the Born-Haber cycles for the formation of MX and MX3
M+ 1/2X2→MX
M(s)→M(g) ∆Hatom
1/2X2(g)→X(g) 1/2∆Hdiss
M(g)+X(g)→MX(g) ∆Hbond
If ∆Hbond (the exothermic term) is greater than the sum of∆Hatom and 1/2∆Hdiss (the endothermic
terms) then the reaction will proceed.
And for MX3
M+ 1½ X2→MX3
M(s)→M(g) ∆Hatom
M(g)→M*(g) ∆Hprom
1½ X2(g)→X(g) 1½ ∆Hdiss
M(g)+3X(g)→MX3(g) 3∆Hbond
Introduction of an additional endothermic term (the promotion of an electron to an excited state)
influences the reaction.
M(g)→M*(g)
s2px1→s1px
1py1
If 3∆Hbond is greater than the sum of the 3 endothermic terms the reaction will proceed. It is only
feasible to form MX3 if sufficient energy from bond formation is present to offset the energy
required for promotion. For elements of higher atomic number the promotion energies do not
s px py pz s px py pz
change significantly but energies regained from bond formation decrease, as a result the lower
oxidation states become more stable upon descending the group.
The same result is obtained by lattice energy calculations for ionic compounds.
This effect is important in the chemistry of many main group elements as it governs the relative
stability of each oxidation state.
The Unique Properties of the First Row Elements
1) Small size
Elements of the first row are able to form strong pπ-pπ bonds (i.e.C=C is much stronger than Si=Si,
also C=O compared to the practically unheard of Si=O bond).
For σ bonds the p-orbitals bonding overlap is at a minimum when they are apart, they peak when
one lobe of each orbital overlaps (with the same sign, +ve or –ve), and then hits a second minima
when they are fully overlapped as there is anti-bonding
behaviour between the two sets of oppositely aligned
orbitals.
For π bonds the minima is again seen when the two orbitals
are apart, there is an increase as you bring them together,
and finally a maximum when the two are fully overlapped as
there are two bonding interactions between the two sets of
orbital lobes.
Since the first row elements are smaller than second row
elements they can form much more efficient overlaps and as
such form much stronger π bonds.
The small size of first row elements can also lead to inter-electronic repulsion effects, take the
following series of bond strengths (in kJmol-1);
N-O 163 Big step then
steady decrease
C-F 485 Same trend
observed
O-H 467 Doesn’t follow trend
P-O 368 Si-F 582 S-H 347
As-O 331 Ge-F 473 Se-H 246
2)The Inter-Electronic Repulsive Effect
The inter-electronic repulsion effect occurs between first row elements and elements with lone pairs
of electrons, weakening the bond. The small first row elements are able to minimise the internuclear
distance leading to a strong repulsion between the non-bonding pairs of electrons. It is this property
that allows the use of hydrazine as the principle component in some rocket fuels, with a ∆Gof of
+149.7kJmol-1 and being a light molecule a large amount of energy is released per gram.
σ
π
Distance
Overlap
Integral
3)Unable to Perform Hypervalency
The first row elements have a maximum coordination number of 4, for the second row elements the
coordination number isn’t restricted as above the occupied 3p orbitals are easily accessed 3d
orbitals, since there are no 2d orbitals and the energy gap between 2p and 3d is too large the same
promotion cannot be achieved by first row elements.
When describing the shapes of molecules a commonly employed model is that of hybridisation (sp3
tetrahedral, sp2 trigonal planar etc.), these available d-orbitals are also subject to hybridisation (in
turn allowing hypervalency through occupation of the hybrid d-orbital). An sp3d2 orbital will give an
octahedral geometry, sp3d a trigonal bypyramidal geometry. d-orbitals may also be used in the
formation of π-bonds (i.e. POX3).
These properties also effect structure and reactivity, for example N(CH3)3 displays a different
geometry to the planar N(SiH3)3 structure due to π-bonding d-orbital overlap. Also CCl4 is less
reactive than SiCl4 due to the available coordination sites on the Si allowing the formation of new
bonds to proceed or occur simultaneously as the breaking of the Si-Cl bonds.
Group XIII
Boron
Boron displays some very interesting chemistry; most important is the chemistry of boron halides.
These halides are monomeric (where other group XIII halides tend to dimerise) and strongly Lewis
acidic due to the empty valence orbital;
BF3 is such a strong acid it can react with HF to form H+BF4- and will form adducts with both NH3 and
H2O (an adduct is the product of a direct addition of two or more distinct molecules, incorporating
all atoms of all components leading to a net reduce in bond multiplicity in at least one of the
reactants through the formation of two chemical bonds) and it can also act as the electrophile in a
Friedel-Craft style reaction where;
BF3+RF→Rδ+FBF3δ-
The acceptor strength of the borohalides does not vary as expected, on the basis of electronegativity
it would be expected that BF3>BCl3>BBr3>BCl3 since it would be sensible to suggest that boron is
more electropositive in BF3 than BCl3 and so forth and thus be more able to accept electrons
however this fails to take into account the effect of π-bonding. In BF3 the result of this π-bonding is
the strongest ‘single’ bond known (rB+F 0.13nm observed, 0.15nm predicted).
B F
F
F
This overlap leads to a bonding molecular orbital of π symmetry. As the halides get larger down the
group this effect becomes weaker as the overlap becomes less efficient.
Aluminium
Aluminium trifluorides are predominantly ionic in nature, forming complex lattice structures in
which the aluminium is 6-coordinate;
Al
F
Al
F
F
F
F
F
F
Al
Al
F
F
F
F
F
F
F F
F
F
F
F
F
Al
Al
Al
As the halides are descended this behaviour changes, AlCl3 is 6-coordinated as with AlF3, however
upon melting at 192oC the structure becomes that of a 4-coordinate dimer and upon further heating
these dimers dissociate to form the 3-coordinate monomer. AlBr3 and AlI3 are both then 4-
coordinate dimers.
M
ClCl
Cl
M
Cl Cl
Cl
The dimeric form is a common structure for group XIII halides. The bridge bonding in a dimeric tri-
halide is not the same as in alkyls or hydrides, instead involving a 3-centre 4-electron bond of
reasonable strength (30-40kJ mol-1).
The trihalides of all group XIII elements are Lewis acids of varying strength and can form
coordination complexes with upto 3 charged or uncharged donor ligands, as such they are widely
used as starting materials for the synthesis of other derivatives.
The dihalides are known, but very rarely contain MII centres, most are in fact ionic in their behaviour,
i.e. GaX2 is more accurately represented as Ga+[GaCl4-].
The monohalides are most stable for thallium, although they are also known for gallium and indium
Thallium
Thallium is the heaviest of the group XIII elements and also displays interesting properties. Due to
the inert pair effect TlIII is a very powerful oxidising agent, so much so that TlIII is not the state found
in TlI3 as it may be expected as the iodine would instantly be oxidised to form I2 and 2 electrons, the
two species cannot co-exist. Instead TlI3 contains TlI and the tri-iodide ion [I-I-I]-.
Due to its similar size and identical charge, TlI will act very much like Ag+, K+ and Rb+ in terms of
solubility and reactivity.
As a result of this thallium is very toxic as, within the body, it can replace K+ accelerating or disabling
enzyme action.
Group XIV
Carbon
Carbon has 3 known allotropes;
Graphite
Diamond
Fullerenes
Graphite is most commonly found in impure and finely divided forms such as soot, lampblack and
charcoal. Charcoal can be converted into ‘activated carbon’ (also known as ‘active’ carbon), this is
carbon with a very large, adsorbent surface area (just one gram can have a surface area of 500m2)
suitable for performing reactions upon. It is converted by passing steam, air or CO2 at elevated
temperatures through the charcoal. It can be further treated to enhance the adsorbent properties.
Under an electron microscope it can be seen that there are individual particles convoluted,
displaying various kinds of porosity with flattened areas (similar to graphite) separated by only a few
nanometres, providing the correct environment for adsorption.
Graphite in its purest form is found as an offset lamellar structure of hexagonal rings (formed from
sp2 carbon in trigonal planar arrangements with angles of 120o) with the layers aligned as ABABAB.
The remaining electrons are held in multicentre molecular orbitals derived from the 2pz orbitals on
the carbon, extending across the lattice. This leads to electron delocalisation and thus graphite is a
conductive material (in a similar mechanism as with metals). The measurement of the magnetic
susceptibility of finely divided graphite will allow the average radius of the the path of one electron
to be determined, imperfections in the structure will give rise to magnetic fields that repel the
electrons from their path. The average radius of the electron path is 30 rings.
The van der Waals forces holding the layers together are relatively weak and so they are free to
slide over each other. It is this characteristic trait that leads to two of graphite’s important
commercial uses, as a lubricant and as pencil ‘lead’.
Graphite undergoes few chemical reactions. Carbon can be fluorinated to form (CF)X
F
F
F
FF
F
F
F
F
F
F
F
C(graphite) + F2 (CF)x
This reaction is problematic when considering the synthesis of pure F2 via the electrolysis of Na3AlF6.
The carbon anodes gradually corrode due to surface layers of (CF)x forming and dropping off under
the heat (600-1000oC).
(CF)6→CnF2n+2
In the presence of HF a different product is formed.
F2+HF+C(graphite)→(C4F)x (at room temperature)
This preserves the layer structure of graphite;
F
Since very few carbons are ‘lost’ there are still a sufficient number of singly occupied 2pz orbitals
present to form delocalised molecular orbitals. The C-F bond involves a 2pz orbital and so removes
conductivity but still allows for the layers to slide.
Graphite also reacts with the vapours of alkali metals to form intercalation compounds.
K
K
K
K K
K
C(graphite) + K(g) C8K
K0.54nm
(Also C24K, C36K, C48K, C60K)
K
The bonding involves transfer of electrons from K to vacant M.O’s in graphite resulting in an ionic
type of structure. The layers in the intercalary species are slightly further apart (0.54nm compared to
0.335nm) not only due to the potassium ions placed between layers, but also since the layers change
alignment, becoming an AAAA structure, leading to an increase in repulsion.
This C8K can then undergo several reactions, if water is rapidly added the mixture will explode,
however if it is added in a controlled manner KOH and H2 are formed. It will also react with MXn
(where M is a transition metal) to form C8nM and nKX. There are many graphite intercalates formed
with small molecules, for example AlCl3, HF, CuBr2 and more. The formulae of these intercalates is
not always well defined;
Graphite + Conc. H2SO4 → C24+HSO4
-.2H2SO4
Diamond has a cubic unit cell, it is a 3d structural form with each carbon arranged in a tetrahedron
with 4 others and each carbon with 4 sp3 hybridised orbitals. Diamond is thermodynamically less
stable than graphite, the favoured allotrope, (CDiamond→CGraphite ∆H=-1.90kj mol-1).
Diamond has two main uses that take advantage of its appearance and hardness;
Jewellery – prized for rarity, coloured diamonds are the result of impurities in the lattice
Cutting tools/abrasives – on the Moh’s scale Diamond has a hardness of 10, the hardest
naturally occurring substance known.
Due to its wear resistance and optical properties diamond is also used in bearings, laser optics,
resistors, thermistors, radiation detection equipment and wire dies (to form very precise wire
gauges).
Fullerenes are only a recently discovered allotrope of carbon but possess the potential to
revolutionise reaction mechanics. Carbon nanotubes (or ‘Buckytubes’, after Buckminster Fuller, the
designer of buildings of a similar geodesic dome structures in the early 1900’s) can be used as
reaction chambers, small enough to only allow one molecule of each reactant contact at once.
Boron-Nitrogen Compounds
There are several methods of forming BN and several potential resultant structures.
Na2B4O7.10H2O + NH4Cl → BN
B(OH)3 + (NH2)2CO → BN (in the presence of NH3, at 500-950oC)
BCl3 + NH3 → BN (700oC)
Among the structures possible is hexagonal BN (similar to graphite)
BN
BN
BN
BN
B
N
NB
B
NB
N
NB
N
BN
BN
BN
BN
B
N
NB
B
NB
N
NB
N
BN
BN
BN
BN
B
N
NB
B
NB
N
NB
NN
B N
B
NB
B N
B
N
B
NB
N
B
N
B
NBNB
N
This is a colourless insulator and an effective lubricant, it is chemically inert to most reagents but will
react with fluorine or hydrogen fluoride.
HF+BN→NH4+BF4
-
F2+BN→BF3+N2
The second possible structure is of cubic BN (similar to diamond)
B
NB
N
BN
B
NB
N
B
N
N
NB
N
BN
B
BN
N
These structures analogous to carbon structures can be rationalised considering the components,
boron and nitrogen have covalent radii of 88pm and 70pm respectively, compared to 77pm for
carbon, it seems sensible therefore that boron and nitrogen can form similar structures to those of
carbon.
Boron-nitrogen compounds may also take on a ring structure, similar to benzene, known as
borazine.
H3B NH3
200oC
B3N3H6
B
N
B
N
B
N
H
H
H
H
H
H
The similarity in structure in this case is accompanied with a change in chemistry, where before the
covalent radii explained similarities it is now the electronegativity that plays an important role.
Boron and nitrogen have electronegativities of 2.0 and 3.0 respectively, whereas carbon has an
electronegativity of 2.5 (and obviously no difference between electronegativities in a C-C bond),
therefore where a benzene ring is susceptible to electrophilic attack a borazine ring is in fact
susceptible to nucleophilic attack.
B
N
B
N
B
N
HN
HB
NH
BH
NH
HB B
N
B
N
B
N
H
H
Cl
H
H H
Cl
H
H
H
Cl H
B
N
B
N
B
N
Cl
Cl
H
Cl
H H
sp2
sp3
HCl B
N
B
N
B
N
R
R
H
R
H H
LiR
Boron-Oxygen Compounds
Boron shows a strong affinity for oxygen, most likely due to π-bonding.
Boron can form both weak and strong acids. Boric acid (B(OH)3) is weak, however, boric acid when
mixed with glycol will create a strong acid (i.e. propylene glycol and boric acid form a non-toxic
version of anti-freeze).
HO(CH2)2OH + B(OH)3 →
Another boron-oxygen compound, B2O3 (which contains sp2 Boron) is important in the glass industry.
B
O
B
O
B
O O
B
O
B
O
B
O
O
B
O
B
O
B
O
SP2 BORON
Boron-oxygen compounds can be formed of BO33-
units, BO45- units or a combination of the two.
Silicon
Oxides of silicon are the main area of study when considering the chemistry of silicon.
Silica, SiO2, is known to have over 22 phases and at least a dozen polymorphs are known of the pure
compound. Silica has many commercial uses, including;
Silica gel
Quartz
-piezoelectric devices
-crystal oscillators
-frequency controllers
Vitreous silica
-glassware (boro-glasses absorb UV, vitreous silica glasses do not, they therefore each have
their uses)
Fumed silica
-thickening agent
-reinforcing filler
Diatomaceous earth (formed from the skeletons of ‘diatoms’, single celled animals), used in
filtrations
There are a wide range of naturally occurring silicates known, often with complex formulae. The
structural makeup of silicates has been extensively studied by X-ray crystallography and the basic
unit discovered to be a tetrahedron, a silicon atom surrounded by 4 oxygen atoms.
Silicate structures are then made up of shared corners on adjacent tetrahedrons, they can take on
island, sheet, ribbon or chain structures.
Island;
These structures with ionic lattices correspond to harder minerals, such as zircon and garnet. They
are discrete anions of defined size.
Ribbon/Chain;
SiO32- SiO3
2-Si4O11
6- Si2O52-
These are the structures that give rise to fibrous materials such as asbestos.
Sheet;
Si2O52-
These structures give minerals which can cleave along sheet boundaries, such as mica.
There are also very large classes of related solids in which some of the silicon atoms are replaced
with aluminium - the aluminosilicates. Clays and zeolites are materials of this type.
Group XIV
As the group is descended from carbon to lead metallic character increases, the +IV state decreases
in stability while the +II state increases in stability. As in other groups there is a large jump between
characteristics of the first row element and the heavier elements of the group.
The presence of d-orbitals allows higher coordination numbers to be achieved through hybridisation
with s and p-orbitals.
sp3d sp
3d
2
This effects the ground state structures and reactivity’s of analogous compounds. D-orbitals may
also become involved in π-bonding leading to a change in geometry and thermodynamic changes.
N(CH3)3 will act as a strong base, a pyramidal geometry allows the lone pair on the nitrogen the
space to become involved in reactions. In N(SiH3)3 the geometry is planar due to overlap between
the 2pz and 3dxz orbitals (on the nitrogen and silicon respectively),leaving the lone pair less available
to act as a base, thus making it a much weaker base. Similar can be said of siloxanes. Also, upon
comparing CCl4 and SiCl4 it is found that the silicon compound is more readily reactive, this is also
due to d-orbitals but this time because they offer a site through which a reaction can begin with the
formation of bonds and then commence bond-breaking, making energy differences usually less
favourable much more achievable.
Trend in Bond Energies
For σ-bonds formed by group XIV elements the strength will decrease as the atomic number
increases, upon descending the group.
In some cases however this trend isn’t followed (i.e. C-F cf. Si-F, C-O cf. Si-O, N-O cf. P-O), these
exceptions are the result of two phenomena, interatomic lone-pair/lone-pair repulsion and pπ-dπ
bonding. When considering Si-F and C-F this is due to the fluorine being able to donate electrons into
the empty 3d orbitals of silicon, changing the bond length. This effect is most pronounced for the 2nd
row as 4d and 5d orbitals are more diffuse than 3d orbitals.
C=Si, Si=Si, Si=O and Si=N bonds are extremely rare. This is due to 2nd row elements not forming
strong pπ-pπ bonds due to very inefficient overlap of p-orbitals. The end result is that there are
almost no silicon analogues of alkenes, alkynes, carbonyls or aromatics, there are very few
compounds known to contain Si=Si bonds.
These can only be prepared when stabilisation is achieved ether by the use of bulky protecting
groups which make σ bonding stearically unfavourable protecting the bond, kinetically, from attack,
or by low temperature isolation in an inert gas matrix.
Group XV
Nitrogen
Nitrogen can exhibit a wide range of oxidation states, +5 → -3.
Ox. State Examples
+5 HNO3 NO3- NO2
+
+4 N2O4 NO2.
+3 N2O3 HNO2 NO+
+2 NO.
+1 N2O HON=NOH
0 N2
-1 NH2OH NH3OH+ NH2O-
-2 N2H4 N2H5+
-3 NH3 NH4+ NH2
-
+5
N2O5 may be considered to be the anhydride of nitric acid. Solid N2O5 is found as [NO2]+ and [NO3]- in
a rock-salt structure, however in the gas phase it is found as a molecular form.
N
O
ON
O
OO
There are two pathways by which it will readily dissociate;
N2O5 2NO2 + ½O2
N2O5 NO3 + NO2 NO + NO2 + O2
It is a powerful oxidising and nitrating agent due to the weakness of the N-O bonds which in turn is
due to the small atom radii and the large inter-electronic repulsion of the lone-pairs present.
The other main +5 compound of nitrogen is HNO3, used commercially to produce nitrates for
fertilisers, explosives and other pyrotechnics.
+4
N2O4 can act as a source of both NO2 and the [NO+] and [NO3-] ions, possibly due to an unstable
isomeric form of N2O4 present. The liquid form of N2O4 has very low conductivity and there is little
dissociation into ions.
In the presence of a solvent with a high dielectric constant or one that can act as a donor the
equilibrium will shift to the right, forming the ions.
There is no evidence of heterolytic cleavage of N2O4 into [NO2+] and [NO2
-].
N2O4 is a powerful oxidising and nitrating reagent, especially for free radical nitration of organic
compounds, it can also react with inorganic species such as NaCl forming NaNO3 and ClNO.
N2O4 can also be used in conjunction with hydrazine derivatives as an oxidising agent, providing a
very compact and efficient fuel.
+3
N2O3 is an intensely blue coloured liquid which will readily decompose to N2O and NO. For this
reason the reactivity of N2O3 is difficult to characterise.
O
N N
OO
N2O3 is the anhydride of HNO2, nitrous acid. This is also unstable and is prepared in situ by
acidification of a nitrite salt.
+2
NO gas is a very reactive free radical species. It has a bond order of 2½, with a bond length of
0.115nm (the intermediate of the N-O bond in NO+ and NO-).
NO is a biologically active gas, acting as a chemical messenger, for example in the dilation of blood
vessels. For this reason NO derivatives can be used to treat for angina.
+1
N2O is manufactured via the careful thermal decomposition of ammonium nitrate melts, it is a
comproportionation process involving nitrogen in both the +V and –III oxidation states.
Unlike the higher oxides, N2O is not the anhydride of an acid form. It will not react with water to
form hyponitric acid (H2N2O2).
N2O is used as a general anaesthetic and is better known as ‘laughing gas’ though is rarely used
anymore as much more preferable options are readily available. It is also used as a propellant and
aerating agent in things like whipped cream, this is due to its high solubility under pressure in
vegetable fats, also as it is non-toxic and tasteless.
+0
N2 is the elemental form of nitrogen. It’s primary uses are all related to its inertness, it is used as a
carrier gas regularly in medicine and industry and can be used as a cryogenic agent.
N2 is most importantly processed into NH3 via either biological fixation or catalysis (the Haber
process).
Recently the reactivity of N2 with transition metals has been studied, with many N2 complexes now
known. They can be synthesised through three main methods;
1. Direct ligand replacement by N2
2. Reduction of a metal salt under N2 gas 3. Conversion of a ligand with N-N bonds to bound N2
Phosphorus
As with the differences between carbon and silicon chemistry, due to d-orbital availability and a
change in electronegativity, there are several differences between the chemistry of nitrogen and
phosphorus.
There is, in fact, a sort of ‘diagonal relationship’ to be found in the periodic table. Phosphorus can
therefore be said to share many of its chemical traits with carbon (the element one above and one
to the left of phosphorus).
Elemental phosphorus displays many allotropes with five crystalline forms having been isolated. All
of these allotropes contain σ-bonds only, 3pπ-3pπ overlap is inefficient and so multiple bonds are
not observed.
The relative stabilities of these allotropes vary with temperature and pressure and so upon
adjustment the desired form can be obtained. Polymeric forms (‘black’ phosphorus and ‘red’
phosphorus) are formed in this manner, other forms such as ‘violet’ or ‘grey’ phosphorus can be
formed at high temperatures in the presence of metals. The simplest form, P4 or ‘white’ phosphorus,
are present in melts.
Oxides of Phosphorus
5O2+P4→P4O10
The stoichometry of the oxide formed will depend upon the conditions applied to the reaction, the
series P4O6+n is known for n=1,2,3,4.
P P
P
P
PO
P
OO
P
P
OO
P P
OO
P
O
P
OO
O
P P
OO
P
O
P
OO
O OP P
OO
P
O
P
OO
O
OOP
OP
OO
P
P
OO
O
O
O
O
The lower oxides will burn in air to eventually yield P4O10 and will condense out of the vapour phase
as the hexagonal form (containing tetrahedral molecules) as above.
Within the P4O10 molecule the PO4 unit is clearly evident: this forms the principal building block of
phosphate chemistry. Thermal modification of P4O10 results in sheets of interlocking PO4 units fused
into heterocyclic rings (cf. silicates).
The principle uses of P4O10 are in the production of phosphoric acid and phosphate esters.
P4O10 + 6H2O 4H3PO4
P4O10 + 6Et2O 4PO(OEt)3
Phosphorus will form oxo-acids with a large range of formulae and oxidation states and is similar to
nitrogen in this manner.
Name Formula Ox
Hypophosphorous acid H3PO2 +I
Orthophosphorous acid H3PO3 +III
Hypophosphoric acid H4P2O6 +IV
Orthophosphoric acid H3PO4 +V
Metaphosphoric acid HPO3 +V
Pyrophosphoric acid H4P2O7 +V
All oxoacids of phosphorus can be classified using a few basic principles:
P is always 4 coordinate - thus H3PO3 is not P(OH)3 even though it is made by hydrolysing PCl3
Acids ending in -OUS are reducing and have P-H bonds
All remaining protons will be on OH groups and this will determine the basicity of the acid.
Tautomerism is possible for H-P=O/P-OH
Condensation can occur to form systems with -P-O-P-O-P-O- repeat units, similar to those seen
in silicates and silicones. The P-O bonds are strengthened by pd bonding. Phosphorus-Nitrogen Halide Compounds Phosphonitrilic compounds are formed when NH4Cl and PCl5 are heated together in the presence of an inert solvent (e.g. C2H2Cl4)
PCl5 + NH4Cl (NPCl2)x + 4HCl
Over several hours these compounds form a buttery mass, the progress of this conversion can be measured by its conductivity.
PCl5 + NH4Cl NPCl2.PCl5 (NPCl2)x
As the reaction progresses the ionic intermediate is consumed. The (NPCl2)x mixture consists of cyclic oligomers with values of x up to 17. Individual compounds are
separated by extraction with benzene (for x = 3, 4), then fractional distillation.
A description of the P-N bonding in these compounds has to take account of the following
observations:
The compounds are thermally and chemically very stable
Skeletal interatomic distances are identical unless the compounds have asymmetric substitution
P-N bonds are shorter than expected for ‘single’ bonds. (0.147-0.158 nm)
N-P-N angles are usually ca. 120o but P-N-P angles vary from 120 - 150o
Skeletal N atoms are weakly basic and can be protonated, especially if the groups on adjacent P atoms are electron releasing
The skeleton is hard to reduce electrochemically unlike aromatics
There is no evidence of bathochromic shifts in UV spectra which are associated with delocalisation changes
There is clearly some double bond character to the P-N bond. This presumably arises from a pπ-dπ
interaction. It does not appear however, that electrons (formally the N lone pairs) are truly
delocalised in these compounds.
This type of bonding is sometimes described as ‘pseudoaromatic’.
N
P NP
PN P
N
N
P NP
P
N PN
Chair Boat
N
P
N
P
N
P
=Cl
Although one structure is puckered the bonding is almost identical in terms of ‘delocalisation’.