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’.