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;M(g) M(g) +
B Al Ga In Tl
798 578 579 556 590
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
the delta positive/negative charges in the bond, in a C-H bond
the C is - however in Si-H it is the H that is -, the Si is +).
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. = 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; a b = 0.017 D A B 1 (D
A A + D(B B) 2
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 Slaters rules
while the surface area of an atom can be taken as the covalent
radius squared, expressed in Angstroms. = 0.704 + 0.359 x Z 2
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 Fajans
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 (BTl, CPb). 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 Tl I 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 s 2 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+ 1/2X2MX
M(s)M(g) 1/2X2(g)X(g) M(g)+X(g)MX(g) Hatom 1/2Hdiss Hbond
If Hbond (the exothermic term) is greater than the sum ofHatom
and 1/2Hdiss (the endothermic terms) then the reaction will
proceed. And for MX3 M+ 1 X2MX3 M(s)M(g) M(g)M*(g) 1 X2(g)X(g)
M(g)+3X(g)MX3(g) Hatom Hprom 1 Hdiss 3Hbond
Introduction of an additional endothermic term (the promotion of
an electron to an excited state) influences the reaction. M(g)M*(g)
s2px1s1px1py1s px py pz s px py pz
If 3Hbond 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
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 p-p bonds (i.e.C=C is much stronger than
Si=Si, also C=O compared to the practically unheard of Si=O bond).
For 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 Overlap orbitals. Integral For 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. Distance such form much
stronger 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 P-O As-O 163 368 331Big step then
steady decrease
Since the first row elements are smaller than second row
elements they can form much more efficient overlaps and as
C-F Si-F Ge-F
485 582 473
Same trend observed
O-H S-H Se-H
467 347 246
Doesnt follow trend
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 G of of
+149.7kJmol-1 and being a light molecule a large amount of energy
is released per gram.
3)Unable to Perform Hypervalency The first row elements have a
maximum coordination number of 4, for the second row elements the
coordination number isnt 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 (sp 3 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 -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 -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+RFR+FBF3The 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 -bonding. In BF3 the result of
this -bonding is the strongest single bond known (rB+F 0.13nm
observed, 0.15nm predicted).
F B F F
This overlap leads to a bonding molecular orbital of 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 Al F F F Al F F F F Al F F F F Al F F F Al F F
As the halides are descended this behaviour changes, AlCl 3 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 4coordinate dimers.
Cl M Cl
Cl M Cl
Cl Cl
The dimeric form is a common structure for group XIII halides.
The bridge bonding in a dimeric trihalide 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 Tl III is not the state found in TlI3 as it may be expected as
the iodine would instantly be oxidised to form I 2 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, Tl I 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 graphites 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
C(graphite) + F2F F F F
(CF)x
F F F
F F
F
This reaction is problematic when considering the synthesis of
pure F 2 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)6CnF2n+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 2p z 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.
C(graphite) + K(g)
C8K(Also C24K, C36K, C48K, C60K)
K K K0.54nm
K
K
K
K KThe bonding involves transfer of electrons from K to vacant
M.Os 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, (CDiamondCGraphite 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 Mohs 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 1900s) 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)B B N B B N N N B B N N N N N B N B B N N N B
B N N N B B B N N N B B N N B B B N N B B B N N N N B N B N B B N N
B B N B N B N B N B N B N B B N B N
N
BB
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+BNNH4+BF4F2+BNBF3+N2 The second possible
structure is of cubic BN (similar to diamond)
N N B N N B N B N B
B
N B N N B B B N B N
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.
H H N B B N B HThe 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. Cl H B HN HB N H NH BH H B N B
N Cl H H N B H H H H H B N B Cl N H N B Cl R H B N B N B
H
H3B NH3
200oC
B3N3H6N H H
Cl H R B N B N H N B R H
N
HCl
H H
LiR
Cl
sp
2
sp
3
Boron-Oxygen Compounds Boron shows a strong affinity for oxygen,
most likely due to -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 B B O O O O B O O B
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-
SiO32-
Si4O116-
Si2O52-
These are the structures that give rise to fibrous materials
such as asbestos.
Sheet;
Si2O52These 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
sp3d2
This effects the ground state structures and reactivitys of
analogous compounds. D-orbitals may also become involved in
-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 -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 isnt 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 p-d 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 2 nd 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 p-p 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 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 +5 +4 +3 +2 +1 0 -1 -2 -3 Examples HNO3
N2O4 N2O3 NO. N2O N2 NH2OH N2H4 NH3 NH3OH+ N2H5+ NH4+
NH2NH2OHON=NOH NO3NO2. HNO2 NO+ NO2+
+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.
O N O N
O
OThere are two pathways by which it will readily dissociate;
N2O5 2NO2 + O2
O
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 N 2O4 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 NO
N O
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 2, 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.
Its 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 N 2 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 -bonds only, 3p-3p 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+P4P4O10 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.O P P P P O P O P P O P O O O P O O O O P P O P O O O
O P P O O O O O O P O O P O O P P P O P O O O O P P O P P 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 Hypophosphorous acid Formula Ox
H3PO2 +I +III +IV +V +V +V
Orthophosphorous acid H3PO3 Hypophosphoric acid Orthophosphoric
acid Metaphosphoric acid Pyrophosphoric acid H4P2O6 H3PO4 HPO3
H4P2O7
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 p-d 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 N P P N
N P N P P N N P P
N P N P N N
P
Chair
Boat
=Cl
Although one structure is puckered the bonding is almost
identical in terms of delocalisation.