Download Topic 4 Chemistry of the Elements of the Main Group

Topic 4
Chemistry of the Elements of the
Main Group
Content
4.1. INTRODUCTION ..................................................................................................................... 83
4.1.1 Natural abundance ......................................................................................................... 83
4.1.2 Classification of Main Group elements. .......................................................................... 84
4.1.3. Physical forms of the Main Group Elements .................................................................. 84
4.1.4. Properties of the Main Group compounds ...................................................................... 87
4.3 CHEMISTRY OF HYDROGEN H ................................................................................................ 89
4.3.1. Properties...................................................................................................................... 89
4.3.2. Compounds of Hydrogens .............................................................................................. 90
4.3.3. Hydrogenation reactions ............................................................................................... 92
4.3.4. Production of Hydrogen ................................................................................................ 93
4.3.5. Industrial uses of Hydrogen ........................................................................................... 94
4.3.6. Reactions of Main Group Hydrides ................................................................................ 95
4.3.7 . Hydrogen Bonding........................................................................................................ 96
4.4. CHEMISTRY OF THE HALOGENS (GROUP 17) .......................................................................... 96
4.4.1. Properties...................................................................................................................... 96
4.4.2. Production of Halogens ................................................................................................. 97
4.4.3. Industrial use of Halogens ............................................................................................. 97
4.4.4. Common stoichiometries of the Halides of period 2 and 3 .............................................. 99
4.4.5. Hydrogen halides ........................................................................................................ 101
4.4.6. Halogen oxyacids ........................................................................................................ 101
4.4.7. Interhalogen Fluorides ................................................................................................ 102
4.5. CHEMISTRY OF THE CHALCOGENS (GROUP 16: O, S, SE) ...................................................... 103
4.5.1. Production .................................................................................................................. 103
4.5.2. Stoichiometry of most common oxides of period 2 and 3 .............................................. 103
4.5.3. Reactivity of Alkali metals with O2, H2O, ..................................................................... 103
4.5.4. Oxides of group 14 (C, Si) ........................................................................................... 104
4.5.5. Oxides of group 15 (N, P, As) ...................................................................................... 106
4.5.6. Oxides of Sulfur ........................................................................................................... 108
4.6. ENVIRONMENTAL IMPACT OF MAIN GROUP COMPOUNDS ................................................... 109
4.6.1. Acid rain ..................................................................................................................... 109
4.6.2. Ozone depletion........................................................................................................... 110
4.6.3. Greenhouse Effect. ...................................................................................................... 112
Prerequisite knowledge
Covalent bonding; octet rule; VSEPR; electronegativity; hypervalency; hybridization.
Learning objectives
Elemental abundance and physical form ; definition of main group elements ; valence orbital
availability ; properties of main group compounds.
Descriptive chemistry of hydrogen, the halogens, oxygen and sulfur.
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
4.1. Introduction
The main group elements include the s-block elements (group 1 and 2) and the p-block
elements (group 13 to 18).
Electronegativity Table of the Main Group Elements
H
He
2.1
G1
G2
G13
G14
G15
G16
G17
G18
Li
Be
B
C
N
O
F
Ne
1.0
1.5
2.0
2.5
3.0
3.5
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.2
1.5
1.8
2.1
2.5
3.0
K
Ca
Ga
Ge
As
Se
Br
0.8
1.0
1.6
1.8
2.0
2.4
2.8
Rb
Sr
In
Sn
Sb
Te
I
0.8
1.0
1.7
1.8
1.9
2.1
2.5
Cs
Ba
0.7
0.9
Metals
<2
Metalloids
~2
Non-Metals
>2
Ar
Kr
Xe
These elements are the most abundant elements in the solar system and in the universe and
are also the major components of the human body.
4.1.1 Natural abundance
The six most abundant elements in the universe: H (for 91%), He, O, C, N, Ne, are all
main group elements.
The mass of earth is made principally of Fe (32%), O (30%), Si (15%), Mg (14%), S (3%),
Ni (1.8%), Ca (1.5%), Al (1.4%). All except for Fe and Ni, are main group elements.
The human body is composed principally of: O (65%), C (18%), H (10%), N (3%), Ca
(1.5%), P (1.2%), K (0.2%), S(0.2%), Cl (0.2%), Na (0.1%), Mg (0.05%), all these are main
group elements. The transition metal element: Fe, Co, Cu, Zn account for less than 0.05%
(and are called trace elements).
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4.1.2 Classification of Main Group elements.
The main group elements are classified according to their conduction properties into three
groups: Metals, Metalloids and non-Metals
Non-metals are not electrical conductors, they are characterised by a high electronegativity
value (over 2). The valence electrons of non-metals are strongly attracted to their positively
charged nucleus and are not available to conduct electricity.
Metals are electrical conductors; they have a low electronegativity value (under 2). The
valence electrons of metals are more loosely attracted to the nucleus and are free to conduct
electricity. Metals make crystal lattice structures in which electrons can flow freely.
Metalloids or semi-metals show intermediate conduction properties (they are semiconductors). Their electronegativity values are close to 2. The valence electrons of
metalloids are localised around the nucleus but can also be moved (and therefore conduct
electricity) when excited by small inputs of energy. Semi-conductors elements such as Si,
As, Sb play an important part in microelectronic devices.
4.1.3. Physical forms of the Main Group Elements
S-BLOCK ELEMENTS, groups 1 and 2 (H excepted) are typically metallic and form lattice
structures.
Lattice structure. The most stable form for lithium metal at 298K is the body-centered cubic
(bcc) structure. Under standards conditions, all of the group 1 (alkali metals) elements have
a bcc structure (bcc is made of two interlocked cubic arrays of atoms).
304 pm
electron
lithium ion in lattice
every Li is surrounded by eight other Li
The elements of the first two groups (such as alkali metals: Li, Na, K, and alkaline earth
metals: Be, Mg, Ca) are silver-coloured, soft, low-density metals. They have respectively one
and two valence electrons which are easily lost. They form lattice structures in which the
valence electrons are free to conduct electricity. The s-block elements are characterised by
low ionisation energies and hence are very reactive.
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P-BLOCK ELEMENTS
Monoatomic structure (noble gases). The noble gases under standard conditions are all
odourless, colourless, monoatomic gases with very low chemical reactivity.
Diatomic structures can be gas, liquids or solids, depending on temperature. For example
hydrogen (H2), nitrogen (N2), oxygen (O2), chlorine (Cl2).
H
H
N
N
O
O
Cl
Cl
Lewis structures of diatomic molecules: H2, N2, O2, Cl2.
Polyatomic structures
B, P, S, C have several molecular allotropes; discrete polyatomic molecules such as B 12, P4,
S8, C60.
Boron has many molecular allotropes such as B 12, B28, B50, B105-108, B192.
Alfa-Boron has a unit cell of twelve B atoms: B 12. The structure is icosahedral. Each boron
(B: [He] 2s2 2p1) has three valence electrons, and the icosahedral structure is achieved by
three centers two electrons bridging. Each bridge is made of one electron pair from a B-B
bond and one empty orbital: a pair of electron is shared between three borons: this makes a
3 centre 2-electron bond. In B12 structure, each boron atom is bonded to five others.
empty orbital
B
B
B
B
B
B
B
B
bp B
B
B
B
B
B
B
B
3 centers 2 electrons bond
B12 icosahedron.
Phosphorus can exist in several molecular allotropes such as P 2, P4.
White phosphorus, tetraphosphorus P 4 is a molecule made up of four P atoms and has a
tetrahedral structure. Each phosphorus (P: [Ne] 3s 2 3p3) has a sp3 hybridisation. In the P4
structure, the four lone pairs are pointing outside the structure and six sigma bonds link the
four phosphorus atoms together.
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lp
P
P
P
P
tetraphosphorus molecule
Carbon is known for its two main allotropes: diamond and graphite and also for forming
carbon nanotubes and many other molecular allotropes such as fullerenes C 20, C60, C70.
In Fullerene C60, each carbon has sp2 hybridisation and is bonded to three neighboring
carbons. The ball shape of fullerene is achieved by a series of six and five carbon rings (just
like a football).
Fullerene, C60 is made of 60 sp2 carbons organized in series of 5 and 6 member rings.
Extended lattice structure
G14
G15
G16
Si
P
S
1.8
2.1
2.5
Ge
As
Se
1.8
2.0
2.4
Sn
Sb
Te
1.8
1.9
2.1
C
2.5
Carbon exists in various structural forms and like some other main group elements (Si, P, S,
Ge, As, Se, Sn, Sb, Te, Bi) can make lattice structures (infinite networks) such as diamond
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(3D structures of sp3 carbons made of sigma-bonds), graphite and nanotubes (2D structures
of sp2 carbons involving sigma and -bondings).
2
sp
3
C
sp
C
Diamond structure
Graphite, Fullerene, nanotube structure
Note. Graphite is made of layers of extended lattices (principally flat molecular sheets of sp 2
carbon made of six carbon rings).
Tin has few allotropes: andtin.
Alfa-tin (-tin) or grey tin is a non metallic allotrope; it is brittle. It has a diamond cubic crystal
structure with covalently bonded atoms: the valence electrons in alfa-tin are not available for
conduction and the structure is brittle and non metallic.
Beta-tin (-tin) or white tin is a metallic allotrope: it had a lattice structure which is stable and
malleable.
4.1.4. Properties of the Main Group compounds
Bonding polarity (ionic, polar covalent, covalent)
The chemical bonds formed by the elements of the main group cover a very broad range of
polarities. Depending on the difference in electronegativity between bonded atoms, the main
group elements can form ionic bonding (when the difference in  is over 2), polar covalent
bonding (the difference in  is more than 0.5) and covalent bonding (the difference in  is
between 0.5 and 0).
Table: Chemical bond polarity: (rule of thumb): bonds between S-Blocks elements,
Metaloids, Non-metals)
s-Block metals
Metalloids (and H)
Non-Metals
Non-Metals
Ionic
Polar Covalent
Covalent
Metalloids (and H)
Polar covalent
Covalent/Metallic
Polar Covalent
S-block metals
Metallic bond
Polar covalent
Ionic
The ionic bond is considered achieved when the difference of electronegativity of the atoms
is greater than 2.
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The s-block elements have low electronegativity values ( is less than 1.5) and form polar
bonding with the metalloids (of which  is close to 2.0) and form ionic bonding with the nonmetals (2.0 to 4.0).
Example. Caesium fluoride, CsF (= 4.0 – 0.7 = 3.3); sodium chloride, NaCl ( = 3.0 –
0.9 = 2.1) are regarded as having typical ionic bonding.
The s-block elements form ionic hydrides with hydrogen. Formally ionic hydrides contain a
metal cation and hydride ion H-, (the s-block element give its electron to a hydrogen which
get reduced into H- a hydride of oxidation number -1). Ionic hydrides are extremely reactive
towards water, and are moisture sensitive (they can be used as drying agents).
Example: NaH ( = 2.1- 0.9 = 1.2)
NaH + H 2O
H2 + NaOH H = -83.6 kJ/mol
Acid/base properties (Brønsted-Lowry and Lewis acidity)
Brønsted-Lowry definition of acids/bases is based on proton exchange (acids are proton
donors, base are proton acceptors). The strength of an acid or a base depends on its
dissociation in water.
Strong acids dissociate totally in water giving H3O+.
Example: Sulfuric acid (H2SO4).
H2SO4 + H2O
HSO 4- + H3O+
Strong bases dissociate totally in water giving OH-.
Example: Sodium hydroxide (NaOH) dissociates totally in water giving OH- and Na+
NaOH
Na+ + OH-
Weak acids dissociate only partially in water, and are in equilibrium with H 3O+.
Example: Hydrofluoric acid, HF, is the only hydrohalic acid that is not considered a strong
acid. HF is a highly corrosive acid, capable of dissolving many materials including oxides
and glass but it is a weak acid because of its lower dissociation constant in water:
HF + H 2O
F- + H3O+
pKa = 3.17
Weak bases dissociate only partially in water, and are in equilibrium with OHExample: Ammonia NH3
NH 3 + H2O
NH 4+ + OH-
Lewis acid/base concept is based on electron pair (ep) donor/acceptor properties: acids are
ep acceptors, bases are ep donors:
Acids: electron deficient species are acceptors of electron pairs: BF3, H+.
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Bases: electron rich species – with at least one lone pair available – are donors of
electrons pairs: NH3, H2O.
Redox properties
Oxidising agent (reduced during the redox process: they gain electrons). A strong oxidant
increases its stability by gaining extra electron(s).
E.g. Permanganate ion MnO4-:
-
+
-
MnO 4 + 8H + 5 e
Mn
VII
2+
+ 4 H2O
II
Bromate ion: BrO3-
BrO3- + 6H+ + 6e-
Br- + 3 H2O
V
-I
Fluorine: F2
-
-
F2 + 2e
2F
0
-I
Reducing agents (oxidised during the redox process: they lose electrons). A strong reducer
increases its stability by losing electron(s).
Sodium metal: Na.
Na + + e-
Na
0
I
Hydride H-.
-
H
-I
1
-
/2 H2 + e
0
4.3 Chemistry of Hydrogen H
4.3.1. Properties
Hydrogen is the most abundant element in the universe. Most of the hydrogen on earth
occurs bonded to oxygen (H2O) or to carbon (hydrocarbons and other organic molecules).
Electronic configuration: 1s1
Bond energy. The bond energy is the amount of energy required to break one mole of
molecules into their individual atoms. The H-H bond is one of the strongest single bonds:
only H-F (567kJ/mol) and O-H (463 kJ/mol) are stronger than H-H. The molecule of H2 is
thus a relatively stable molecule.
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Bond energy
H2
E
2H
H = 436 kJ/mol
Ionisation energy. The amount of energy required to remove an electron from the atom in
its gaseous state. H ionisation energy is very high; in fact similar to O and Cl IE :1313
and1251 kJ/mol respectively. Therefore, H + is difficult to form and has in fact no separate
chemical existence: it is always in a combined form and will migrate from one atom to
another. For example it forms H3O+ with water.
Ionisation energy
+
H
-
H +e
IE
H = 1309 kJ/mol
Electron affinity. The amount of energy released when an electron is added to a neutral
atom to form a negative ion.
Electron affinity
H+ e
-
-
H
EA H = -71 kJ/mol
Hydrogen forms ionic hydrides with the reactive s-block metals (groups 1 and 2) and forms
covalent hydrides with the p-group metals, e.g. Al and Sn (group 13 and 14).
Electronegativity  = 2.1. The value is intermediate in the electronegativity scale that spans
from 0.7 to 4.0.
H can form hydrides (negatively charged H) with a whole range of polarity, from ionic
hydrides Cs-H ( = 1.4) to non-polar covalent hydrides B-H ( = 0.1).
H can form hydrogen compounds (H is positively charged) with a whole range of polarity,
from ionic hydrogen halides HF ( = 1.9) to non-polar covalent C-H ( = 0.4).
4.3.2. Compounds of Hydrogen
4.3.2.1. Ionic hydrides (compounds of H-) form with very reactive s-block metals (NaH,
CaH2 from Na and Ca), the metal atoms increase their stability by giving all their valence
electrons to the hydrogen forming a negatively charged hydride ion H-. Hydrides ions are
strong reducing agent (they readily give away their electron).
4.3.2.2. Covalent hydrides (compounds of H-)
The covalent hydrides form infinite lattice solids with p-group metals
Example: with aluminum, AlH3 ( = 0.59)
They form discrete molecular species (which could be solids, liquids or gases) with the nonmetals.
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Example: Phosphine is a odourless, colourless, flammable and toxic gas.
PH3 ( = 0.01)
H
H
P
H
phosphine
4.3.2.3. The stability of hydrides decreases down any group of the periodic table.
PH3 and AsH3 are stable but decompose on warming: SbH3 and BiH3 are unstable at room
temperature. As the differences in electronegativity increase, the strength of the M-H bond
decreases due to poorer orbital overlap.
Phosphine, PH3 ( = 0.01) is a flammable but stable gas at room temperature, it
decomposes into P4 gas at 650oC.
Arsine, AsH3 ( = 0.02) is stable kinetically but not thermodynamically. It decomposes to
produce arsenic at 230oC but is stable at room temperature.
Stibine, SbH3 ( = 0.15) decomposes slowly at room temperature to form antimony metal.
Bismuthine, BiH3 (= 0.18) is unstable and decomposes to produce bismuth metal well
below 0oC.
Stoichiometry of hydrogen compounds of periods 2 and 3.
Period 2
G1
G2
G13
G14
G15
G16
G17
LiH
BeH2
BH3
CH4
NH3
H2O
HF
SiH4
PH3
H2S
HCl
B2H6
Period 3
NaH
MgH2
Ionic
hydrides
AlH3
Polymeric
Covalent
Solid
Discrete Covalent
Molecules
Hydrogen compounds (except with group 13) follow the stoichiometry expected from the
octet rule.

With group 1 and 2: both the metal and hydrogen achieve noble gas structures by
forming ionic hydrides.

With group 14 to 17: both atoms achieve noble gas structures by forming covalent
molecules.

With group 13: there is an anomaly due to the lack of valence electrons to satisfy the
octet rule and the valence of hydrogen.
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With boron, the stoichiometry is 3 but borane (BH 3) forms gaseous dimmers: diborane (B2H6)
with bridging H atoms (3 centre 2-electron bonds).
bridging H
empty orbital
H
H
H
B
+
H
H
H
H
B
B
H
H
H
H
H
H
B
B
H
131 pm
o
97
o
120
B
H
o
119 pm H
H
83
H
dimer B2H6
With aluminium, the stoichiometry is 3, aluminium hydride (or alane) is a polymeric covalent
solid of stoichiometry: AlH3. In alane, each aluminium atom is surrounded by six hydrogen
atoms that bridge to six other aluminium atoms
4.3.2.4. Interstitial hydrides
Interstitial hydrides exist within metals or alloys, they are usually non-stoichiometric, with
variable amounts of hydrogen atoms in the lattice.
Example: Palladium absorbs up to 900 times its own volume of hydrogen at room
temperature, forming palladium hydride PdHx (potential application in H2 storage).
4.3.3. Hydrogenation Reactions
Hydrogenation is a chemical reaction between molecular hydrogen and other compounds.
Hydrogen is used as a reducing agent: hydrogenation generally occurs in the presence of a
catalyst. The process is employed to reduce or saturate organic compounds (ketone,
alkynes) into alcohols and alkanes respectively.
Hydrogen is an oxidizing agent when reacting with metals to form hydrides.
HX
hydrogen halides
H
C
H
X2
N2 cat
NH 3
ammonia
high P
H2
, cat
H
C
H
alkane
O
, cat
M

MxOy

H
C
MHx
OH
hydrides
alcohol
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4.3.4. Production of Hydrogen
About 50 million tons of hydrogen is produced every year, more than 90% of which is made
from fossil fuels (by steam reforming, partial oxidation of oil, coal gasification) and about 4%
is by electrolysis of water.

Steam reforming of hydrocarbons is a common method of producing bulk
hydrogen. It is performed at high temperature (600 to 1000 oC) in the presence of a
metal based catalyst: steam reacts with methane to give carbon monoxide and
hydrogen.
600 - 1000oC
CH4 + H2O
CO + 3 H2
Cat (Ni/Al2O3)
Additional hydrogen can be recovered by a lower temperature gas-shift reaction with
the carbon monoxide produced.
1) 350oC
o
2) 190-210 C
CO + H2O

Partial oxidation of heavy oil occurs when a sub-stoichiometric fuel-air mixture is
partially combusted in a reformer, creating a hydrogen-rich syngas (mixture of
hydrogen and carbon monoxide). Syngas can be used as fuel for internal combustion
engines. Alternatively, more hydrogen can be recovered by the lower temperature
gas-shift reaction.
2 CnH2n+2 (l) + n O2 (g)
2(n+1) H2 (g) + 2n CO(g)
CO (g) + H2O (g)

H2 (g) + CO2 (g)
Coal gasification is the process of producing syngas from coal and water with air
and oxygen.
3 C(s) + O2 (g) +

CO2 + H2
H2O (l)
H2 (g) + 3 CO (g)
Electrolysis of water is the decomposition of water into oxygen and hydrogen gas
resulting from an electric current.
H2O
H2 (g) + 1/2 O2 (g)
This is an expensive process but the majority of the hydrogen produced through
electrolysis is in fact a side product in the production of chlorine.
2 NaCl + 2 H2O

Cl2 + H2 + 2 NaOH
Biological production, using bioreactors with green algae in water, bacteria and
archaea growing on biomass. These processes are still being developed.
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Green algae, such as C. reinhardtii can switch from oxygen production (as part of the
photosynthetic metabolism) to hydrogen production when growing in absence of sulfur.
Some bacteria and archaea produce hydrogen gas as part of their normal metabolism;
they can produce hydrogen from biomass.
4.3.5. Industrial uses of Hydrogen
Ammonia synthesis. The largest application of H2 is in the production of ammonia; the
hydrogen reacts with nitrogen in presence of a catalyst to form anhydrous liquid ammonia.
3 H2 + N2
catalyst
2 NH3
Oil refineries. Large quantities of H2 are needed in the petroleum industry, for the upgrading
of fossil fuels.
Hydrodealkylation (example: hydrodealkylation of toluene into benzene)
CH3
+ H2
40-60 atm
+ CH4
o
500-600 C
benzene
toluene
Hydrodesulfurisation (example: hydrodesulfurisation of ethanethiol into ethane and
hydrogen sulfide)
C2H5SH + H 2
catalyst
C2H6 + H2S
Hydrocracking is a catalytic cracking process obtained at high partial pressures of
hydrogen. Hydrocraking is used for the removal of sulfur and nitrogen hetero-atoms from
hydrocarbon streams.
MeOH synthesis. Hydrogen is used as a hydrogenating agent to produce methanol from
carbon monoxide. The process takes place at high pressure (50-100 atm) and at 250oC.
CO + 2 H2
50-100 atm
o
CH3OH
250 C
Hydrogenation of vegetable oil into saturated fats. Hydrogen is used in the food
processing industry, as a hydrogenating agent, to increase the level of saturation of
unsaturated vegetable fats and oils.
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4.3.6. Reactions of Main Group Hydrides
The chemistry of hydrogen compounds changes with periodic trends (electronegativity and
size) across the periodic table.
Hydrides of group 1 and 2 are strong reducing agents and can react violently with water.
NaH(s) + H2O(l)
Na+(aq) + OH-(aq) + H2(g)
Hydrides of group 13 are mild reducing agents, Diborane reacts with water to form
hydrogen and boric acid (antiseptic and insecticidal).
B2H6(g) + 6 H2O(l)
2 H3BO3(aq) + 6 H2(g)
Hydrogen compounds of group 14 are covalent and their polarity increases down the
group. As the electronegativity decreases down the group (C = 2.5; Si =1.8, Ge =1.8) the
partial charge on the hydrogen goes from positive (with C) to negative (with Si, Ge, Sn) as
the polarity increases.
Example: comparing methane and silane.
Methane (CH4) is a very stable molecule, the C-H bond is covalent with very little polarity
(hydrogen has a small partial (+) charge) and methane reacts with oxygen (combustion of
methane) but does not react with water.
Silane (SiH4) ignites spontaneously in air and reacts instantly with water. The Si-H bond is
more polar than C-H and hydrogen has a partial (-) charge. As the central atom gets larger,
water molecules can access it easily.
SiH4(g) + 2 H2O(l)
SiO2(s) + 4 H2(g)
Hydrogen compounds of group 15. Hydrogen forms covalent bonds with nitrogen (in
which the small polarity gives a slight positive charge to the hydrogen), and forms hydrides
of increasing polarities down the group (from P, to As, and Sb) as the electronegativity
decreases.
Ammonia (NH3) dissolves in water to give weakly basic solutions.
NH3(g) + H2O(l)
NH4+(aq) + OH-(aq)
Phosphine (PH3) is a very weak base with a low solubility in water. It spontaneously ignites
in air.
Hydrogen compounds of group 16. Hydrogen forms polar covalents bonds with the
elements of group 16, the polarity of which decreases down the group (from O, S, Se, to Te).
They do not react with water. For example, hydrogen telluride (H2S) is only weakly acidic
and is a gas above -1oC.
Hydrogen compounds of group 17. Hydrogen forms polar bonds with the elements of
group 17: all are acids in water (HF is a weak acid due to its poor dissociation in water, but
all the others are strong acids.
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
4.3.7 . Hydrogen Bonding
Covalent bonds between hydrogen and N, O, F or Cl are highly polarised resulting in the
hydrogen holding a large (+) partial charge.
Hydrogen bonding is a strong electrostatic interaction that occurs between the positively
charged hydrogens and negatively charged atoms (N, O, F, Cl) within the same molecule
(intramolecular bonding) or between two molecules (intermolecular bonding).
Intramolecular hydrogen bonding is responsible for the shapes and functions of many
macromolecules such as proteins and DNA.
Intermolecular hydrogen bonding is the strongest intermolecular force (20 to 100 kJ.mol-1).
Intermolecular hydrogen bonds are responsible for the 3D structure of ice water and for the
high boiling point of water.
When comparing boiling points down the elements of the group 16:
(H2O (100oC); H2S (-60oC); H2Se (-41.25oC), H2Te (-1oC)), H2O seems abnormally high. This
is because the hydrogen bonds create a cohesion amongst the water molecules in the liquid
form which prevents them from escaping into a gaseous form. This must be overcome to
generate gaseous H2O, hence the higher temperature required.
Boiling point (oC)
100
H2O
mass
0
H2Te
H2Se
-100
H2S
4.4. Chemistry of the Halogens (group 17)
4.4.1. Properties
Halogens have seven valence electrons and a valency of one: they form diatomic halogen
molecules such as F2, Cl2, Br2, etc .
Example: Fluorine (F2) has a very weak F-F bond (BDE is only 159 kJ.mol-1). The small bond
length results in inter-electronic repulsions between lone pairs, weakening the F-F bond.
Halogens have the highest electronegativity values in their respective period going from At
( = 2.2) to F ( = 4.0).
They are strong oxidants. Halogens have a high affinity for electrons (EA) and form stable
halide ions (X-) with hydrogen and metals.
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
4.4.2. Production of Halogens
Fluorite (CaF2) is the main mineral source of commercial fluorine.
Cryolite (Na3AlF6) is the least abundant but the most concentrated mineral source of fluorine.
Fluorine (F2) is produced by electrolysis of a KF/HF mixture. HF alone cannot be
electrolysed, the presence of some KF is critical even though it is not consumed in the cell
and remains in solution. When combined, potassium fluoride (KF) and hydrogen fluoride
(HF) produce potassium bifluoride (KHF2), which increases the conductivity of the solution.
2 HF(l)
H2(g) + F2(g)
Chlorine (Cl2) is produced by electrolysis of NaCl aqueous solution.
-
reaction at anode
-
2Cl (aq)
Cl2(g) + 2e
reaction at cathode 2 H2O + 2e-
2 OH-(aq) + H2(g)
-
-
2 Cl (aq) + 2 H2O
Cl2(g) + 2 OH (aq) + H2(g)
Note. During the process H2 and NaOH are also produced.
2 Na+(aq) + 2 Cl-(aq) + 2 H2O
Cl2(g) + 2 NaOH(aq) + H2(g)
Bromine (Br2) and Iodine (I2) are produced by chlorine oxidation of halide compounds
containing (Br- or I-).
o
-
2 Br (aq) + Cl2(g)
100-120 C
-
Br2(g) + 2 Cl (aq)
Note. It is necessary to heat the system to prevent the bromine from remaining in solution.
4.4.3. Industrial use of Halogens
Fluorine (F2)
Over half of the production of fluorine is used in uranium separation. Uranium is turned into
gaseous uranium hexafluoride to separate uranium-235 (required for nuclear fission) from
the heavier isotope uranium 238.
Polytetrafluoroethylene (PTFE) production. PTFE (DuPont’s Teflon) is an important
substance with applications in electrical insulation, as a dielectric. It is also used in the
chemical process industry in coating pipes, tubing, jointing materials, etc - where corrosion
resistance in needed.
Chlorofluorocarbons (CFCs) used as refrigerant gases, were the major class of fluorinated
organic chemical in the 1980’s. Hydrochlorofluorocarbons and hydrofluorocarbons (HFCs)
have now replaced CFC as refrigerants, following concerns about their environmental impact
(see chapter 4.6).
Several million tons per years of sodium fluoride (NaF) is produced globally. NaF has many
applications: it is used in dentistry, in water treatment and as a cleaning agent.
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
Chlorine Cl2
The production of industrial and consumer products, such as plastics, solvents, textiles,
insecticides, are the principal applications of chlorine (70% of production).
Organochlorine intermediates are used in the production of many important industrial
materials such as polycarbonates and polyurethanes.
Sanitation and water treatment uses about 20%, in bleaches and disinfection products
preparation.
The production of inorganic material (such as hydrogen chloride (HCl), dichlorine monoxide
(Cl2O), hypochlorous acid (HOCl), etc) uses about 10%.
Bromine Br2
Brominated flame retardants (BFRs) represent the largest commercial use of bromine (27%).
BFRs are organobromide compounds that have inhibitory effects on the ignition of
combustible organic materials. They are effective in plastics and textile applications.
Ethylene di-bromide (BrCH2-CH2Br) was historically used as an anti-knock fuel additive. It
reacts with lead residues to generate volatile lead bromides, preventing fouling of engines.
About 15% of bromine is used for the preparation of agricultural chemicals. For example,
methyl bromide or ethylene bromide, both poisonous, have been used as pesticide to treat
soils and houses.
Iodine I2
The largest part of iodine production is consumed by the production of ethylenediammonium
diiodide (EDDI) a nutritional supplement for livestock.
Iodine is also consumed in the synthesis of acetic acid in which iodine is a co-catalyst.
About half of the iodine produced is used in the production of a range of organic chemicals.
Potassium iodide (KI) is produced for medical application and consumes about 15% of the
iodine production. KI has been used as an expectorant and is currently used to treat acute
thyrotoxicosis.
15% used directly as I2 ; 15% other Iodide salts.
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
4.4.4. Common stoichiometries of the Halides of period 2 and 3
Stoichiometries of the fluorides
G1
G2
G13
G14
G15
G16
G17
LiF
BeF2
BF3
CF4
NF3
OF2
F2
NaF
MgF2
AlF3
SiF4
PF3
SF2
ClF
PF5
SF4
ClF3
SF6
ClF5
G18
Stoichiometries of the chlorides
G1
G2
G13
G14
G15
G16
G17
LiCl
BeCl2
BCl3
CCl4
NCl3
OCl2
ClF
AlCl3 SiCl4
PCl3
S2Cl2
Cl2
PCl5
SCl2
NaCl MgCl2




G18
The stoichiometries of halides are mostly similar to that of the hydrides.
The chemical reactivity of halides varies dramatically along the periodic table (NaF
and OF2). (Do not mistake OF2 for NaF! NaF is used in toothpaste and OF2 is a
powerful fluorinating agent).
The halides of group 1 and 2 are ionic halides, when the halides of other groups are
polar covalents.
The elements of the second period strictly obey the octet rule but the elements from
the third period (and higher) have access to extra orbitals to expand their electronic
configuration, resulting in possible hypervalency.
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
Example: PF3 (obeys the octet rule) and PF5 (hypervalent).
upe
F
[He] 2s2 2p5
p
s
Ground state P
2
3
d
0
[Ne] 3s 3p 3d
sp3
Hybridisation sp 3
P
F
F
F
Phosphorus in PF3 obeys the octet rule: P sp 3 hybridisation provides four valence orbitals
(three bonding to F and a lone pair). The total number of electron around the P is 8.
Because an empty 3d-orbital is available, the phosphorus can also undergo excitation to
promote one electron from the 3s-orbital to one empty 3d-orbital, providing five unpaired
electrons and allowing the formation of a PF 5 molecule (here the P is hypervalent: it is
surrounded by 10 valence electrons.
p
s
Excited state P
d
[Ne] 3s1 3p3 3d1
sp3d
Hybridisation sp 3d of P
10 ve arround P
F
F
P
F
F
F
PF5 does not obey the octet rule: phosphorus sp 3d hybridisation provides five valence
orbitals for bonding with five F. As a result 10 valence electrons surround the central
phosphorus.
Hypervalency is possibe for the elements of period 3 and over and is common for the
elements of group 15, 16 and 17 (and also 18: for example the fluorides of noble gases such
as XeF4)
More on hypervalency (see previous topic: Chapter 3.2.2)
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4.4.5. Hydrogen halides
Hydrogen halides are strong acids (except H-F which is a weak acid as it does not dissociate
well in water).
H-F and H-Cl are major industrial chemicals.
Hydrofluoric acid (H-F) is produced by the treatment of the mineral fluorite (CaF2) with
concentrated sulfuric acid:
CaF2(s) + H2SO4(l)
CaSO4(s) + 2HF(g)
H-F solution reacts with SiO2, (H-F attacks glass, used in etching):
6 H-F (aq) + SiO2 (s)
2 H+ (aq) + SiF62- (aq) + 2 H2O (l)
H-F is used as a catalyst in oil refining and is an intermediate in the synthesis of many
organofluorine compounds.
Hydrogen chloride H-Cl is prepared by direct reaction of H2 and Cl2 gases.
UV light
Cl2(g) + H2(g)
2 HCl(g)
H-Cl can also be prepared in the laboratory (in small amounts) from the reaction of NaCl with
H2SO4.
NaCl + H2SO4
NaHSO4 + HCl
4.4.6. Halogen oxyacids
Halogen oxyacids are polyatomic acids that contain oxygen, halogen and at least one
hydrogen atom.
Acid
Name of Acid
Anion
Name of Anion
H-O-Cl
hypochlorous
ClO -
hypochlorite
1+
H-O-ClO
chlorous
ClO 2-
chlorite
3+
H-O-ClO 2
chloric
ClO 3-
chlorate
5+
H-O-ClO 3
perchloric
ClO 4-
perchlorate
7+
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Cl ox. no.
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
Applications: Halogen oxyacids are strong oxidizing agents.
 Calcium hypochlorite [Ca(ClO)2] is used in swimming pools for water treatment.
 Sodium chlorite (NaClO2) is a major component of liquid bleach.
 Sodium chlorate (NaClO3) is used in the paper making industry as a bleaching agent
for pulp.
 Sodium perchlorate (NaClO4) was used as a solid-fuel booster for the US space
shuttle: - each launch used about 700 tons.
4.4.7. Interhalogen Fluorides
Interhalogen fluorides contain fluorine and other halogen atoms (chlorine, bromine, iodine or
astatine).
They are subject to hydrolysis and ionize into polyhalogen ions.
The general formula is XYn, with the most electronegative halogen placed last.
Example: Interhalogen fluorides (XFn).
Oxidation
number of X
0
1
3
5
7
Fn
ClFn
BrFn
IFn
ClF
ClF3
ClF5
BrF
BrF3
BrF5
IF
IF3
IF5
IF7
F2
The geometry of interhalogen fluorides obeys the VSEPR theory.
Formula
Lewis structure
XF
X
F
VSEPR geometry
AXE3, tetrahedral
X
F
F
F
F
XF3
X
F
X
AX 3E2, trigonal bipyramidal
F
F
F
XF5
F
F
F
X
F
AX5E, octahedral
X
F
F
F
XF7
F
F
F
F
F
X
F
F
F
F
AX 7, pentagonal bipyramidal
F
F
F
X
F
F
F
F
Note. In the interhalogen general formula (XYn); X represents the less electronegative
halogen (with fluorides, X = Cl, Br, I). In the VSEPR geometry AXnEm, X represents the
number of F. Not to be confused.
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
4.5. Chemistry of the Chalcogens (group 16: O, S, Se)
The chalcogens are the elements of the oxygen family or group 16. All chalcogens have six
valence electrons and their most common oxidation states are -2, +2, +4 and +6.
The focus of this course is comparison between the elements O and S.
4.5.1. Production
The main method of oxygen (O 2) production is by fractional distillation of liquefied air. About
100 million tons of O2 are produced annually for industrial uses.
Sulfur is either mined or recovered from oil refineries; about 68 million tons are produced
annually.
4.5.2. Stoichiometry of most common oxides of period 2 and 3
G1
G2
G13
G14
G15
G16
G17
Li2O
BeO
B2O3
CO2
N2O
O2
F2O
CO
NO
O3
F2O2
P4O6
SO2
Cl2O
P4O10
SO3
Cl2O2
G18
N2O3
NO2
N2O5
Na2O2
MgO
Al2O3
SiO2
4.5.3. Reactivity of Alkali metals with O 2, H2O,
Alkali metals react with O2 to give oxide (M2O), peroxide (M2O2) or superoxide (MO2).
The relative stability of these species depend on the size of the metal ion (larger cations
stabilize larger anions):
 Small ions (Li) form stable oxides (Li2O)
 Medium size ions (Na) form peroxides (Na2O2)
 Large ions (K, Rb, Cs) form superoxides (KO2, RbO2, CsO2)
The oxides of group 1 and 2 react with water to give basic solutions.
Example with Li2O.
Li2O (s) + H2O (l)
2 Li+ (aq)
+
2 OH- (aq)
The oxides of non-metals (except CO) react with water to give acidic solutions.
Example with SO3.
SO3 (s) + H2O (l)
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2 H+ (aq)
+
SO42- (aq)
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
The oxides of the other metals do not react with water, but react with both acids and
bases: they are amphoteric.
Example of Al2O3.
Al2O3 (s) + 6 H+ (aq)
2 Al3+ (aq) + 3 H2O (l)
Al2O3 (s) + 2 OH- (aq) + 3 H2O (l)
2 Al(OH)4- (aq)
4.5.4. Oxides of group 14 (C, Si)
The elements of group 14 form very strong bonds with oxygen.
The oxides of carbon (CO2 and CO) are both discrete molecules of gases under standard
conditions of temperature and pressure, and are both thermodynamically stable.
Carbon dioxide (CO2) is a byproduct of fossil fuels combustion. Carbon dioxide is also a
useful product: about 30 million ton a year is used industrially: 50% as a refrigerant (dry ice)
and 25% for the preparation of fizzy drinks.
Carbon dioxide is a potent greenhouse gas: it traps solar radiation in the atmosphere. (ref:
see later discussion in chapter 4.6.3).
O
C
O
carbon dioxide
Carbon monoxide (CO) is a highly toxic gas. CO forms when carbon compounds burn in
deficiency of O2. The toxicity is due to the fact that CO molecules bind to the iron centre of
haemoglobin and prevent O2 from binding to it: this affects oxygen transport throughout the
body.
C
O
C
carbon monoxide
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O
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
Silicon dioxide (SiO2). is not a discrete molecule like CO2 and CO, but makes extended
lattice solids. SiO2 can be amorphous or crystalline and has many different crystal
structures. SiO2 is most commonly found in nature as sand or quartz.
Unlike carbon, silicon does not form strong -bonds with oxygen, because the overlap
between the Si 3p and the O 2p is relatively poor. Instead, Si forms four -bonds with four O
atoms and each O is further linked to another Si by -bonds.
O
Si
O
O
O
Si
Si
O
Si O
O
O
O
O
3D structure of silicon dioxide (SiO2)
Zeolites are microporous aluminosilicate minerals, used as commercial adsorbents and
molecular sieves. They are used in water purification which is the single biggest use of
zeolites: to "soften" water (to reduce the concentration of ions in water). Zeolites are also
used in gas separation and as catalysts in nuclear reprocessing.
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
4.5.5. Oxides of group 15 (N, P, As)
The most important members of this group are nitrogen and phosphorus.
Oxides of Nitrogen.
There are five oxides of nitrogen, with N ranging in oxidation number from +1 to +5
N2O, NO, N2O3, NO2, N2O5.
Ox.
No.
Formula
Note
Lewis structure
+1
N2O
Laughing gas
N
+2
+3
NO
N2O3
N
toxic
Blue liquid
N
O
N
N
O
O
+4
NO2
O
O
In equilibrium
with N2O4
O
2 O
N
O
O
N
N
O
+5
O
N2O5
O
N
O
N
O
O
O
O
N
N
O
O
In equilibrium with NO2+ +NO3-
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O
O
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
There are also two important oxoacids of nitrogen: nitrogen(III) acid (nitrous acid) HONO and
nitrogen(V) acid (nitric acid) HNO3.
Nitric acid (HNO3) is produced in nature when NO 2 reacts with water. Nitric acid is highly
reactive, and behaves as an oxidising agent and a nitrating agent as well as an acid.
Industrial quantities of nitric acid are produced using the Ostwald Process.
NH3 (g) + 5/4 O2 (g)
NO (g)
+ 1/2 O2 (g)
3 NO2 (g) + H2O (l)
Pt catalyst
5 atm
850oC
NO (g) + 3/2 H2O (g)
NO2 (g)
2 HNO3 (aq)
+ NO (g)
Nitric acid is an intermediate in the synthesis of NH 4NO3, an important fertiliser.
Oxides of Phosphorus.
The two most common oxides of phosphorus are phosphorus(III) oxide P4O6 and
phosphorus(V) oxide P4O10. Both have a P with sp 3 hybridisation. Phosphorus(V) oxide also
has p-d -bonds between P and O.
Reaction with water.
O
P4O6 (s) + 6 H2O (l)
4 H3PO3 (aq)
HO
HO
P
H
phosphorous acid
O
P4O10 (s) + 6 H2O (l)
4 H3PO4 (aq)
HO
HO
phosphoric acid
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P
OH
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
4.5.6. Oxides of Sulfur
Sulfur dioxide (SO2) and sulfur trioxide (SO3) are commercially and environmentally
important. Both have a sp 2 hybridised S and both have p : d bonding between S and O. S
is hypervalent in both products.
Sulfur dioxide is a toxic gas. It is released by volcanic activities and because coal and
petroleum contain sulfur compounds; their combustion generates sulfur dioxide.
Sulfur dioxide can further oxidise into SO 3 in the presence of NO2, and forms sulfuric acid
(H2SO4) which is partly responsible for acid rain.
Sulfur dioxide is an intermediate in the industrial production of sulfuric acid H 2SO4. The
conversion of SO2 into H2SO4 is achieved using the “Contact Process”.
S8 (s) + 8 O2 (g)
8 SO2 (g)
2 SO2 (g) + O2 (g)
2 SO3 (g)
SO3 (g) + H2O (l)
H2SO4 (aq)
Application of sulfuric acid.
Annual production was 157 million tons in 2000
About 60% is used in the production of H3PO4 (a fertilizer)
Other uses include oil refining, steel manufacturing .
Table. Comparison of oxygen and sulfur chemistry:
Electronegativity
Formal oxidation number
Hybridisation
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Oxygen
 = 3.44
much more electronegative
mostly 
but with exceptions, e.g.:
0 in O2, -1 in H2O2
sp3 (H2O)
sp2 (CO2)
Sulfur
 = 2.58
variable (-2 to +6)
sp3 (H2S, SF2)
sp3d (SF4)
sp3d2 (SF6)
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
4.6. Environmental Impact of Main Group Compounds
The three mains areas where the main group compounds have a harmful environmental
impact are: acid rain, ozone depletion and greenhouse effect.
All these effects are different and are not to be confused.
4.6.1. Acid rain
Acid rain is caused by emissions of sulfur dioxide and nitrogen oxide. When these react with
water in the atmosphere they produce sulfuric acid. The acid rain cause great damage to the
economy, e.g. steel structure (bridges) corrosion, erosion, paint peeling.
SO2 is produced from the oxidation of sulfur compounds, and is produced during the
combustion of coal and oil when they contain sulfur (this leads to concerns over the use of
high-sulfur coal). Sulfur dioxide reacts with water to give H 2SO3, a relatively weak acid.
SO2 (g) + H2O (l)
H2SO3 (aq)
The problem happens when sulfur dioxide oxidizes in the atmosphere into SO 3. Sulfur
trioxide then reacts with atmospheric water to give sulfuric acid.
2 SO2 (g) + O2 (g)
SO3 (g) + H2O (l)
2 SO3 (g)
H2SO4 (aq)
Nitric acid is formed from various nitrogen oxides “NOx”, primarily NO and NO2.
Acidity of the rain.
 Normal rain is slightly acidic, because of solubilised CO 2: its pH is 5.6.
 Acid rains are much more acidic (pH 4.2, in some part of the USA; pH 2.7has been
measured in Sweden).
 The world record was established in West Virginia (USA) in 1980 with pH 1.8.
Effect of the acidity
 Fish/shellfish start to die at pH 5.0 to 4.5
 Forests are harmed (nutrients are removed and toxic substances are released by
acidity).
Technical solutions
 Scrub exhaust gases to remove acidic oxides.
SO2 can be removed with alkaline materials, such as CaO or NaOH.
NOx converted to HNO3.
 Chemically treat acidified lakes (using CaO)
 Catalyse noxious products into more benign species.
Catalytic converters on cars catalyse the transformation of CO into CO2, NOx into
N2, and O2 and unburnt hydrocarbons into CO2.
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
4.6.2. Ozone depletion
The surface of earth is continually bombarded by solar radiation. The ozone layer absorbs
UVB ultraviolet light (280-320nm) from the sun and therefore protects us. UVB photons have
enough energy to break chemical bonds, and damage our DNA which leads to skin cancer.
It is feared that the ozone depletion will lead to an increase in skin cancer.
Ozone cycle
Three allotropes of oxygen are involved in the ozone-oxygen cycle: O (atomic oxygen), O2
(oxygen gas) and O3 (ozone gas).
Ozone is formed in the stratosphere when oxygen molecules photo-dissociate after
absorbing an ultraviolet radiation ( less than 240 nm). This converts a single O 2 molecule
into two highly reactive oxygen atoms.
O2
UV light
2O
Each oxygen atom then combines with a O2 molecule to create O3.
O + O2
O3
ozone formation
Ozone molecule absorbs UV light and splits into a molecule of O 2 and an oxygen atom.
O3
UVB
O + O2
ozone breakdown
This is a continuing process that terminates when an oxygen atom recombines with an
ozone molecule to make two O2 molecules.
O + O3
2 O2
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termination
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
The effects of chlorofluorocarbons on ozone.
Chlorofluorocarbons (CFCs), such as CF2Cl2, were previously used as aerosol propellants,
plastic foam blowing agents, and air-conditioning refrigerants, - had to be replaced due to
their effect on ozone.
Paradoxically, CFCs are very stable in lower atmosphere. However, when in the
stratosphere; high energy UV radiations break the C-Cl bonds and form chlorine radicals Cl.,
which are extremely reactive.
high UV
CF2Cl2
CF2Cl
+ Cl
(Cl
is extremely reactive)
A chlorine radical Cl. reacts with ozone to form a chlorine monoxide radical (.ClO) and
oxygen. The chlorine monoxide radical can go on to react with an oxygen atom to form
oxygen and generate more chlorine radicals.
O3 + Cl
ClO + O
ClO + O2
Cl + O2
In this way, the overall reaction for the decomposition of ozone into oxygen is a
radical chain process initiated by chlorine radicals. The radical chain process via
chlorine radical has a 2 year lifetime.
Technical solutions.
CFCs are replaced as aerosol by other gases such as butane, and as refrigerants by HFCs
(hydrofluorocarbons such as CF3CH2F) although these too can be replaced.
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TOPIC 4. CHEMISTRY OF THE ELEMENTS OF THE MAIN GROUP
4.6.3. Greenhouse Effect.
There is a direct correlation between the global average temperature and the atmospheric
concentration of what are now called the greenhouse gases. The three major greenhouse
gases are: water vapour (accounting for 36 to 70% of the greenhouse effect), carbon dioxide
(9 to 26%), methane (4 to 9%). Other contributors include nitrous oxide, ozone and CFCs.
Carbon dioxide
Carbon dioxide is an important contributor to the greenhouse effect.
Technical solutions
 Use less fossil fuel.
It is believed that a reduction of greenhouse effects could be achieved by reducing our
carbon dioxide emission. Reducing fossil fuel consumption could be a direct way to achieve
that outcome; although this is difficult because the worldwide energy requirement is
increasing dramatically and relying even more than before on fossil fuel consumption.
 CO2 capture AFTER combustion.
Post-combustion capture of CO2 is achieved by using amine solvents, using CaO or
membrane separators.
 CO2 capture BEFORE combustion.
Pre-combustion capture of CO2 is achieved by generating H2 from fuel, and H2 not the fossil
fuel itself is combusted to generate power.
 Oxy-fuel combustion
Pure O2 is used instead of air during the combustion of a fuel. Because pure oxygen is used,
higher temperatures can be achieved and consequently less fuel is necessary. Oxy-fuel
combustion produces about 75% less flue gases compared with air-fuel combustion.
Another benefit - the reduction of nitrogen oxides produced (because nitrogen from air is not
present in the combustion).
Methane
Methane is a very effective greenhouse gas; 21 times more efficient than CO2. and its
importance as a greenhouse gas could have been underestimated. Methane is produced
naturally by methanogenic bacteria in the gut of ruminants.
Note. CFCs are even more effective greenhouse gases, 10,000 times more effective than
CO2. But their small concentration in atmosphere means they are not a significant
contributor to greenhouse effect.
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