The Chemistry of Oxygen and Sulfur
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The Chemistry of Oxygen and Sulfur
The Chemistry of Oxygen and
Sulfur
The Chemistry of
The Chemistry of
Oxygen as an Oxidizing
Oxygen
Ozone
Agent
Peroxides
Methods of Preparing
O2
The Chemistry of Sulfur
The Effect of
The Effect of
The Effect of Differences in
Differences in the
Strength of X-X
and X=X Bonds
Differences in the
Electronegativities of
Sulfur and Oxygen
the Abilities of Sulfur and
Oxygen to Expand Their
Valence Shell
The Chemistry of Oxygen
Oxygen is the most abundant element on this planet. The earth's crust is 46.6% oxygen by
weight, the oceans are 86% oxygen by weight, and the atmosphere is 21% oxygen by volume.
The name oxygen comes from the Greek stems oxys, "acid," and gennan, "to form or
generate." Thus, oxygen literally means "acid former." This name was introduced by Lavoisier,
who noticed that compounds rich in oxygen, such as SO 2 and P4O 10, dissolve in water to give
acids.
[He] 2s2 2p 4
The electron configuration of an oxygen atom
suggests that neutral oxygen
atoms can achieve an octet of valence electrons by sharing two pairs of electrons to form an
O=O double bond, as shown in the figure below.
According to this Lewis structure, all of the electrons in the O 2 molecule are paired. The
compound should therefore be diamagnetic
it should be repelled by a magnetic field.
Experimentally, O 2 is found to be paramagnetic
it is attracted to a magnetic field. This can
be explained by assuming that there are two unpaired electrons in the
* antibonding
molecular orbitals of the O 2 molecule.
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This photograph shows that the liquid
O2 is so strongly attracted to a
magnetic field that it will bridge the
gap between the poles of a horseshoe
magnet.
At temperatures below -183oC, O 2 condenses to form a liquid with a characteristic light blue
color that results from the absorption of light with a wavelength of 630 nm. This absorption is
not seen in the gas phase and is relatively weak even in the liquid because it requires that
three bodies
two O 2 molecules and a photon
collide simultaneously, which is a very
rare phenomenon, even in the liquid phase.
The Chemistry of Ozone
The O 2 molecule is not the only elemental form of oxygen. In the presence of lightning or
another source of a spark, O 2 molecules dissociate to form oxygen atoms.
spark
O2 (g)
2 O(g)
These O atoms can react with O 2 molecules to form ozone, O 3,
O2 (g) + O(g)
O3 (g)
whose Lewis structure is shown in the figure below.
Oxygen (O 2) and ozone (O 3) are examples of allotropes (from the Greek meaning "in another
manner"). By definition, allotropes are different forms of an element. Because they have
different structures, allotropes have different chemical and physical properties (see table
below).
Properties of Allotropes of Oxygen
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Oxygen (O2 ) Ozone (O3 )
Melting Point
-218.75oC
-192.5oC
Boiling Point
-182.96oC
-110.5oC
Density (at 20 oC) 1.331 g/L
O-O bond order
O-O bond length
1.998 g/L
2
0.1207 nm
1.5
0.1278 nm
Ozone is an unstable compound with a sharp, pungent odor that slowly decomposes to oxygen.
3 O3 (g)
3 O2 (g)
At low concentrations, ozone can be relatively pleasant. (The characteristic clean odor
associated with summer thunderstorms is due to the formation of small amounts of O 3.)
Exposure to O 3 at higher concentrations leads to coughing, rapid beating of the heart, chest
pain, and general body pain. At concentrations above 1 ppm, ozone is toxic.
One of the characteristic properties of ozone is its ability to absorb radiation in the ultraviolet
portion of the spectrum ( > 300 nm), thereby providing a filter that protects us from exposure
to high-energy ultraviolet radiation emitted by the sun. We can understand the importance of
this filter if we think about what happens when radiation from the sun is absorbed by our skin.
Electromagnetic radiation in the infrared, visible, and low-energy portions of the ultraviolet
spectrum ( < 300 nm) carries enough energy to excite an electron in a molecule into a higher
energy orbital. This electron eventually falls back into the orbital from which it was excited and
energy is given off to the surrounding tissue in the form of heat. Anyone who has suffered
from a sunburn can appreciate the painful consequences of excessive amounts of this radiation.
Radiation in the high-energy portion of the ultraviolet spectrum ( 300 nm) has a different effect
when it is absorbed. This radiation carries enough energy to ionize atoms or molecules. The
ions formed in these reactions have an odd number of electrons and are extremely reactive.
They can cause permanent damage to the cell tissue and induce processes that eventually
result in skin cancer. Relatively small amounts of this radiation can therefore have drastic
effects on living tissue.
In 1974 Molina and Rowland pointed out that chlorofluorocarbons, such as CFCl 3 and CF 2Cl2,
which had been used as refrigerants and as propellants in aerosol cans, were beginning to
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which had been used as refrigerants and as propellants in aerosol cans, were beginning to
accumulate in the atmosphere. In the stratosphere, at altitudes of 10 to 50 km above the
earth's surface, chlorofluorocarbons decompose to form Cl atoms and chlorine oxides such as
ClO when they absorb sunlight. Cl atoms and ClO molecules have an odd number of electrons,
as shown in the figure below.
As a result, these substances are unusually reactive. In the atmosphere, they react with ozone
or with the oxygen atoms that are needed to form ozone.
Cl + O3
ClO + O2
ClO + O
Cl + O2
Molina and Rowland postulated that these substances would eventually deplete the ozone shield
in the stratosphere, with dangerous implications for biological systems that would be exposed
to increased levels of high-energy ultraviolet radiation.
Oxygen as an Oxidizing Agent
Fluorine is the only element that is more electronegative than oxygen. As a result, oxygen
gains electrons in virtually all its chemical reactions. Each O 2 molecule must gain four electrons
to satisfy the octets of the two oxygen atoms without sharing electrons, as shown in the figure
below.
Oxygen therefore oxidizes metals to form salts in which the oxygen atoms are formally present
as O 2- ions. Rust forms, for example, when iron reacts with oxygen in the presence of water to
give a salt that formally contains the Fe 3+ and O 2- ions, with an average of three water
molecules coordinated to each Fe 3+ ions in this solid.
H 2O
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4 Fe(s) + 3 O2 (g)
2 Fe2 O3 (s) 3 H 2 O
Oxygen also oxidizes nonmetals, such as carbon, to form covalent compounds in which the
oxygen formally has an oxidation number of -2.
C(s) + O2 (g)
CO2 (g)
Oxygen is the perfect example of an oxidizing agent because it increases the oxidation state
of almost any substance with which it reacts. In the course of its reactions, oxygen is reduced.
The substances it reacts with are therefore reducing agents.
Peroxides
It takes four electrons to reduce an O 2 molecule to a pair of O 2- ions. If the reaction stops
after the O 2 molecule has gained only two electrons, the O 22- ion shown in the figure below is
produced.
This ion has two more electrons than a neutral O 2 molecule, which means that the oxygen
atoms must share only a single pair of bonding electrons to achieve an octet of valence
electrons. The O 22- ion is called the peroxide ion because compounds that contain this ion are
unusually rich in oxygen. They are not just oxides
they are (hy-)peroxides.
The easiest way to prepare a peroxide is to react sodium or barium metal with oxygen.
2 Na(s) + O2 (g)
Ba(s) + O2 (g)
Na2 O2 (s)
BaO 2 (s)
When these peroxides are allowed to react with a strong acid, hydrogen peroxide (H2O 2) is
produced.
BaO 2 (s) + 2 H +(aq)
Ba2+(aq) + H 2 O2 (aq)
The Lewis structure of hydrogen peroxide contains an O-O single bond, as shown in the figure
below.
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below.
The VSEPR theory predicts that the geometry around each oxygen atom in H2O 2 should be
bent. But this theory cannot predict whether the four atoms should lie in the same plane or
whether the molecule should be visualized as lying in two intersecting planes. The
experimentally determined structure of H2O 2 is shown in the figure below.
The H-O-O bond angle in this molecule is only slightly larger than the angle between a pair of
adjacent 2p atomic orbitals on the oxygen atom, and the angle between the planes that form
the molecule is slightly larger than the tetrahedral angle.
The oxidation number of the oxygen atoms in hydrogen peroxide is -1. H2O 2 can therefore act
as an oxidizing agent and capture two more electrons to form a pair of hydroxide ions, in which
the oxygen has an oxidation number of -2.
H 2 O2 + 2 e-
2 OH-
Or, it can act as a reducing agent and lose a pair of electrons to form an O 2 molecule.
H 2 O2
O2 + 2 H + + 2 e-
Reactions in which a compound simultaneously undergoes both oxidation and reduction are
called disproportionation reactions. The products of the disproportionation of H2O 2 are
oxygen and water.
2 H 2 O2 (aq)
O2 (g) + 2 H 2 O(l)
The disproportionation of H2O 2 is an exothermic reaction.
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The Chemistry of Oxygen and Sulfur
2 H 2 O2 (aq)
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O2 (g) + 2 H 2 O(l)
H o = -94.6 kJ/mol H 2 O
This reaction is relatively slow, however, in the absence of a catalyst, such as dust or a metal
surface. The principal uses of H2O 2 revolve around its oxidizing ability. It is used in dilute (3%)
solutions as a disinfectant. In more concentrated solutions (30%), it is used as a bleaching
agent for hair, fur, leather, or the wood pulp used to make paper. In very concentrated
solutions, H2O 2 has been used as rocket fuel because of the ease with which it decomposes to
give O 2.
Methods of Preparing O 2
Small quantities of O 2 gas can be prepared in a number of ways.
1. By decomposing a dilute solution of hydrogen peroxide with dust or a metal surface as the
catalyst.
2 H 2 O2 (aq)
O2 (g) + 2 H 2 O(l)
2. By reacting hydrogen peroxide with a strong oxidizing agent, such as the permanganate ion,
MnO 4-.
5 H 2 O2 (aq) + 2 MnO 4 -(aq) + 6 H +(aq)
2 Mn2+(aq) + 5 O2 (g) + 8 H 2 O(l)
3. By passing an electric current through water.
electrolysis
2 H 2 O(l)
2 H 2 (g) + O2 (g)
4. By heating potassium chlorate (KClO3) in the presence of a catalyst until it decomposes.
MnO 2
2 KClO 3 (s)
2 KCl(s) + 3 O2 (g)
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The Chemistry of Sulfur
Because sulfur is directly below oxygen in the periodic table, these elements have similar
electron configurations. As a result, sulfur forms many compounds that are analogs of oxygen
compounds, as shown in the table below. Examples in this table show how the prefix thio- can
be used to indicate compounds in which sulfur replaces an oxygen atom. The thiocyanate
(SCN -) ion, for instance, is the sulfur-containing analog of the cyanate (OCN-) ion.
Oxygen Compounds and Their Sulfur Analogs
Oxygen Compounds
Sulfur Compounds
Na2 O (sodium oxide) Na2 S (sodium sulfide)
H 2 O (water)
H 2 S (hydrogen sulfide)
O3 (ozone)
SO2 (sulfur dioxide)
CO2 (carbon dioxide) CS 2 (carbon disulfide)
OCN- (cyanate)
SCN - (thiocyanate)
OC(NH 2 ) 2 (urea)
SC(NH 2 ) 2 (thiourea)
There are four principal differences between the chemistry of sulfur and oxygen.
1. O=O double bonds are much stronger than S=S double bonds.
2. S-S single bonds are almost twice as strong as O-O single bonds.
3. Sulfur (EN = 2.58) is much less electronegative than oxygen (EN = 3.44).
4. Sulfur can expand its valence shell to hold more than eight electrons, but oxygen cannot.
These seemingly minor differences have important consequences for the chemistry of these
elements.
The Effect of Differences in the Strength of X-X and X=X Bonds
The radius of a sulfur atom is about 60% larger than that of an oxygen atom.
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As a result, it is harder for sulfur atoms to come close enough together to form bonds. S=S
double bonds are therefore much weaker than O=O double bonds.
Double bonds between sulfur and oxygen or carbon atoms can be found in compounds such as
SO 2 and CS2 (see figure below). But these double bonds are much weaker than the equivalent
double bonds to oxygen atoms in O 3 or CO 2. The bond dissociation enthalpy for a C=S double
bond is 477 kJ/mol, for example, whereas the bond dissociation enthalpy for a C=O double
bond is 745 kJ/mol.
Elemental oxygen consists of O 2 molecules in which each atom completes its octet of valence
electrons by sharing two pairs of electrons with a single neighboring atom. Because sulfur does
not form strong S=S double bonds, elemental sulfur usually consists of cyclic S8 molecules in
which each atom completes its octet by forming single bonds to two neighboring atoms, as
shown in the figure below.
S8 molecules can pack to form more than one crystal. The most stable form of sulfur consists of
orthorhombic crystals of S8 molecules, which are often found near volcanoes. If these crystals
are heated until they melt and the molten sulfur is then cooled, an allotrope of sulfur consisting
of monoclinic crystals of S8 molecules is formed. These monoclinic crystals slowly transform
themselves into the more stable orthorhombic structure over a period of time.
The tendency of an element to form bonds to itself is called catenation (from the Latin catena,
"chain"). Because sulfur forms unusually strong S-S single bonds, it is better at catenation than
any element except carbon. As a result, the orthorhombic and monoclinic forms of sulfur are
not the only allotropes of the element. Allotropes of sulfur also exist that differ in the size of
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the molecules that form the crystal. Cyclic molecules that contain 6, 7, 8, 10, and 12 sulfur
atoms are known.
Sulfur melts at 119.25 oC to form a yellow liquid that is less viscous than water. If this liquid is
heated to 159oC, it turns into a dark red liquid that cannot be poured from its container. The
viscosity of this dark red liquid is 2000 times greater than that of molten sulfur because the
cyclic S8 molecules open up and link together to form long chains of as many as 100,000 sulfur
atoms.
When sulfur reacts with an active metal, it can form the sulfide ion, S2-.
16 K(s) + S8 (s)
8 K 2 S(s)
This is not the only product that can be obtained, however. A variety of polysulfide ions with a
charge of -2 can be produced that differ in the number of sulfur atoms in the chain.
The Effect of Differences in the Electronegativities of Sulfur and Oxygen
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Because sulfur is much less electronegative than oxygen, it is more likely to form compounds in
which it has a positive oxidation number (see table below).
Common Oxidation Numbers for Sulfur
Oxidation
Number
Examples
-2
Na2 S, H 2 S
-1
Na2 S2 , H 2 S2
0
S8
+1
S2 Cl2
+2
S2 O3 2-
+2 1 / 2
S4 O6 2-
+3
S2 O4 2-
+4
SF 4 , SO2 , H 2 SO3 , SO3 2-
+5
S2 O6 2-
+6
SF 6 , SO3 , H 2 SO4 , SO4 2-
In theory, sulfur can react with oxygen to form either SO 2 or SO 3, whose Lewis structures are
given in the figure below.
SO2
SO3
In practice, combustion of sulfur compounds gives SO 2, regardless of whether sulfur or a
compound of sulfur is burned.
S8 (s)
+ 8 O2 (g)
8 SO2 (g)
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CS 2 (l)
+ 3 O2 (g)
3 FeS2 (s) + 8 O2 (g)
CO2 (g)
+ 2 SO2 (g)
Fe3 O4 (s) + 6 SO2 (g)
Although the SO 2 formed in these reactions should react with O 2 to form SO 3, the rate of this
reaction is very slow. The rate of the conversion of SO 2 into SO 3 can be greatly increased by
adding an appropriate catalyst.
V2O 5/K2O
2 SO2 (g)
2 SO3 (g)
Enormous quantities of SO 2 are produced by industry each year and then converted to SO 3,
which can be used to produce sulfuric acid, H2SO 4. In theory, sulfuric acid can be made by
dissolving SO 3 gas in water.
SO3 (g) + H 2 O(l)
H 2 SO4 (aq)
In practice, this is not convenient. Instead, SO 3 is absorbed in 98% H2SO 4, where it reacts
with the water to form additional H2SO 4 molecules. Water is then added, as needed, to keep
the concentration of this solution between 96% and 98% H2SO 4 by weight.
Sulfuric acid is by far the most important industrial chemical. It has even been argued that
there is a direct relationship between the amount of sulfuric acid a country consumes and its
standard of living. More than 50% of the sulfuric acid produced each year is used to make
fertilizers. The rest is used to make paper, synthetic fibers and textiles, insecticides,
detergents, feed additives, dyes, drugs, antifreeze, paints and enamels, linoleum, synthetic
rubber, printing inks, cellophane, photographic film, explosives, automobile batteries, and
metals such as magnesium, aluminum, iron, and steel.
Sulfuric acid dissociates in water to give the HSO4- ion, which is known as the hydrogen
sulfate, or bisulfate, ion.
H 2 SO4 (aq)
H +(aq) + HSO4 -(aq)
10% of these hydrogen sulfate ions dissociate further to give the SO 42-, or sulfate, ion.
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HSO4 -(aq)
2H +(aq) + SO4 (aq)
A variety of salts can be formed by replacing the H+ ions in sulfuric acid with positively charged
ions, such as the Na + or K + ions.
NaHSO 4 = sodium hydrogen sulfate
Na2 SO4 = sodium sulfate
Sulfur dioxide dissolves in water to form sulfurous acid.
SO2 (g) + H 2 O(l)
H 2 SO3 (aq)
Sulfurous acid doesn't dissociate in water to as great extent as sulfuric acid, but it is still
possible to replace the H+ ions in H2SO 3 with positive ions to form salts.
NaHSO 3 = sodium hydrogen sulfite
Na2 SO3 = sodium sulfite
Sulfuric acid and sulfurous acid are both examples of a class of compounds known as
oxyacids, because they are literally acids that contain oxygen. Because they are negative ions
(or anions) that contain oxygen, the SO 32- and SO 42- ions are known as oxyanions. The Lewis
structures of some of the oxides of sulfur that form oxyacids or oxyanions are given in the
table below.
OXYACIDS
OXYANIONS
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One of these oxyanions deserves special mention. This ion, which is known as the thiosulfate
ion, is formed by the reaction between sulfur and the sulfite (SO 32-) ion.
8 SO3 2-(aq) + S8 (s)
8 S2 O3 2-(aq)
The Effect of Differences in the Abilities of Sulfur and Oxygen to Expand Their Valence
Shell
The electron configurations of oxygen and sulfur are usually written as follows.
O = [He] 2s 2 2p4
S = [Ne] 3s 2 3p4
Although this notation shows the similarity between the configurations of the two elements, it
hides an important difference that allows sulfur to expand its valence shell to hold more than
eight electrons.
Oxygen reacts with fluorine to form OF2.
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O2 (g) + 2 F 2 (g)
2 OF2 (g)
The reaction stops at this point because oxygen can hold only eight electrons in its valence
shell, as shown in the figure below.
Sulfur reacts with fluorine to form SF 4 and SF 6, shown in the figure below, because sulfur can
expand its valence shell to hold 10 or even 12 electrons.
S8 (s) + 16 F 2 (g)
8 SF 4 (g)
S8 (s) + 24 F 2 (g)
8 SF 6 (g)
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