Download CHEMISTRY OF p-ELEMENTS - Львівський національний

DANYLO HALYTSKY LVIV NATIONAL MEDICAL UNIVERSITY
DEPARTMENT of GENERAL, BIOINORGANIC, PHYSICAL and
COLLOIDAL CHEMISTRY
V.V. Ogurtsov, O.V. Klenina, O.M. Roman, O.I. Marshalok
CHEMISTRY OF p-ELEMENTS
Lectures for the 1st year students
оf pharmaceutical faculty
(Module 2. Inorganic chemistry)
LVIV – 2011
ЛЬВІВСЬКИЙ НАЦІОНАЛЬНИЙ МЕДИЧНИЙ УНІВЕРСИТЕТ
імені Данила Галицького
КАФЕДРА ЗАГАЛЬНОЇ, БІОНЕОРГАНІЧНОЇ ТА ФІЗКОЛОЇДНОЇ ХІМІЇ
В.В. Огурцов, О.В. Кленіна, О.М. Роман, О.І. Маршалок
ХІМІЯ p-ЕЛЕМЕНТІВ
Тексти лекцій для студентів I курсу
фармацевтичного факультету
(Модуль 2. Неорганічна хімія)
ЛЬВІВ – 2011
2
Посібник обговорено і схвалено до друку цикловою методичною
комісією з фізико-хімічних дисциплін фармацевтичного факультету
(протокол № 2 від 23 червня 2011 р).
Рецензент: проф. Й.Д. Комариця – професор кафедри фармацевтичної, органічної та біоорганічної хімії ЛНМУ імені
Данила Галицького.
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Introduction
The elements of p-series of the Periodic table are the representative elements of
III – VIII groups.
Most of p-elements are of a vital significance for living systems. Five of them
(carbon, oxygen, nitrogen, sulfur and phosphorus) as well as hydrogen (s-element)
are the structural material for bio systems formation.
Carbon is key to life and by definition is present in all organic compounds
including proteins, nucleic acids, hydrocarbons, enzymes, vitamins. The study of
life is known as biological chemistry or biochemistry.
Oxygen atoms are present in water (H2O) and water is essential to all life.
Oxygen is present in many organic compounds. Most organisms use oxygen for
respiration. It takes part in all kinds of matter renewal. The reactions of oxygen
with lipids, proteins and with glucose for the first time are the source of energy for
living organisms. While oxygen (O2) is necessary for life, oxygen as ozone (O3) is
highly toxic. On the other hand, ozone is an important component of the
atmosphere (the ozone layer) and helps to shield us from harmful ultraviolet rays
from the sun.
Nitrogen is a key component of biological molecules such as proteins (which
are made from amino acids), and nucleic acids (DNA, RNA). The nitrogen cycle in
nature is very important.
Sulphur is also essential to life. It is a minor constituent of fats, body fluids, and
skeletal minerals. Sulphur is a key component in most proteins since it is contained
in the amino acids methionine and cysteine. Sulphur-sulphur interactions (-S-S
bridges) are important in determining protein tertiary structure. Hydrogen sulphide
(H2S) replaces H2O in the photosynthesis of some bacteria.
Phosphorus is a key component of biological molecules such as DNA and
RNA. Phosphorus is a component of bones, and teeth, and many other compounds
required for life.
Some other p-elements are also of a great importance for living systems.
Chlorine, bromine, iodine, fluorine, boron, aluminium, tin and lead are among
them. Blood, skin, bones, tissues, liver, lungs, brain and other organs contain small
amounts of these elements. Some of them take part in bio chemical processes.
Many inorganic compounds of p-elements have a wide range of applications in
pharmacy, dentistry and medicine.
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1. Group IIIA. Boron, Aluminum, Gallium, Indium, and
Thallium
There are five chemical elements in group IIIA of the periodic table: boron B,
aluminum Al, gallium Ga, indium In and thallium Tl.
The atoms of these elements have the following configuration:
– 1s22s22p1
[He]2s22p1
5B
2 2
6 2
1
– 1s 2s 2p 3s 3p
[Ne]3s23p1
13Al
2 2
6 2
6
10 2
1
– 1s 2s 2p 3s 3p 3d 4s 4p
[Ar] 4s24p1
31Ga
2 2
6 2
6
10 2
6
10 2
1
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p
[Kr] 5s25p1
49In
2 2
6 2
6
10 2
6
10 14 2
6
10 2
1
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s 6p
[Xe]6s26p1
81Tl
Table 1.1 lists some physical properties of group IIIA elements.
Table 1.1 Properties of Group IIIA Elements
Property
Electron configuration
Melting point, °C
Boiling point, °C
Density, g/cm3
Electronegativity
(Pauling scale)
Covalent radius, Å
Ionic radius, Å
Boron
Aluminum
Gallium
Indium
Thallium
[He]2s22p1 [Ne]3s23p1 [Ar]4s24p1 [Kr]5s25p1 [Xe]6s26p1
2076
660
30
157
304
3927
2519
2204
2072
1473
2.46
2.7
5.9
7.3
11.85
2.0
1.6
1.8
1.7
1.6
0.91
0.20
1.43
0.57
1.39
0.62
1.66
0.92
1.71
1.05
Physical Properties and Occurrence
Boron, symbol B, hard, brittle semimetallic element with an atomic number of
5. Pure boron, as usually prepared, is an amorphous powder. A crystalline form can
be prepared, however, by dissolving boron in molten aluminum and cooling
slowly. The atomic weight of boron is 10.81; the element melts at about 2076 °C,
boils at about 3927 °C, and has a specific gravity of 2.46.
Compounds of boron, notably borax, have been known since early times, but
the pure element was firstly prepared in 1808 by the French chemists Joseph GayLussac and Baron Louis Thinard, and independently by the British chemist Sir
Humphry Davy. Important boron ores are ulexite (NaCaB5O9 · 8H2O), colemanite
(Ca2B6O11 · 5H2O), and kernite (Na2B4O7 · 4H2O). Boron ranks about 38th in
natural abundance among the elements in the earth's crust.
Aluminum (in Canada and Europe, aluminium), symbol Al, the most abundant
metallic element in the earth's crust. The atomic number of aluminum is 13. It is a
lightweight, silvery metal. The atomic weight of aluminum is 26.9815; the element
melts at 660 °C, boils at 2519 °C, and has a specific gravity of 2.7.
5
Hans Christian Oersted, a Danish chemist, was the first who isolated aluminum
in 1825, using a chemical process involving potassium amalgam. Between 1827
and 1845, Friedrich Wohler, a German chemist, improved Oersted's process by
using metallic potassium. He was the first to measure the specific gravity of
aluminum and show its lightness. In 1854 Henri Sainte-Claire Deville, in France,
obtained the metal by reducing aluminum chloride with sodium. Aided by the
financial backing of Napoleon III, Deville established a large-scale experimental
plant and displayed pure aluminum at the Paris Exposition in 1855.
Aluminum is the most abundant metallic constituent in the earth’s crust; only
the nonmetals namely oxygen and silicon are more abundant. Aluminum is never
found as a free metal; commonly as aluminum silicate or as a silicate of aluminum
mixed with other metals such as sodium, potassium, iron, calcium, and magnesium.
However, the extracting of aluminum from these silicates is chemically difficult
and therefore an expensive process that is why they are not useful ores. Bauxite
(consisting of aluminum oxide of various degrees of hydration) is the commercial
source of aluminum and its compounds.
Gallium, symbol Ga, is a metallic element that remains in the liquid state over a
wider range of temperatures than any other element. Gallium is blue-gray in color
as a solid and silvery as a liquid. Its atomic number is 31. Gallium melts at 30 °C,
boils at about 2204 °C, and has specific gravity of 5.9; the atomic weight of the
element is 69.72.
Gallium was discovered spectroscopically by the French chemist Paul Imile
Lecoq de Boisbaudran in 1875; a year later he isolated the element in its metallic
state. The element is about 34th in order of abundance in the earth’s crust. Gallium
occurs in small quantities in some varieties of zinc blende, bauxite, pyrite,
magnetite and kaolin.
Indium, symbol In, is a soft, malleable, silvery white metallic element. Its
atomic number is 49. Indium melts at about 157 °C, boils at about 2072 °C, and
has a specific gravity of 7.3. The atomic weight of indium is 114.82.
Indium was discovered spectroscopically in 1863 by the German chemists
Hieronymus Theodor Richter and Ferdinand Reich. It ranks 63rd in order of
abundance of the elements in the earth’s surface. Indium never occurs as a free
metal and is usually found as the sulfide In2S3; in certain zinc blendes; and in
tungsten, tin, and iron ores.
Thallium (Greek thallos, “young shoot”), symbol Tl, soft, malleable metallic
element that acquires a bluish-gray color upon exposure to the atmosphere. The
atomic number of thallium is 81. Thallium melts at about 304 °C, boils at about
1473 °C, and has a specific gravity of 11.85. The atomic weight of thallium is
204.38.
Thallium was discovered spectroscopically in 1861 by the British chemist Sir
William Crookes. It was isolated by Crookes and, independently, by the French
chemist Claude August Lamy in 1862. Thallium ranks 60th in abundance among the
elements in the earth’s crust. Thallium occurs in combination with pyrites, zinc
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blende, and hematite and is often recovered from the flue dust produced by pyrites
ovens in which sulfur and iron are separated.
Preparation of the Elements
The most common sources of boron are tourmaline, borax
(Na2B4O5(OH)4.8H2O), and kernite [Na2B4O5(OH)4.2H2O]. It is difficult to obtain
pure boron. It can be made through the magnesium reduction of B2O3 oxide. The
oxide is made by melting boric acid, B(OH)3, which in turn is obtained from borax.
B2O3 + 3Mg → 2B + 3MgO.
Some amounts of high purity boron are available through the thermal
decomposition of such boron compounds as BBr3 with hydrogen gas using a heated
tantalum wire. Results are better when hot wires at temperatures over 1000 °C are
used:
2BBr3 + 3H2 → 2B + 6HBr.
Aluminium is mined in huge scales as bauxite (typically Al2O3·2H2O). Bauxite
contains Fe2O3, SiO2, and other impurities. In order to isolate pure aluminium,
these impurities must be removed from the bauxite. This is done by the Bayer
process. It involves bauxite treatment with sodium hydroxide solution, which
results in a solution of sodium aluminate and sodium silicate:
Al2O3 + 6NaOH + 3H2O → 2Na3[Al(OH)6];
SiO2 + 2NaOH → Na2SiO3 + H2O.
The iron remains behind as a solid. When CO2 is blown through the resulting
solution, the sodium silicate stays in solution while the aluminium is precipitated
out as aluminium hydroxide:
2Na3[Al(OH)6] + 3CO2 → 2Al(OH)3↓ + 3Na2CO3 + 3H2O.
The hydroxide can be filtered off, washed, and heated to form pure alumina,
Al2O3:
∆ Al O + 3H O.
2Al(OH) →
3
2
3
2
The next stage is formation of pure aluminium. It is obtained from the pure
Al2O3 by an electrolytic method.
In aqueous solution aluminium oxide dissociates into ions:
On anode:
Al3+ + 3e− → Al
Al2O3 ⇄ Al3+ + AlO33-.
On cathode:
4AlO33- - 12e− → 2Al2O3 + 3O2↑
electrolysis, ∆
2Al2O3    → 4Al + 3O2↑.
Electrolysis is necessary as aluminium is very electropositive. It seems these
days that electrolysis of the hot oxide in a carbon lined steel cell acting as the
cathode with carbon anodes is most common.
Gallium is normally a byproduct of the manufacture of aluminium. The
purification of bauxite by the Bayer process results in concentration of gallium in
its ratio from 5000 to 300 in the alkaline solutions from an aluminium. Electrolysis
using a mercury electrode provides a further concentration and following
7
electrolysis of the resulting sodium gallate using a stainless steel cathode affords
liquid gallium metal.
Preparation of very pure gallium requires a number of further processes ending
with zone refining to make very pure gallium metal.
Indium is a byproduct of the formation of lead and zinc. Indium metal is
isolated by the electrolysis of indium salts in water. Further processes are required
to make very pure indium for electronics purposes.
Crude thallium is present as a component in flue dust along with arsenic,
cadmium, indium, germanium, lead, nickel, selenium, tellurium, and zinc.
Thallium is prepared by dissolving of flue dust in diluted acid, precipitating out
lead sulfate, and then adding HCl to precipitate thallium chloride, TlCl. Further
purification can be achieved by electrolysis of soluble thallium salts.
Chemical Properties of Boron and its Compounds
Boron is the first element in group III. In many of its properties it differs from
the next element in the group, aluminum, which is a metal. Although it is
somewhat metallic in appearance, it is a very poor conductor of electricity. It is
best regarded as a semimetal, like silicon.
CHEMICAL PROPERTIES OF BORON
Boron shows trivalency in its compounds. It is difficult to get B3+ because large
amount of energy is required and, therefore, boron forms tricovalent compounds. It
shows common oxidation state of +3 in majority of its compounds. However boron
shows also an oxidation state of –3 in the metal borides, e.g., in Mg3B2
(magnesium boride) oxidation state of boron is –3.
Boron forms oxide B2O3 when heated in oxygen atmosphere at high
temperature:
T = 700 0 C
4B + 3O2  
→ 2B2O3 (Boron oxide or Boric anhydride).
Boron can form trichloride either by passing chlorine over the heated boron or
by passing chlorine over the heated mixture of its oxide and charcoal:
∆
→
2BCl3;
∆
B2O3 + 3C + 3Cl2 
→ 2BCl3 + 3CO.
2B + 3Cl2
BCl3 is hydrolyzed by water:
BCl3 + 3H2O → H3BO3 + 3HCl.
Boron forms also nitride BN when heated in the atmosphere of nitrogen or
ammonia and sulfide B2S3 when heated with sulfur:
∆
→
2BN;
∆
2B + 2NH3 
→ 2BN + 3H2;
∆
2B + 3S 
→
B2S3.
2B + N2
The nitride and sulfide undergo hydrolysis with steam:
8
BN + 3H2O → H3BO3 + NH3;
Boric acid
B2S3 + 6H2O → 2H3BO3 + 3H2S.
Boron reacts with steam at red heat liberating hydrogen:
∆
2B + 3H2O 
→ B2O3 + 2H2.
Boron reacts with concentrated H2SO4 and evolves sulfur dioxide, SO2:
2B + 3H2SO4(conc.) → 2H3BO3 + 3SO2↑.
It can also dissolve in alkalies and evolve hydrogen:
2B + 6NaOH → 2Na3BO3 + 3H2↑.
sodium ortho-borate
Boron acts as powerful reducing agent:
4B + 3CO2 → 2B2O3 + 3C;
4B + 3SiO2 → 2B2O3 + 3Si.
BORIC ACID AND BORATES
Boric acid and the borates are among the simplest and most important of the
boron compounds. Boric acid, B(OH)3, is a stable, colorless crystalline compound
that forms thin, plate like crystals. It consists of planar molecules with an
equilateral triangular AHal3 geometry around boron. The molecules are held
together in flat sheets by hydrogen bonds:
.
Boric acid is a very weak monoprotic acid (Ka = 6.0×10–10). It ionizes in water
in an unusual way. Instead of donating one of its hydrogen atoms to a water
molecule; it removes an OH− from a water molecule, leaving an H+ ion, which
combines with another water molecule to give an H3O+ ion:
H3BO3 + 2H2O → B(OH)4– + H3O+ .
The boron atom has a vacant 2p orbital, which can accept an electron pair
from a water molecule. This water molecule, simultaneously or in a subsequent
step, gives a proton to another water molecule. In this way, boron completes its
octet by forming the tetrahedral borate ion B(OH)4–. When it is heated, boric acid
loses water to form various condensed boric acids, such as cyclic metaboric acid:
Further heating gives boric oxide, B2O3 which is often formed in an amorphous
form, or glass. The crystalline form has a complex network structure based on
planar BO3 groups. A few of the salts of boric acid, the borates, contain the simple
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anion BO33– but most are derived from condensed forms of the acid. The most
common and most important is sodium tetraborate, Na2[B4O5(OH)4]·8H2O, which
is commonly called borax and for which the formula is often written as
Na2B4O7·10H2O. The structure of the tetraborate anion, B4O5(OH)42– is shown
below:
Borax and other borates are the only important boron minerals. They are found
only in a very few places, for example, in California and Turkey, but the deposits
in those places are extremely large.
Aqueous solutions of borax are basic, because tetraborate is the anion of a
weak acid and, therefore, it is a weak base:
B4O5(OH)42– + 5H2O → 4H3BO3 + 2OH–.
Since calcium and magnesium borates are insoluble, borax is used as a water
softener and as a component of washing powders.
BORON HALIDES
The boron halides are typical covalent nonmetal halides. Boron trifluoride, BF3,
and boron trichloride, BCl3, are gases at room temperature; BBr3 is a liquid, and
BI3 is a solid. They all consist of molecules with the expected AX3 planar triangular
structure. Although the electronegativity difference between boron and fluorine is
2.1, boron trifluoride is a covalent molecular compound with polar B–F bonds
rather than an ionic crystal containing B3+ and F– ions. The electronegativity
difference between two elements is not always a reliable guide to finding out
whether they will form an ionic crystal or a molecular compound.
Because of the presence of the vacant 2p orbital on the boron atom in the boron
halides, they are rather reactive compounds. For example, BF3 reacts with an F– ion
to form BF4– in which the valence shell of boron is completed
BF3 + NaF → NaBF4.
Boron trifluoride reacts in a similar way with many other molecules and ions
which have unshared electron pairs that can be donated to its vacant 2p orbital.
Thus it forms compounds with ammonia and other amines and with ethers and
alcohols. In all these cases BF3 is behaving as a Lewis acid, that is, as an electron
pair acceptor:
.
The boron halides react with water forming boric acid and the hydrogen
halides. For example,
10
BCl3 + 3H2O → H3BO3 + 3HCl.
It occurs readily because the vacant 2p orbital of the boron atom can accept an
electron pair from a water molecule. In contrast, the carbon tetrahalides have no
vacant orbital and they do not react with water.
BORANES
At least twenty five compounds of boron and hydrogen have been prepared.
They are known as the boranes. They all have unexpected formulas such as B2H6,
B4H10, B5H9, and B10H14. We would expect the formula of the simplest borane to
be BH3, but this molecule is unknown as a stable species. The simplest borane that
can be isolated is diborane, B2H6; its structure is
The unusual and unexpected feature of this structure is that there are two
hydrogen atoms, called bridging hydrogens, shared between the two borons.
However, there are not enough electrons for each of the lines shown in the
structure to represent an electron pair. Each atom of boron contributes 3 electrons
and each atom of hydrogen 1 electron, making a total of 12 electrons, or six pairs
for the molecule. Thus there can be a maximum of only six ordinary covalent
bonds, whereas the structure appears to have eight bonds. Because B2H6 has too
few electrons for all the atoms to be held together by normal electron pair bonds
between two nuclei, it is often described as an electron-deficient molecule. The
bonding in diborane is best described as involving two three-center bonds, in
which one electron pair holds together three rather than two nuclei. Each boron
atom is surrounded by four electron pairs, which have the expected tetrahedral
arrangement. But two of these electron pairs form three-center bonds in which one
electron pair holds together two boron nuclei and a hydrogen nucleus. In the
following diagram of the structure the two electron pairs forming the three-center
bonds are denoted by :
The following alternative but equivalent representation of the bonding in
diborane emphasizes its close relationship to the bent-bond model of ethene:
In this way each boron atom completes on octet in its valence shell, whereas in a
BH3 molecule the boron atom would have only six electrons in its valence shell
11
and an empty 2p orbital. Another way to represent the structure of this molecule is
by means of the two resonance structures.
A molecule of diborane readily adds two hydride ions forming two borohydride
ions, BH4–:
B2H6 + 2H– → 2BH4–;
B2H6 + 2LiH → 2LiBH4.
This ion has the expected tetrahedral AX4 structure and ordinary two-center,
electron pair bonds. Salts such as lithium borohydride, LiBH4, and sodium
borohydride, NaBH4, are good reducing agents with wide application in organic
chemistry. The higher boranes have unusual and fascinating structures (see Figure
1.1) that can not be explained in terms of simple bonding theories, so they are
usually discussed in terms of the molecular orbital theory. There are also many
anions derived from the boranes that are more complicated than BH4–. A
particularly fascinating example is B12H122– (Figure 1.2). It has the shape of an
icosahedron, which has 12 equivalent vertices, 20 equilateral triangular faces, and
30 equivalent edges.
Figure 1.1 Structures of Some Higher Boranes
12
Figure 1.2 Structure of the B12H122– Anion
The bonding in the electron-deficient B12H122– can be satisfactorily described
only by means of molecular orbital theory. The structures of these remarkable
compounds serve to illustrate the fact that although the simple ideas of chemical
bonding presented in this book enable us to understand the structures of many
compounds, they are inadequate for many others. A more detailed interpretation of
the chemical bond is necessary to enable the understanding of such compounds.
Although our understanding of chemical bonding has progressed considerably
since Lewis first proposed the idea of the shared electron pair, many aspects of this
subject are still not fully understood. Chemists continue to prepare new
compounds, the structures of which present a challenge for even the most
sophisticated theories.
Chemical Properties of Aluminum and its Compounds
Like Boron Aluminium shows trivalency. Aluminium also forms covalent
compounds. However it can form electrovalent compounds when it combines with
strong electron accepting atoms or groups. The size of Al3+ ion is very small. On
account of small size and high charge, it has high polarizing power. These accounts
for a covalent character even in the case of electrovalent compounds, for example,
AlCl3, AlBr3, and AlI3 have covalent nature. Aluminium shows common oxidation
state of +3 in a majority of its compounds.
Aluminium is not affected by dry air but in moist air a thin film of oxide is
formed over its surface. It burns in oxygen producing brilliant light:
4Al + 3O2 →2Al2O3 + ∆Q.
The reaction is highly exothermic.
Besides oxygen, aluminium reacts with nonmetals directly to form
corresponding compounds. When heated in the atmosphere of nitrogen or ammonia
it forms aluminium nitride:
∆
→
2AlN;
∆
2Al + 2NH3 
→ 2AlN + 3H2.
2Al + N2
Aluminium powder when fused with sulfur forms aluminium sulfide:
13
∆
2Al + 3S 
→ Al2S3.
Finely powdered heated aluminium combines with halogens or by passing
halogens over heated mixture of its oxide and charcoal to form corresponding
halides:
∆
→
2AlHal3;
∆
Al2O3 + 3C + 3Hal2 
→ 2AlHal3 + 3CO.
2Al + 3Hal2
(Hal2 = F2, Cl2, Br2, I2)
All these compounds are hydrolysed with water:
AlN + 3H2O → Al (OH)3 + NH3;
Al2S3 + 6H2O → 2Al (OH)3 + 3H2S;
Al2Hal3 + 3H2O → 2Al (OH)3 + 3HHal.
Pure aluminium is not affected by pure water. The impure aluminium is readily
corroded by water containing salts (sea water). Aluminium decomposes boiling
water evolving hydrogen:
2Al + 6H2O → 2Al(OH)3 + 3H2↑.
Aluminium does not interact with hydrogen directly. Aluminium hydride can
be prepared by the reaction of LiH with aluminium chloride in the etheric solution:
AlCl3 + 3LiH → AlH3 + 3LiCl.
The oxidation potential of aluminium is 1.66 V. Thus, it is strongly
electropositive, very reactive and a powerful reducing agent. It dissolves in HCl
(dilute and concentrated) and dilute sulfuric acid, evolving hydrogen:
2Al + 6HCl → 2AlCl3 + 3H2↑;
2Al + 3H2SO4 → Al2(SO4)3 + 3H2↑.
The reaction with dilute H2SO4 is very slow probably on account of the
insolubility of the oxide film in this acid.
Hot concentrated sulfuric acid dissolves Al with the evolving of SO2:
2Al + 6H2SO4 → Al2(SO4)3 + 3SO2↑ + 6H2O
Dilute and concentrated HNO3 has no effect on Al, i.e., Al is rendered passive
by nitric acid due to surface oxidation and formation of a thin film of oxide on its
surface.
Aluminium is attacked by caustic alkalies with the evolving of hydrogen:
2Al + 6NaOH + 3H2O → 2Na3[Al(OH)6] + 3H2↑;
solution
2Al + 6NaOH
Sodium hexahydroxoaluminate (III)

→ 2Na3AlO3 + 3H2↑.
fused
Sodium aluminate
It is a good reducing agent and reduces oxides of metals like Cr, Fe, Mn, etc.:
Cr2O3 + 2Al → 2Cr + Al2O3 + Q;
Fe2O3 + 2Al → 2Fe + Al2O3 + Q;
3Mn3O4 + 8Al → 9Mn + 4Al2O3 + Q.
It reduces oxides of non-metals also:
3CO2 + 4Al → 2Al2O3 + 3C;
3SiO2 + 4Al → 2Al2O3 + 3Si.
14
Being more electronegative it displaces copper, zinc and lead from the
solutions of their salts:
3ZnSO4 + 2Al → Al2(SO4)3 + 3Zn;
3CuSO4 + 2Al → Al2(SO4)3 + 3Cu.
Aluminium oxide Al2O3 is a white amorphous powder, usually soluble in acids
but when it’s ignited above 850 0C, it becomes dense and insoluble in acids.
Pure Al2O3 is prepared by igniting Al(OH)3, Al2(SO4)3 or ammonium alum,
(NH4)Al(SO4)2×12H2O:
2Al(OH)3
∆
→
Al2O3 + 3H2O;
∆
Al2(SO4)3 →
Al2O3 + 3SO3;
∆
2(NH4)Al(SO4)2×12H2O →
Al2O3 + 2NH3 + 4H2SO4 + 9H2O.
It is amphoteric oxide. Ignited Al2O3 is brought into solution as aluminate by
fusion with alkali and carbonates or as sulfate by fusion with KHSO4:
t
Al2O3 + 3Na2CO3 
→
2Na3AlO3 + 3CO2↑;
t

→
Ca(AlO2)2 + CO2↑;
t
Al2O3 + 6KHSO4 
→ 3K2SO4 + Al2(SO4)3 + 3H2O.
Al2O3 + CaCO3
When ammonium hydroxide is added to any soluble aluminium salt, we get a
gelatinous white precipitate of aluminium hydroxide:
Al2(SO4)3 + 6NH4OH → 2Al(OH)3↓ + 3(NH4)2SO4.
Aluminium hydroxide is amphoteric and dissolves both in alkalies and acids:
Al(OH)3 + 3NaOH → Na3[Al(OH)6];
Al(OH)3 + 3HCl → AlCl3 + 3H2O.
Aluminium salts which are formed by weak acids (sulfide, carbonate, etc.) are
completely hydrolysed:
Al2(CO3)3 + 3H2O → 2Al(OH)3↓ + 3CO2↑;
Al2S3 + 6H2O → 2Al(OH)3↓ + 3H2S↑;
so when acting with (NH4)2S of an aluminium salt in aqueous solution aluminium
hydroxide Al(OH)3 is formed instead of aluminium sulfide Al2S3:
2AlCl3 + 3(NH4)2S + 6H2O → 2Al(OH)3↓ + 6NH4Cl + 3H2S↑.
Uses
The production of aluminium consumes over 90 % of the world's production of
bauxite. The remainder is used by the abrasive, refractory, and chemical industries.
Bauxite is used also in the production of high-alumina cement, as an absorbent or
catalyst in the oil industry, in welding rod coatings and fluxes, and as a flux in steel
and ferroalloys production.
Usage of aluminium include the following electrical equipment and industrial
applications where a strong, light, easily constructed material is needed: car, ship,
aircraft construction; metallurgical and chemical processes; domestic and industrial
construction; packaging (aluminium foil, cans); kitchen utensils (cutlery, pans),
15
outside building decoration. Although its electrical conductivity is only about 60%
that of copper per area of cross section, it is used in electrical transmission lines
owing to its lightness and price. Alloys of aluminium are of vital importance in the
construction of modern aircraft and rockets
Aluminium, evaporated in a vacuum, forms a highly reflective coating for both
visible light and radiant heat. These coatings form a thin layer of the protective
oxide and do not deteriorate as silver coatings do. These coatings are used for
telescope mirrors, decorative paper, packages, toys, and in many other manners.
The aluminium oxide, alumina, occurs naturally as ruby, sapphire, corundum, and
emery, and is used in glass making and refractories. Synthetic ruby and sapphire
are used in the construction of lasers.
Amorphous boron is used in pyrotechnic flares (distinctive green colour), and
rockets (as an igniter). Boric acid is used as a mild antiseptic. Borax,
Na2B4O7.10H2O, is a cleansing flux in welding and a water softener in washing
powders. Boron compounds are used in production of enamels for covering steel of
refrigerators, washing machines, etc. They are extensively used in the manufacture
of enamels and borosilicate glasses. Boron nitride is as hard as diamond. It behaves
like an electrical insulator, but conducts heat like a metal. It also has lubricating
properties similar to graphite. The hydrides are sometimes used as rocket fuels.
Boron filaments, a high-strength, lightweight material, are used for advanced
aerospace structures. 10B is used as a control isotope for nuclear reactors, as a
shield for nuclear radiation, and in instruments used for detecting neutrons.
Biological role
Aluminium may be involved in the action of enzymes such as succinic
dehydrogenase and δ-aminolevulinate dehydrase (involved in porphyrin synthesis).
Aluminium compounds are toxic to most plants and somewhat toxic to mammals.
Aluminium has been linked to Alzheimer's disease (senile dementia).
Boron is probably not required in the diet of humans but it might be a necessary
"ultratrace" element. Boron is required by green algae and higher plants.
Daily dietary intake: 2.45 mg of Aluminium and 1-3 mg of Boron. Total mass
of element in average (70 kg) person: 60 mg of Aluminium and 18 mg of Boron.
2. Group IVA. Carbon, Silicon, Germanium, Tin and Lead
In group IVA of the periodic table there are five chemical elements: carbon C,
silicon Si, germanium Ge, tin Sn and lead Pb.
The atoms of these elements have the following configuration:
– 1s22s22p2
[He]2s22p2
6C
2 2
6 2
2
– 1s 2s 2p 3s 3p
[Ne]3s23p2
14Si
2 2
6 2
6
10 2
2
– 1s 2s 2p 3s 3p 3d 4s 4p
[Ar] 4s24p2
32Ge
2 2
6 2
6
10 2
6
10 2
2
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p
[Kr] 5s25p2
50Sn
2 2
6 2
6
10 2
6
10 14 2
6
10 2
2
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s 6p
[Xe]6s26p2
82Pb
Table 2.1 lists some physical properties of group IVA elements.
16
Table 2.1 Properties of Group IVA Elements
Property
Carbon
Silicon
Germanium
Electronic
configuration
[He]2s22p2
[Ne]3s23p2
[Ar]4s24p2
1415
937
232
327
2900
2830
2620
1745
2.33
5.36
7.31
11.34
2.0
1.5
1.6
1.7
1.8
0.77
–
–
1.77
–
0.34
1.22
0.65
0.44
1.40
1.02
0.67
–
1.26
0.76
Melting point, °C
Boiling point, °C
Density, g/cm3
Electronegativity
(Pauling scale)
Covalent radius, Å
Ionic radius E2+, Å
Ionic radius E4+, Å
3500
(diamond)
–
diamond
3.514
graphite
2.255
Tin
Lead
[Kr]5s25p2 [Xe]6s26p2
Physical Properties and Occurrence
In IVA group in the middle of the periodic table there is a considerably greater
change in properties from top to bottom than there is, for example, in the alkali
metals group on the left and the halogen group on the right.
Carbon is a typical nonmetal, silicon and germanium are semimetals, and tin
and lead are metals. However, the elements of group 4 have one important feature
in common, namely each of them have four valence electrons. In its compounds
carbon almost invariably completes its valence shell by forming four covalent
bonds, which have a tetrahedral arrangement around the carbon atom. Silicon and
germanium also usually form four covalent bonds with a tetrahedral arrangement,
whereas tin and lead, which are larger and have smaller ionization energies than
2+
carbon and silicon, often, lose just two of their valence electrons to form the Sn
2+
and Pb ions.
Silicon, symbol Si, is a semimetallic element that is the second most common
element on earth, after oxygen. The atomic number of silicon is 14. It was firstly
isolated from its compounds in 1823 by the Swedish chemist Baron Jцns Jakob
Berzelius.
Crystalline silicon has a hardness of 7, compared to 5 to 7 for glass. Silicon
melts at about 1410 °C, boils at about 2900 °C, and has a specific gravity of 2.33.
The atomic weight of silicon is 28.086.
Silicon constitutes about 28 percent of the earth's crust. It does not occur in the
free elemental state, but it is found in the forms of silicon dioxide and complex
silicates. Silicon-containing minerals constitute nearly 40 percent of all common
minerals, including more than 90 percent of igneous-rock-forming minerals. The
mineral quartz (composed of silicon dioxide, or silica, SiO2), varieties of quartz
17
(such as carnelian, chrysoprase, onyx, flint, and jasper), and the minerals
cristobalite and tridymite are the naturally occurring crystal forms of silica. Silicon
dioxide is the principal constituent of sand. The silicates (such as the complex
aluminum, calcium, and magnesium silicates) are the main constituents of clays,
soils, and rocks in the form of feldspars, amphiboles, pyroxenes, micas, and
zeolites, and of semiprecious stones, such as olivine, garnet, zircon, topaz, and
tourmaline.
Germanium, symbol Ge, is a hard, brittle, grayish-white, crystalline
semimetallic element. The atomic number of germanium is 32. Germanium melts
at about 937 °C, boils at about 2830 °C, and has a specific gravity of 5.3; its atomic
weight is 72.59.
The Russian chemist Dmitry Mendeleyev predicted the existence and chemical
properties of germanium in 1871; he called it ekasilicon because of its position
under silicon in the periodic table. The element was actually discovered in the
silver-sulfide ore argyrodite by the German chemist Clemens Alexander Winkler in
1886.
Germanium ranks 54th in order of abundance of the elements in the earth's
crust. Germanium occurs in small quantities in the ores of silver, copper, and zinc,
and in the mineral germanite, which contains 8 percent of germanium.
Tin, symbol Sn, is a metallic element which has been used by people since
ancient times. The atomic number of tin is 50. Ordinarily a silver-white metal, at
temperatures below 13 °C it often changes into an allotropic (distinctly different)
form known as gray tin, which is an amorphous, grayish powder with a specific
gravity of 5.75. Tin melts at about 232 °C, boils at about 2260 °C, and has a
specific gravity of 7.31. The atomic weight of tin is 118.69.
Tin has been found in the tombs of ancient Egyptians and was exported to
Europe in large quantities from Cornwall, England, during the Roman period. The
ancient Egyptians considered tin and lead different forms of the same metal.
Tin ranks 49th in abundance of the elements in the earth's crust. The principal
ore of tin is the mineral cassiterite (or tinstone), SnO2, found abundantly in
Cornwall, England, and in Germany, the Malay Peninsula, Bolivia, Brazil, and
Australia.
Lead, symbol Pb (Latin plumbum, a lead weight), is a dense, bluish-gray
metallic element which was the one of the first known metals. The atomic number
of lead is 82. Lead melts at 327 °C, boils at 1745 °C, and has a specific gravity of
11.34; the atomic weight of lead is 207.20.
Lead was mentioned in the Old Testament. It was used by the Romans for
making water pipes, soldered with an alloy of lead and tin.
Lead is widely distributed all over the world in the form of its sulfide, the ore
galena (PbS). Lead ranks about 36th in natural abundance among elements in the
earth's crust. Ores of the secondary importance are cerussite and anglesite.
18
Carbon and its Compounds
Although living matter is composed of an enormous variety of carbon
compounds, carbon is only the fourteenth most abundant element. It composes jnly
0.08 % of the earth's crust, and about half of this amount is in the form of the
carbonate ion, CO32–. The most common metal carbonates are calcium carbonate,
CaCO3, in its various forms such as chalk, limestone, and marble; magnesium
carbonate, MgCO3; and dolomite, CaMg(CO3)2. The remaining amount of carbon
is present in vegetable and animal matters, in coal and petroleum, and as carbon
dioxide in the atmosphere and the oceans.
There are an enormous number of carbon compounds, and the vast majority of
them are classified as organic compounds. At one time it was believed that many
compounds of carbon could only be obtained from organic, or living, matter; hence
these compounds were called organic compounds. Compounds occurring in the
nonliving or mineral world were called inorganic compounds. Today we know
that carbon compounds found in living matter can be synthesized from substances
that are not found in living matter. Nevertheless, we retain the name organic to
denote the compounds of carbon with the exception of CO2, carbonates, and a few
other compounds traditionally regarded as inorganic compounds. With a very few
exceptions we may define organic chemistry as the chemistry of carbon
compounds and inorganic chemistry as the chemistry of all the other elements and
their compounds. The division between the inorganic and organic compounds of
carbon is, however, an artificial one. In the process of photosynthesis carbon
dioxide and water are converted by plants to organic compounds. We convert
organic compounds in our bodies back to carbon dioxide and water. Indeed, life
itself appears to have originated from simple inorganic compounds such as H2,
H2O, NH3, and HCN, as well as CH4, which were present in the original
atmosphere of our planet. Probably these compounds were combined under the
action of lightning discharges or ultraviolet radiation to form the amino acids and
other organic compounds that form the basic of life.
Physical Properties of Carbon
Carbon occurs in several different solid forms, the most important of which are
diamond and graphite. The different forms of an element are called allotropes.
Thus diamond and graphite are allotropes of carbon.
19
Diamond and graphite are so different that if we did not know that they both
consist of carbon, and carbon only, we would find it difficult to believe that they
are forms of the same element. Pure diamond forms beautiful, transparent,
colorless, very hard, and highly refractive crystals. In contrast, graphite is a soft,
black substance which is used as a lubricant, and as the "lead" in lead pencils when
mixed with varying amounts of clay. Diamond has a much higher density (3.53
g/cm3) than graphite (2.25 g/cm3) and is an electrical insulator, whereas graphite is
a fairly good conductor. Being a fairly good conductor of electricity graphite is
exceptional among the nonmetals.
Carbon black is a pure form of soot which is deposited when hydrocarbons are
burned in a very limited supply of air. For example,
2C2H2 + O2 → 4C + 2H2O.
Ethyne
Carbon black has a very intense color and is used in large quantities as a pigment for paint, paper, and printer's ink, and to reinforce and color the rubber used
in automobile tires.
Charcoal is made by heating wood and other organic materials to a high
temperature in the absence of air. We are familiar with the fact that charcoal is
much lighter than the wood from which it is made; in other words, it appears to
have a very low density. The density is low because charcoal is extremely porous;
it has a structure resembling a sponge but with holes that are too small to be visible
to the eye. This porous structure means that it has a very large surface area relative
to its volume. This large surface area can adsorb considerable quantities of other
substances. Charcoal that has been thoroughly cleaned by heating with steam is
known as activated charcoal. It has many applications, such as removing unburnt
hydrocarbons from automobile exhaust, unpleasant and dangerous gases from the
air, and impurities from water. Many municipal water treatment plants pass water
through beds of activated charcoal.
Coke is made by heating coal in the absence of air. Coal is a very complex
material consisting of many organic compounds. It contains 60%–90% C together
with H, O, N, S, Al, Si, and some other elements. When coal is heated to a high
temperature, it decomposes, producing a variety of gaseous and liquid products.
The mixture of gases, methane and hydrogen mainly, is known as coal gas. The
mixture of liquid products, which includes many hydrocarbons and other organic
compounds, is called coal tar. The solid residue contains 90%–98% C and is
known as coke. Coke is used in enormous quantities in industry as a reducing agent
for the production of metals, phosphorus, and other substances.
In 1985, scientists vaporized graphite to produce a stable form of carbon
molecule consisting of 60 carbon atoms (C60) in a roughly spherical shape, looking
like a soccer ball. The molecule was named buckminsterfullerene – “buckyball” for
short – in honor of R. Buckminster Fuller, the inventor of the geodesic dome. The
molecule may be common in interstellar dust.
20
Chemical Properties of Carbon and its Compounds
Carbon is rather chemically passive while its reactivity increasing with the
increasing of temperature. It reacts vigorously with oxygen when heated in the air.
In the lack of oxygen carbon monoxide forms while carbon dioxide is formed in
the excess of oxygen:
2C + O2 → 2CO↑;
C + O2 → CO2↑.
Carbon reacts with some metals giving the metals carbides under the high
temperatures:
2C + Ca → CaC2.
Carbon combines with hydrogen at the temperature of electric arc producing
methane:
C + 2H2 → CH4↑.
When water steam is passing through red heated charcoal or coke (at the
temperature of 1200 °C) the mixture of carbon monoxide and hydrogen gasses
forms:
C + H2O → CO↑ + H2↑.
There are other products of this reaction at low temperatures (1000 °C):
C + 2H2O → CO2↑ + 2H2↑.
These reactions make the basis of conversion method of hydrogen preparation.
Concentrated sulfuric and nitric acids oxidize carbon to carbon dioxide when
heated:
C + 2H2SO4 → CO2↑ + 2SO2↑ + 2H2O;
3C + 4HNO3 → 3CO2↑ + 4NO↑+ 2H2O.
CARBON MONOXIDE
When carbon is burned in a limited supply of air, carbon monoxide, CO is
obtained:
2C + O2 → 2CO.
Industrially, carbon monoxide mixed with hydrogen is made on a large scale by
passing water steam over red-hot coke:
C + H2O → CO + H2.
This mixture of hydrogen and carbon monoxide is called water gas.
Another increasingly important method of carbon monoxide preparation is the
high-temperature reaction of methane (natural gas) with steam in the presence of a
catalyst:
CH4 + H2O Catalyst

→ CO + 3H2.
The mixture of carbon monoxide and hydrogen prepared in this way is known as
synthesis gas, because of its importance as the basic material for synthesis of many
organic compounds. It is also used as a fuel.
Carbon monoxide is a colorless, odorless, and tasteless gas that has only a
slight solubility in water. It is very toxic because it combines with hemoglobin in
the blood, thus preventing hemoglobin from carrying out its function as an oxygen
21
carrier. It is particularly dangerous because it is odorless and therefore not easily
detected. Carbon monoxide is produced when tobacco burns in cigarettes and it is
present in the exhaust gases from automobiles.
Carbon monoxide is a reducing agent which can be oxidized to carbon dioxide.
It burns in the air forming carbon dioxide:
2CO + O2 → 2CO2.
It reduces steam at high temperature, giving an equilibrium mixture of CO2 and
H2:
CO + H2O ⇄ CO2 + H2.
Carbon monoxide reduces metal oxides to metals. For example,
∆ 2Fe + 3CO ↑.
Fe O + 3CO →
2
3
2
CARBON DIOXIDE AND CARBONIC ACID
The combustion of carbon and carbon-containing compounds in the excess of
oxygen leads to the formation of carbon dioxide, CO2. It is produced also by
heating alkaline-earth metal carbonates, such as calcium carbonate:
∆ CaO + CO ↑.
CaCO →
3
2
Beer and wine making also produce carbon dioxide as a by-product because the
fermentation of a sugar such as glucose gives ethanol and carbon dioxide:
C6H12O6 fermentati
  on
→ 2C2H5OH + 2CO2↑.
In the laboratory we can prepare small amounts of carbon dioxide by adding a
dilute aqueous acid to a metal carbonate:
CaCO3 + 2H+ → Ca2+ + CO2↑ + H2O.
Carbon dioxide is a linear AX2 molecule with two double bonds. It is a colorless
gas with a very slight odor. When cooled, it forms a white solid (dry ice) that
sublimes at –78 0C at 1 atm pressure. Liquid carbon dioxide can be obtained only
under pressure.
Carbon dioxide dissolves readily in water giving a solution that contains a
small equilibrium amount of carbonic acid, H2CO3:
CO2 + H2O ⇄ H2CO3.
This solution is commonly called soda water; it has a slightly acidic taste.
Carbonic acid is a weak acid that is partially ionized in two stages:
H2CO3 + H2O ⇄ H3O+ + HCO3–;
HCO3– + H2O ⇄ H3O+ + CO32–.
Carbonic acid and its two ions, HCO3– and CO32–, are represented by the
following structures:
O C
O C
O
O
O H
O H
O C
O
H
O
Carbonate ion, CO32–, and hydrogen carbonate ion, HCO3–, being the anions of
a weak acid, H2CO3, are weak bases. In the presence of an acid CO32– adds a
proton forming its conjugate acid HCO3–:
22
And in turn
HCO3–
CO32– + H+ ⇄ HCO3–.
adds a proton giving its conjugate acid H2CO3:
HCO3– + H+ ⇄ H2CO3.
But H2CO3 is not very stable and it almost completely decomposes to CO2 and
water:
H2CO3 ⇄ H2O + CO2.
Carbon dioxide is used as a fire extinguisher because it does not burn or easily
support combustion of other substances and because it has a higher density than
air. It sinks down on the fire, forming a blanket of carbon dioxide that excludes air,
extinguishing the fire. Only a few very reactive metals such as sodium, potassium,
and magnesium will burn in carbon dioxide. For example, a previously ignited
piece of magnesium will burn in carbon dioxide with a spluttering flame producing
white magnesium oxide and black carbon particles:
CO2 + 2Mg → 2MgO + C.
CARBON DISULFIDE
Carbon combines with sulfur at heating and gives carbon disulfide, CS2:
∆ CS .
C + 2S →
2
Carbon disulfide is a linear AX2 molecule with double bonds, just like carbon
dioxide:
:S=C=S:
Carbon disulfide is a toxic, flammable liquid, but it is a useful solvent for sulfur
and rubber.
TETRACHLOROMETHANE (CARBON TETRACHLORIDE)
When carbon disulfide is heated with chlorine, tetrachloromethane, CCl4, and
disulfur dichloride, S2Cl2, are formed:
∆ CCl + S Cl .
CS + 3Cl →
2
2
4
2
2
The two products can be separated by distillation. Tetrachloromethane is a
useful solvent. Because it is a good solvent for oils and grease, it may be used as a
cleaning agent. But it should be used with suitable precautions because it is toxic
and the liquid can pass through the skin.
HYDROGEN CYANIDE
When heated to approximately 1000–1200 °C in the presence of a catalyst, a
mixture of methane and ammonia burns in the air and gives hydrogen cyanide in an
exothermic reaction:
∆ 2HCN + 6H O.
2CH + 2NH + 3O →
4
3
2
2
Ammonia and methane react also in the absence of air. The reaction is
endothermic and requires a temperature of 1200–1300 °C and a platinum catalyst:
∆ → HCN + 3H .
CH + NH Catalyst,


4
3
2
23
Hydrogen cyanide is a colorless liquid that boils just above room temperature
(25.6 °C). It has the odor and taste of bitter almonds and is highly toxic. In the
form of a gas it is used in execution chambers; a concentration of only
0.2 %
by volume causes death within minutes.
HCN can be conveniently prepared in the laboratory by adding acid to a metal
cyanide such as NaCN or KCN:
NaCN + HCl → NaCl + HCN.
HCN is a weak acid and therefore the position of the equilibrium lies far to the
right:
CN– + H3O+ ⇄ HCN + H2O.
Because of its low boiling point HCN is evolved as a gas. For this reason acids
should not be added to metal cyanide except in a fume hood. The HCN molecule
has a triple C≡N bond. It has a linear geometry as expected for an AX2 molecule.
Hydrogen cyanide behaves as a weak acid in water. An aqueous solution is
called hydrocyanic acid. Its salts are the cyanides, such as sodium cyanide, NaCN.
−
They contain the cyanide ion : C ≡ N : .
Hydrogen cyanide has many important applications in the plastics industry.
Sodium cyanide is used for the extraction of gold and silver from their ores:
4Au + 8NaCN + O2 + 2H2O → 4Na[Au(CN)2] + 4NaOH .
When an alkali metal cyanide is fused with sulfur, a thiocyanate is formed, for
example:
∆ NaSCN.
S + NaCN →
The thiocyanate ion, SCN–, has the following structure:
−
..
: S− C ≡ N : .
..
CARBIDES
In addition to its compounds with nonmetals such as oxygen, sulfur, and
nitrogen, carbon reacts with many metals. The compounds of carbon with the alkali
and alkaline earth metals contain the carbide ion, C22–, for example, Na2C2 and
CaC2.
Commercially, calcium carbide is made by heating lime with coke in a furnace:
∆ CaC + CO↑.
CaO + 3C →
2
Calcium carbide reacts with water giving ethyne (acetylene), C2H2:
CaC2 + 2H2O → Ca(OH)2 + C2H2↑.
These carbides are salts of ethyne, which are too weak as acid to ionize in
water. Therefore, the carbide ion, C22–, is a strong base in water. It removes protons
from water and gives C2H2, which bubbles off as a gas:
C22– + 2H2O → C2H2↑ + 2OH–.
There are many other types of metal carbides that do not contain the C22– ion,
for example Al4C3:
Al4C3 + 12H2O → 4Al(OH)3 + 3CH4.
Often they have unexpected formulas, such as iron carbide, Fe3C:
24
Fe3C + 6HCl → 3FeCl2 + CH4 + H2,
which is called cementite and is an important component of some steels.
With more electronegative elements such as silicon, carbon forms covalent
rather than ionic carbide. Silicon carbide, SiC, is a covalent compound with the
diamond structure, except that each alternate atom is a silicon atom rather than a
carbon atom. Silicon carbide, commonly known as carborundum, is almost as hard
as diamond and is used as an abrasive; for example, carbide sandpaper and
grinding wheels are coated with silicon carbide.
Silicon and its Compounds
Silicon follows carbon in group IVA. Whereas carbon is a typical nonmetal,
silicon and the next element germanium, are semimetals, and they are followed by
tin and lead, which are metals. Unlike the elements in the groups on the left and
right sides of the periodic table, those in the main groups in the middle of the table
show a considerable variation in properties on descending the group. Silicon is a
shiny, silvery solid that looks like a metal but has only a low electrical
conductivity; moreover, its conductivity increases with increasing of temperature,
whereas the conductivity of a metal decreases with increasing of temperature.
Substances that have small electrical conductivity in the solid state which increases
appreciably with increasing temperature are known as semiconductors.
Silicon occurs not only as silicates but also as silicon dioxide, SiO2, which has
been known for centuries as silica, which is familiar, in an impure form, as sand.
Preparation of Silicon
The element can be prepared from silica by it heating with coke to a
temperature of about 3000 °C in an electric arc furnace:
SiO2 + 2C → Si + 2CO↑.
The reactants are added continuously at the top of the furnace. Carbon monoxide escapes from the furnace and burns giving carbon dioxide, while the molten
silicon (Tm= 1414 °C) runs out from the bottom of the furnace and solidifies. This
silicon is pure enough for many purposes, such as the manufacture of alloys with
metals, but ultra pure silicon, needed in many electronic devices is obtained by
preliminary converting impure silicon into silicon tetachloride by heating it with
chlorine:
Si + 2Cl2 → SiCl4 (Tb= 57.6 °C) .
The obtained SiCl2 is then purified by distillation and reduced to silicon by heating
with hydrogen or magnesium:
SiCl2 + H2 → Si + 2HCl ;
SiCl4 + 2Mg → Si + 2MgCl2 .
The magnesium chloride is removed from the silicon by washing it out with hot
water. The silicon can be further purified by zone refining.
25
Chemical Properties of Silicon and its Compounds
Silicon is a rather uncreative element and is not attacked by acids. Amorphous
silicon burns brilliantly in oxygen when heated to 450 °C. It reacts with fluorine at
400 °C, with chlorine at 450 °C and with bromine at 500 °C. It reacts at red heat
with iodine:
Si + 2F2
400 C

→ SiF4;
o
450 C

→ SiCl4;
500 o C
Si + 2Br2 → SiBr4.
Si + 2Cl2
o
Silicon decomposes water at red heat:
Si + 2H2O → SiO2 + 2H2↑.
It reacts with hot, concentrated aqueous hydroxide solutions and with molten
hydroxides forming silicates and evolving hydrogen:
Si + 2NaOH + H2O → Na2SiO3 + 2H2.
Silicon does not react with oxy-acids, it may be dissolved in the mixture of
nitric and hydrofluoric acids:
3Si + 4HNO3 + 18HF → 3H2SiF6 + 4NO↑ + 8H2O.
SILICON DIOXIDE
Silicon dioxide (SiO2), or silica, occurs in several crystalline forms, including
quartz, cristobalite, and tridymite, and in several amorphous forms, such as agate
and flint. Quartz is the best known of the silica minerals. It is one of the commonest minerals in the earth's crust and is often found in the form of colorless,
transparent crystals, which are sometimes beautifully formed and of enormous size.
Quartz is a constituent of many rocks, which are often complex mixtures of
different minerals. Granite, for example, is a mixture of the three silicon minerals,
quartz, mica, and feldspar. Some forms of quartz that contain traces of impurities
are beautifully colored and are often used as semiprecious gemstones. Amethyst,
for example, is quartz that is colored violet by traces of Fe(III). Onyx, jasper,
carnelian, and flint are colored forms of noncrystalline silicon dioxide.
Silicon dioxide is an acidic oxide that reacts with NaOH or Na2CO3 on heating
giving silicates:
SiO2 + 4NaOH → Na4SiO4 + 2H2O;
SiO2 + Na2CO3 → Na2SiO3 + CO2↑.
These are only simplified representations of the reactions; there are many silicates with complex structures, as we will see in the next section.
A concentrated solution of sodium silicate in water is called wafer glass. It is
used for fireproofing wood and cloth, for adhesives, and for preserving eggs. It can
be used to prepare other metal silicates.
SILICIC ACID AND SILICATES
26
Silicon dioxide is the anhydride of silicic acid. When finely powdered silicon
dioxide is shaken with water for a long time, a very slightly acidic solution is
obtained, due to the formation of a very small amount of silicic acid, Si(OH)4:
SiO2 + 2H2O → Si(OH)4 or
SiO2 + 2H2O → H2SiO3 .
The very small solubility of silicon dioxide in water is increased at high
pressure and temperature, which accounts for the deposits of silica found around
many hot geysers. The silicon dioxide which is dissolved in water at high pressure
below the surface precipitates from the solution when the water rises to the surface,
where the pressure and temperature are much lower.
More concentrated solutions of silicic acid can be obtained by the reaction of
silicon tetrachloride with water:
SiCl4 + 4H2O → Si(OH)4 + 4HCl
or by the reaction of an aqueous solution of sodium silicate with hydrochloric acid:
Na2SiO3 + H2O + 2HCl → Si(OH)4 + 2NaCl.
However, the product of these reactions is not simply a solution of Si(OH)4. A
gelatinous solid is obtained that consists of polymeric silicic acids formed by
condensation reactions such as:
Continuation of such reactions leads to a complex mixture of many polymeric
acids.
Silicic acid is very weak (Ka = 1×10–10). It is a member of very weak acids
class such as HOCI and B(OH)3 (or H3BO3), it has the general formula X(OH)n. It
is the first member of the third-period oxoacids series. These acids increase in
strength from the very weak silicic acid to the very strong perchloric acid.
If the mixture of polymeric silicic acids is heated, causing further condensation
reactions to occur, a hard, granular, translucent substance called silica gel is
obtained. It has a large surface area and readily adsorbs water and other substances.
It is widely used as a drying agent and as a catalyst. Small bags of silica gel are
frequently packed with delicate scientific apparatus, cameras, and electronic
equipment to protect them from damage by moisture during transportation.
SILICATES
Over one thousand silicates occur naturally. Their number and complexity result from the great variety of ways in which SiO4 tetrahedral can be linked
together. The simplest silicates contain the anion SiO44–, derived from the acid
Si(OH)4. Olivine is an important mineral of this type; it is the principal component
of the earth's mantle, which lies between the core and the crust. Olivine is an iron
magnesium silicate that is often represented by the formula FeMgSiO4, although its
composition may vary from Mg2SiO4 to Fe2SiO4. The SiO44– tetrahedrals are
27
packed together in a compact way in olivine, giving it a greater density compared
with most other silicate minerals.
Many silicates are the salts of the many polymeric forms of silicic acid. The
condensation of two silicic acid molecules gives the acid H6Si2O7:
SILICONE POLYMERS
The wide range of applications of the mineral silicates depends on their great
thermal stability and inertness toward other substances. These properties are
caused by the great strength of the silicon–oxygen bond, which has an average
bond energy (464 kJ/mol) that is much greater than the silicon–silicon bond energy
(196 kJ/mol) or even the carbon–carbon, single-bond energy (348 kJ/mol).
Chemists succeeded in combination of the strength and inertness of the siliconoxygen bond with some of the useful properties of organic polymers in the
synthetic polymers known as silicones.
The simplest silicones are chain polymers that have the general formula
(R2SiO)n, where R is an alkyl or aryl group such as CH3, C2H5, or C6H5. There are
also cyclic polymers (see Figure 2.1) and cross-linked polymers in which chains
are held together by sharing oxygen atoms.
Silicones are prepared by heating silicon with chloroalkanes such as
chloromethane, CH3Cl. The main product of the reaction with CH3Cl is dimethyl
silicon dichloride, (CH3)2SiCI2:
But small amounts of (CH3)3SiCl and CH3SiCl3 are also formed.
Figure 2.1. Structures of Some Silicone Polymers.
(a) a long-chain polymer; (b) a cyclic polymer.
28
SILANES
The great difference between the chemistry of silicon and the chemistry of
carbon is also illustrated by the hydrides of silicon, which are called silanes. In
their general formula SinH2n+2, the silanes are analogous to the alkanes, CnH2n+2,
but only a few silanes are known. They are extremely reactive and are spontaneously flammable in the air:
SiH4 + 2O2 → SiO2 + 2H2O.
In contrast, the alkanes burn only when ignited. The silanes react rapidly with
basic aqueous solutions evolving hydrogen:
SiH4 + OH– + 3H2O → SiO(OH)3– + 4H2↑ .
SILICON HALIDES
Silicon tetrachloride, which is a colorless liquid (Tb= 57 °C), can be made by
passing chlorine through a red-hot mixture of sand and coke:
SiO2 + 2C + 2Cl2 → SiCl4 + 2CO↑ .
Silicon tetrafluoride, which is a colorless gas, can be made by the reaction of
silicon with fluorine,
Si + 2F2 → SiF4
or by heating calcium fluoride with sand and concentrated sulfuric acid:
2CaF2 + SiO2 + 2H2SO4(conc.) → SiF4↑ + 2CaSO4 + 2H2O.
In this reaction hydrogen fluoride is produced by the reaction of calcium fluoride
with concentrated sulfuric acid, and then HF reacts with silicon dioxide:
CaF2 + H2SO4(conc.) →2HF + CaSO4;
SiO2 + 4HF → SiF4↑ + 2H2O .
Both SiCl4 and SiF4 are covalent tetrahedral AX4 molecules. In contrast to the
carbon compounds CF4 and CCl4, which are stable in the presence of water, both
SiCl4 and SiF4 react rapidly with water giving silicic acid, Si(OH)4, and a hydrogen
halide:
SiCl4 + 4H2O → Si(OH)4 + 4HCl ;
SiF4 + 4H2O → Si(OH)4 + 4HF .
Compounds of Tin
Tin shows allotropy among its three allotropic forms:
Grey tin
White tin
Rhombic tin
18 o C
170 o C
←
→
←

→
Specific gravity = 5.8
Specific gravity = 7.8
Specific gravity = 7.56
The common oxidation states of tin are +2 and +4, but Sn+4 compounds are
a little more stable.
When tin metal is heated to whiteness (1500-1600°C) in presence of O2 tin
burns with a bright flame and forms SnO2.
29
When Sn is heated in an atmosphere of C12 or with sulfur, SnCl4 and SnS2 are
formed:
Sn + 2Cl2 → SnCl4;
Sn + 2S → SnS2.
Sn is slowly attacked by dilute HCl and dilute H2SO4 and H2 is evolved in both
the cases. With concentrated HCl, H2 is evolved and with concentrated H2SO4, SO2
gas is obtained:
Sn + 2HCl → SnCl2 + H2↑;
Sn + H2SO4(dil.) → SnSO4 + H2↑;
Sn + 4H2SO4(conc.) → Sn(SO4)2 + 2SO2↑ + 4H2O.
When very dilute HNO3 (6%) reacts with metallic tin, the metal is oxidized to
Sn(NO3)2 but HNO3 reduces to NH4NO3:
4Sn + 10HNO3 → 4Sn(NO3)2 + NH4NO3 + 3H2O.
When hot and concentrated HNO3 reacts with the metal, it is oxidized to
stannic acid (H2SnO3) while the acid gets reduced to NO2:
Sn + 4HNO3 → H2SnO3 + 4NО2↑ + H2O.
When boiling with concentrated alkali solution, the metal dissolves, forming
solution of stannate:
Sn + 2KOH → K2SnO3 + 4H2↑.
DIHYDRATED STANNOUS CHLORIDE, SnCl2⋅2H2O
It is prepared by dissolving Sn in hot and concentrated HCl:
Sn + 2HCl → SnCl2 + H2↑.
The solution after evaporation and cooling gives the transparent monoclinic
crystals of the SnCl2⋅2H2О which is known as tin salt.
Hydrated stannous chloride forms transparent monoclinic crystals which melt at
40 °C. At heating SnCl2⋅2H2O loses HCl and gives stannous oxochloride,
Sn(OH)Cl. Hence anhydrous salt cannot be prepared by heating SnCl2⋅2H2O.
∆
→ Sn(OH)Cl↓ + HCl.
SnCl2 + H2O 
When treated with alkalies, a white precipitate of Sn(OH)2 is obtained. This
precipitate dissolves in the excess of the alkali and forms corresponding stannite.
Na2SnO2 absorbs atmospheric O2 forming sodium stannate, Na2SnO3:
SnCl2 + 2NaOH → Sn(OH)2↓ + 2NaCl;
Sn(OH)2 + 2NaOH → Na2SnO2 + 2H2O;
2Na2SnO2 + O2 → 2Na2SnO3.
With H2S a dark brown precipitate of SnS is obtained. This precipitate is
soluble in yellow ammonium sulfide (NH4)2S forming ammonium thiostannate,
(NH4)2SnS3:
SnCl2 + H2S → SnS↓ + 2HCl;
SnS + (NH4)2S2 → (NH4)2SnS3.
Stannous chloride is an active reducing agent and reduces a number of
compounds:
2FeCl3 + SnCl2 → 2FeCl2 + SnCl4;
30
2HgCl2 + SnCl2 → Hg2Cl2↓ + SnCl4;
Hg2Cl2 + SnCl2 → 2Hg↓ + SnCl4;
2CuCl2 + SnCl2 → 2CuCl↓ + SnCl4.
ANHYDROUS STANNOUS CHLORIDE, SnCl2
It is prepared by heating tin in a steam of dry HCl gas or by heating a mixture
of metallic tin in excess with HgCl2 when Hg volatilizes and SnCl2 remains:
Sn + 2Cl2 → SnCl4;
Sn + 2HgCl2 (excess) → SnCl4 + 2Hg.
Anhydrous SnCl2 is a glass-like transparent substance with the density of 3.95.
It melts at 246 °C and boils at 603 °C, the vapour being associated with Sn2Cl4:
Sn2Cl4 → 2SnCl2.
It is a covalent compound soluble in organic solvents like alcohol and ether. It
combines with NH3 forming various addition compounds like SnCl2 ⋅2NH3,
SnCl2⋅2NH3 and 3SnCl2⋅2NH3 depending the temperature of the reaction.
STANNIC CHLORIDE, SnCl4
It is prepared by passing dry Cl2 over molten tin kept in a retort, or by heating
tin with the excess of mercury (II) chloride:
Sn + 2Cl2 → SnCl4;
Sn + 2HgCl2 (excess) → SnCl4 + 2Hg.
It is a colourless fuming liquid, density 2.229 (at 20°C), its boiling point is
114.1°C. It is a covalent compound, soluble in organic solvents like C6H6 and of
very negligible electrical conductivity. It is miscible with CS2 in all proportions
and can dissolve phosphorus, sulfur, iodine etc. In small quantity of water SnCl4
dissolves with the evolution of heat and forms a clear solution from which several
soluble hydrates like SnCl4⋅3H2O, SnCl4⋅5H2O, SnCl4⋅6H2O etc., can be
crystallized. With the excess of water, SnCl4 is hydrolyzed and basic stannic
chloride, Sn(OH)Cl3 is precipitated. At further hydrolysis Sn(OH)Cl3 gives a
colloidal solution of stannic acid, H4SnO4:
SnCl4 + H2O ⇄ Sn(OH)Cl3↓ + HCl;
Sn(OH)Cl3 + 3H2O ⇄ H4SnO4 + 3HCl.
With NH3, it gives double salt, SnCl4⋅4NH3 which can be sublimed without
decomposition.
Compounds of Lead
Lead is the element of IV group and has the electron configuration
[Xe]4f145d106s26p2, with a valence shell consisting of the four 6s26p2 electrons. Its
valence-shell electron configuration is similar to that of carbon, 2s22p2, and it
forms a few compounds in which all four electrons are used to form four covalent
bonds. In these compounds, which are predominately covalent and resemble the
corresponding compounds of carbon and silicon, lead is in the +4 oxidation state.
However, we have pointed out previously that elements in the central groups of the
31
periodic table resemble each other less than those in the groups at either end. All
elements in group IV have a common valence of 4, but there is a distinct trend in
the group from the typically nonmetallic element carbon at the top to the typically
metallic element lead at the bottom. Because of the large size of the lead atom, its
valence-shell electrons are much less strongly held than are those of carbon. Thus
lead loses two electrons from its ground state configuration quite readily to form
the Pb2+ ion in which lead is in the + 2 oxidation state:
6s
6p
Pb ground state
Pb2 +
These Pb(II) compounds are typical ionic compounds like those formed by
other metals and they are the most common compounds of lead. The less numerous
Pb(IV) compounds are much less ionic and are best described as covalent.
Lead is readily obtained in crystalline state by the process of precipitation. If a
zinc rod is suspended in a weak solution of lead acetate the metallic lead
precipitates as a mass of carbonaceous crystals, (commonly known as the lead
tree):
(CH3COO)2Pb + Zn → Pb + (CH3COO)2Zn.
Lead is unaffected in dry air but in the presence of moisture it tarnishes due to
the formation of a thin film of the hydroxide and finally carbonate. This layer
protects the metal from further influence of air.
When heated in air or oxygen, it is slowly oxidized to lead monoxide, PbO, and
at a high temperature (400-500°C) to red lead, Pb3O4.
2Pb + O2 → 2PbO;
2PbO + O2 → 2PbO2;
2PbO + PbO2 → Pb3O4.
Lead is not attacked by pure water in the absence of air but water containing
dissolved air has a solvent action on it due to the formation of lead hydroxide:
2Pb + O2 + 2H2O → 2Pb(OH)2.
Dilute hydrochloric and sulfuric acids have practically no action on lead. Hot
concentrated hydrochloric acid forms lead chloride with the evolution of hydrogen:
Pb + 2HCl(hot conc.) → PbCl2 + H2↑.
Hot concentrated sulfuric acid liberates sulfur dioxide but the reaction is
retarded by the formation of an insoluble layer of lead sulfate:
Pb + 2H2SO4 → PbSO4 + SO2↑ + 2H2O.
Dilute and concentrated nitric acid dissolves lead forming lead nitrate and the
oxides of nitrogen:
3Pb + 8HNO3(dil.) → 3Pb(NO3)2 + 2NO↑ + 4H2O;
Pb + 4HNO3(conc.) → Pb(NO3)2 + 2NO2↑ + 2H2O.
Lead combines with chlorine and sulfur on heating:
Pb + S → PbS;
32
Pb + 2Cl2 →PbCl4.
LEAD(II) OXIDE
When lead is heated in air, the yellow oxide PbO is formed.
LEAD(II) CHLORIDE, PbCl2
Lead dissolves slowly in boiling, concentrated hydrochloric acid to form PbCl2:
Pb(s) + 2HCl(aq) → PbCl2(aq) + H2(g) .
Lead(II) chloride is more easily prepared by addition of a soluble chloride to a
solution of a lead(II) salt:
Pb2+(aq) + 2Cl–(aq) → PbCl2(s) .
It is only sparingly soluble in cold water but is more soluble in hot water. When
the hot solution is cooled, white crystals of PbCl2 precipitate.
LEAD(II) IODIDE, PbI2
Lead(II) iodide can be precipitated from a solution of a Pb(II) salt. The
precipitate is somewhat soluble in hot water. When the solution is cooled,
sparkling golden yellow crystals are obtained:
Pb2+(aq) + 2I–(aq) → PbI2(s) .
LEAD(II) HYDROXIDE, Pb(OH)2
When NaOH(aq) is added to a solution of a soluble Pb(II) salt a white
precipitate of Pb(OH)2 is obtained:
Pb2+(aq) + 2OH–(aq) → Pb(OH)2(s).
But this precipitate redissolves when the excess of OH – ions is added because
of the soluble Na2[Pb(OH)4] formation, which contains the [Pb(OH)4]2– ions:
Pb(OH)2(s) + 2OH–(aq) → [Pb(OH)4]2–(aq) .
Lead hydroxide is amphoteric and soluble in aqueous acids:
Pb(OH)2(s) + 2HCl(aq) → PbCl2(aq) + 2H2O(l).
LEAD(II) SULFIDE
When lead is heated with sulfur PbS is obtained. When H2S gas is passed
through an aqueous solution of a lead(II) salt black precipitate of PbS is forming.
LEAD(II) SULFATE AND NITRATE
These salts can be prepared from PbO or Pb(OH)2 and the appropriate acid.
Lead(II) sulfate is not very soluble so it is conveniently prepared by addition of a
soluble sulfate to a solution of a soluble Pb(II) salt such as Pb(NO3)2. Lead(II)
nitrate decomposes at heating and gives nitrogen dioxide, NO2, and oxygen:
2Pb(NO3)2(s) → 2PbO(s) + 4NO2(g) + O2(g) .
This reaction is used often to prepare small amounts of NO2.
LEAD(IV) COMPOUNDS
33
If lead is heated in oxygen, or if PbO is strongly heated in air, it is further
oxidized to Pb3O4, red lead:
6PbO(s) + O2(g) → 2Pb3O4(s).
The oxide Pb3O4 is the basis of a red paint that is widely used for protecting
iron and steel structures against rust; apparently, it forms an oxidized, unreactive
layer on the surface of the iron. Red lead, Pb3O4, contains both Pb(II) and Pb(IV)
and its formula may be written as Pb(II)2Pb(IV)O4.
Red lead reacts with concentrated nitric acid giving soluble lead(II) nitrate and
lead dioxide, Pb(IV)O2, a brown insoluble solid in which lead is in the +4
oxidation state:
Pb3O4 + 4HNO3(conc.) → 2Pb(NO3)2 + PbO2 + 2H2O.
The reaction of Pb3O4 with HNO3 is consistent with the formulation of this
oxide as Pb(II)2Pb(IV)O4, because the Pb(II) atoms convert to soluble Pb(NO3)2,
lead(II) nitrate, leaving the Pb(IV) as the insoluble oxide PbO2.
Lead(IV) chloride (lead tetrachloride), PbCl4, is a yellow liquid that
decomposes when heated forming chlorine and PbCl2:
PbCl4(l) → PbCl2(s) + Cl2(g) .
It is a covalent compound consisting of tetrahedral PbCl4 molecules,
analogous to carbon tetrachloride, CCl4, and silicon tetrachloride, SiCl4.
Uses
Carbon is used as a fuel (coal) and as lubricant (graphite). C-14 isotope is used
in archaeological dating. Carbon compounds are important in many branches of the
chemical industry.
Silicon doped with boron, gallium, phosphorus, or arsenic, etc. is applied for
production of silicon which is used in transistors, solar cells, rectifiers, and other
electronic solid-state devices. Silica, as sand, is a principal ingredient of glass, a
material with excellent mechanical, optical, thermal, and electrical properties.
Silicon is used in manufacturing of computer chips and as a lubricant. It is used
also in medicine for silicone implants.
Tin is used in the preparation of a number of alloys which are very useful. Tin
alloys with lead are used in soldering, in making cups, mugs and other utensils. Its
alloys with arsenic and copper are used for making bearings for machinery as well
as for making crockery and table wares. Alloys of tin with copper, lead, bismuth,
and arsenic are applied for making bells, gongs etc., for electrical fuses, for
printing types. Tin is used for tinning of household utensils and for plating of iron
sheets. Tin amalgam is used for mirrors production.
Lead is used for making cable covering, protective sheets for roofs and drains,
water pipes and for lining the chambers in sulfuric acid manufacturing. It is
34
employed in production of accumulator plates, lead shots, fuse wire and such
compounds as red lead, litharge, lead tetraethyl and white lead. It is an excellent
protection from radiation in atomic work.
Compounds of tin and lead are poisonous.
3. Group VA. Nitrogen, Phosphorus, Arsenic, Antimony
and Bismuth
There are five chemical elements in group VA of the periodic table: nitrogen N,
phosphorus P, arsenic As, antimony Sb and bismuth Bi.
The atoms of these elements have the following configuration:
– 1s22s22p3
[He] 2s22p3
7N
2 2
6 2
3
– 1s 2s 2p 3s 3p
[Ne] 3s23p3
15P
2 2
6 2
6
10 2
3
– 1s 2s 2p 3s 3p 3d 4s 4p
[Ar] 4s24p3
33As
2 2
6 2
6
10 2
6
10 2
3
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p
[Kr] 5s25p3
51Sb
2 2
6 2
6
10 2
6
10 14 2
6
10 2
3
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s 6p
[Xe] 6s26p3
83Bi
Table 3.1 lists some physical properties of group VA elements.
Property
Electronic
configuration
Melting point, °C
Boiling point, °C
Density, g/cm3
Electronegativity
(Pauling scale)
Covalent radius, Å
Ionic radius E3–, Å
Ionic radius E5+, Å
Metallic radius, Å
Table 3.1 Properties of Group VA Elements
Nitrogen Phosphorus Arsenic Antimony Bismuth
[He]2s22p3
[Ne]3s23p3
[Ar]4s24p3 [Kr]5s25p3 [Xe]6s26p3
–210.01
–195.79
1.251
44.1
280
2.34
817
630.63
5.7
630
1750
6.7
271
1560
9.8
3.04
2.19
2.18
2.05
2.02
0.77
1.48
0.15
0.71
1.1
1.86
0.35
1.3
1.21
1.92
0.47
1.48
1.41
2.08
0.62
1.61
–
2.13
0.74
1.82
Physical Properties and Occurrence
Nitrogen, symbol N, is a gaseous element that makes up the largest part of the
earth's atmosphere. The atomic number of nitrogen is 7. Nitrogen is a colorless,
odorless, tasteless, nontoxic gas. It can be condensed into a colorless liquid, which
in its turn can be compressed into a colorless, crystalline solid. Nitrogen melts at –
210.01 °C, boils at – 195.79 °C, and has a density of 1.251 at 0 °C and 1
atmosphere pressure. The atomic weight of nitrogen is 14.007.
35
Nitrogen was isolated by the British physician Daniel Rutherford in 1772 and
recognized as an elemental gas by the French chemist Antoine Laurent Lavoisier
about 1776.
Nitrogen composes about four-fifths (78.03 percent) by volume of the
atmosphere. The element occurs in the combined state in minerals, of which
saltpeter (KNO3) and Chile saltpeter (NaNO3) are commercially important
products.
Phosphorus, symbol P, is a reactive nonmetallic element. The atomic number
of phosphorus is 15, and its atomic weight is 30.974. Phosphorus exists in three
main allotropic forms: ordinary (or white) phosphorus, red phosphorus, and black
phosphorus. When freshly prepared, ordinary phosphorus is white, but it turns light
yellow when exposed to sunlight. It is a crystalline, translucent, waxy solid, which
glows faintly in moist air and is extremely poisonous. It ignites spontaneously in
the air at 34 °C and must be stored under the water. It is insoluble in water, slightly
soluble in organic solvents, and very soluble in carbon disulfide. White phosphorus
melts at
44.1 °C and boils at 280 °C. When heated to 230 – 300 °C in the
absence of air, white phosphorus converts into the red form. Red phosphorus is a
microcrystalline, nonpoisonous powder. It sublimates at 416 °C and has a specific
gravity of 2.34. Black phosphorus is made by heating white phosphorus at 200 °C
at very high pressure. It has a specific gravity of 2.69.
Phosphorus was discovered about 1669 by the German alchemist Hennig Brand
in the course of experiments in which he attempted to prepare gold from silver.
Phosphorus is widely distributed in nature and ranks 11th in abundance among
the elements in the crust of the earth. It does not occur in the free state but is found
mostly as a phosphate, like in phosphate rock and apatite (Ca5(PO4)3F). It is found
also in the combined state in all fertile soil and in many natural waters.
Arsenic (from the Greek word "arsenikon" meaning "yellow orpiment"),
symbol As, is an extremely poisonous semimetallic element. The atomic number of
arsenic is 33. A common form of arsenic is gray, metallic in appearance, and has a
specific gravity of 5.7. A yellow, nonmetallic form also exists which has a specific
gravity of 2.0. The atomic weight of arsenic is 74.9216.
Arsenic has been known since ancient times. The pure element can be easily
prepared by heating a common ore called arsenopyrite (FeAsS). Other common
minerals are realgar (As2S2); orpiment (As2S3); and arsenic trioxide (As2O3);
occasionally the pure element is found in nature. Arsenic occurs frequently
together with sulfur in the form of sulfides that are the principal ores of many of
the heavy metals. When these ores are roasted, the arsenic sublimes and can be
collected from the dust in the flues as a by-product. Arsenic ranks about 52nd in
natural abundance among the elements in crustal rocks.
Antimony (from the Greek words "anti + monos" meaning "not alone"),
symbol Sb, is a bluish-white, brittle, semimetalic element. The atomic number of
antimony is 51. The atomic weight of antimony is 121.75; it melts at about 630 °C,
boils at about 1750 °C, and has a specific gravity of 6.7.
36
Antimony's compounds were known in ancient times, and the element was
probably discovered by the German alchemist Basil Valentine about 1450. It was
certainly known by about 1600, but was confused with other elements, such as
bismuth, tin, and lead.
Antimony ranks about 64th in natural abundance among the elements in crustal
rock. It occasionally occurs as a free element, usually associated with silver,
arsenic, or bismuth. The principal ore of antimony is stibnite (antimony trisulfide,
Sb2S3). Considerable amounts of antimony are produced as a by-product in the
refining of ores of copper and lead.
Bismuth (From the German word "bisemutum"), symbol Bi, is a rare metallic
element that has a pinkish tinge. The atomic number of bismuth is 83. It melts at
271 °C, boils at 1560 °C, and has a specific gravity of 9.8. The atomic weight of
bismuth is 208.98.
Bismuth was known in ancient times, but until the middle of the 18th century it
was confused with lead, tin, and zinc. Although bismuth had been discussed many
times before, Claude Geoffroy the Younger showed it to be distinct from lead in
1753. Ranking about 73rd in natural abundance among the elements in the earth's
crust, it is about as rare as silver.
Nitrogen and its Compounds
Nitrogen is the first member of group 5 of the periodic table. It has the electron
configuration 1s22s22p3 with a single electron in each 2p orbital.
Nitrogen exhibits positive oxidation states only in its compounds with oxygen
and fluorine, because only these elements are more electronegative than nitrogen.
In its compounds nitrogen has oxidation numbers from –3 to +5 (Table 3.2).
Table 3.2. Oxidation States of Nitrogen
Oxidation State
Example
+5
HNO3, N2O5, NO3–
+4
NO2, N2O4
+3
HNO2, N2O3, NO2–
+2
NO
+1
N2O
0
N2
–1
NH2OH
–2
N2H4
–3
NH3
37
Preparation of Nitrogen
There is no need to obtain nitrogen in the laboratory as it is readily available
commercially or through in-house air liquefaction plants. However the
decomposition of sodium azide or ammonia nitrite is one route to N2 preparation
and decomposition of ammonium dichromate is the another. Both reactions must
be carried out under controlled conditions by a professional only:
2NaN3 → 2Na + 3N2↑;
NH4NO2 → N2↑ + 2H2O;
(NH4)2Cr2O7 → N2↑ + Cr2O3 + 4H2O.
Nitrogen is made in great scale by liquefaction of air and further fractional
distillation of the liquid air for separating out oxygen and other gases. Very high
purity nitrogen is available by this route.
300 o C
Chemical Properties of Nitrogen and its Compounds
A diatomic molecule on nitrogen N2 is very stable; there is a triple bond
between the nitrogen atoms in it: N≡N. The molecule on nitrogen gas doesn’t
disintegrate into atoms even at high temperature (only 0.1% of molecules are
atomized at 3000 0C).
Nitrogen gas does not react with air under normal conditions.
Nitrogen gas does not react with water. It does, however, dissolve to the extent
of about 2.33 ml in 100 ml of water at 0 °C and 1 atmosphere pressure.
When nitrogen is heated, it combines directly with magnesium, lithium, or
calcium.
6Li + N2 → 2Li3N lithium nitride;
N2 + 2Ca → Ca3N2 calcium nitride.
Nitrides of metals have basic (Na3N, Mg3N2) or amphoteric (AlN) properties,
and nitrides of nonmetals are acidic ones (Si3N4, P3N5, S4N4, Cl3N).
Nitrides of alkali and alkaline earth metals are highly chemically reactive
compounds. They can react with water forming the corresponding alkali and
ammonia:
Li3N + 3H2O → 3LiOH + NH3.
Basic and acidic nitrides interact between themselves giving blended nitrides:
Li3N + AlN → Li3AlN2.
When mixed with oxygen and subjected to electric sparks, it forms nitric oxide
(NO):
N2 + O2 ⇄ 2NO,
and then the dioxide (NO2):
2NO + O2 → 2NO2.
When nitrogen heated under increased pressure with hydrogen in the presence
of a proper catalyst, ammonia forms (Haber process):
N2 + 3H2 ⇄ 2NH3.
38
AMMONIA
Ammonia is a colorless gas which has a typical suffocated odor. It readily
dissolves in water (1150 liters of ammonia may be dissolved in 1 liter of water at 0
0
C).
Ammonia was firstly prepared under the reaction of nitrogen with hydrogen:
N2 + 3H2 ⇄ 2NH3.
The most optimal conditions for the reaction are the pressure of 2×104 kPa,
temperature of 500 0C, and the presence of a proper catalyst.
In the laboratory ammonia is prepared by heating the mixture of an ammonia
salt and alkali:
(NH4)2SO4 + 2NaOH → Na2SO4 + 2NH3↑ + 2H2O,
or by heating ammonia chloride with slake lime:
2NH4Cl + Ca(OH)2 → CaCl2 + 2NH3↑ + 2H2O.
Only three p-electrons of nitrogen take part in the formation of chemical bonds
in the molecule of ammonia, and the nitrogen atom has free unshared pair of
electrons also. Thus the molecule of NH3 is a strong donor of electrons. Selfionization of ammonia is very small, and its ionization constant is 2 × 10-33 (at –50
0
C).
H
H
H
+
H N H
H N : ... H N :
H
+
..
H
:N
H
-
H
H
Dihydronitride-ion
Ammonia cation
The molecule of ammonia can join H+-ion and transfer into ammonia-cation:
..
N H3 + H → [NH4] ;
+
+
NH3 + H2O ⇄ NH4+ + OH−.
The medium of ammonia aqueous solutions is basic one.
Gaseous ammonia and ammonia in aqueous solutions react with acids forming
corresponding salts of ammonia:
NH3 + HCl → NH4Cl.
Ammonia can also interact with metals. In such reactions hydrogen atoms in
the molecule of ammonia are replaced with metals atoms giving the metals amides:
2NH3 + 2Na → 2NH2Na + H2↑.
sodium amide
Ammonia is a strong reducing agent. The nitrogen atom in the reactions may
change its oxidation number from –3 to 0 or +2. Halogens are commonly oxidizing
ammonia to free nitrogen gas:
8NH3 + 3Br2 → N2↑ + 6NH4Br.
Ammonia can burn in the atmosphere of oxygen with greenish-yellow flame:
39
4NH3 + 3O2 → 2N2↑ + 6H2O.
In the presence of the catalyst ammonia interacts with oxygen giving nitrogen
(II) oxide NO:
4NH3 + 5O2 → 4NO↑ + 6H2O.
Ammonia may readily reduce some metals from their oxides:
3CuO + 2NH3 → N2↑ + 3Cu + 3H2O.
Ammonia may be oxidized to nitrate-ion with potassium chlorate at 300 0C:
6NH3 + 8KClO3 + 6NaOH → 6NaNO3 + 8KCl + 12H2O.
Liquid ammonia is a strong ionizing solvent. Thus ammonia salts such as
NH4Cl or NH4NO3 in the medium of liquid ammonia display acidic properties:
KNH2 + NH4Cl → KCl + 2NH3↑;
Ca3N2 + 6NH4Cl → 3CaCl2 + 8NH3↑;
2KNH2 + Zn(NH2)2 → K2[Zn(NH2)4].
Ammonium salts have crystalline structure. They are ionic compounds which
dissociate almost completely into ions in aqueous solutions:
NH4Cl ⇄ NH4+ + Cl−;
NH4+ + H2O ⇄ NH3 + H3O+.
Ammonium salts easily decompose, completely or partially, when heated:
heating
NH4Cl
⇔
cooling
NH3 + HCl;
(NH4)2SO4 ⇄ NH3↑ + NH4HSO4.
Ammonium salts are chemically reactive. While heating ammonium salts with
metals hydroxides solutions ammonia gas releases:
NH4Cl + NaOH → NaCl + NH3↑ + H2O.
Ammonium salts which have the anion of oxidizing nature decompose when
heated:
NH4NO3 → N2O↑ + 2H2O;
NH4NO2 → N2↑ + 2H2O.
AMMONIUM CHLORIDE, NH4Cl
Ammonium chloride is used when metals are soldering for taking off the oxide
layer from the metals surface:
4CuO + 2NH4Cl → 3Cu + CuCl2 + N2↑ + 4H2O,
Fe3O4 + 8NH4Cl → FeCl2 + 2FeCl3 + 8NH3↑ + 4H2O.
HYDRAZINE, N2H4
It is useful to consider hydrazine as being derived from ammonia by replacement of a hydrogen atom by an –NH2 group:
40
Ammonia
Hydrazine
Both nitrogen atoms in N2H4 have the expected pyramidal AX3E geometry.
An aqueous solution of hydrazine can be prepared by oxidizing ammonia with a
solution of sodium hypochlorite:
2NH3 + NaOCl → N2H4 + H2O + NaCl.
The chloramines, H2NCl and HNCl2, are produced also in this reaction. They
both are toxic and explosive. Thus it is dangerous to mix household bleach and an
ammonia cleaning solution.
Hydrazine is a colorless liquid that melts at 2 °C and boils at 114 °C. It is a
weak base, which is protonated giving the N2H5+ and N2H62+ ions. Hydrazine has
an endothermic enthalpy of formation (∆H° = 51 kJ/mol); therefore it is a
somewhat unstable compound.
Hydrazine bums in the air with considerable evolution of heat:
N2H4 + O2 → N2 + 2H2O ∆H° = –622 kJ.
Hydrazine is a good reducing agent and reacts vigorously with strong oxidizing
agents such as halogens, nitric acid, dinitrogen tetraoxide, and hydrogen peroxide
with the evolution of a large amount of heat and the generation of a large volume
of gaseous products:
2N2H4 + N2O4 → 3N2 + 4H2O;
N2H4 + 2H2O2 → N2 + 4H2O.
The oxidation of N2H4 or substituted hydrazines, where one or more of the
hydrogens are replaced by methyl groups, is used to power rockets. For example at
the Apollo missions to the moon, the rocket engines of the command module and
the lunar landing vehicles used N2O4 as an oxidizer and a mixture of hydrazine and
dimethylhydrazine as fuel. The U.S. space shuttle orbiter also used N2O4 and
monomethylhydrazine. They react giving water, nitrogen, and carbon dioxide:
5N2O4 + 4N2H3(CH3) → 12H2O + 9N2 + 4CO2.
This reaction begins immediately when the reactants are mixed which enables
the rocket engines to be started and stopped as required. Since the reaction is
highly exothermic and gives a large number of gaseous molecules, a high thrust is
obtained from minimum mass of fuel.
HYDRAZOIC ACID, HN3
Hydrazoic acid is a liquid that boils at 37 °C and it is dangerously explosive. It
is a weak acid in water. Salts of hydrazoic acid are called azides. The azides of
heavy metals such as lead, mercury, and barium explode when being struck sharply
and are used as detonators.
41
The azide ion, N3–, is isoelectronic with CO2, N2O, and NO2+, and it has the
expected linear AX2 structure with a bond length of 116 pm. It can be represented
by the following two resonance structures:
HYDROXILAMINE, NH2OH
According to its structure NH2OH lies between hydrazine and hydrogen
peroxide:
..
..
.. H
H
H .. .. H
O N
O
O
N N
H
..
H ..
H ..
H
H
hydroxilamine
hydrazine
hydrogen peroxide
NH2OH may be prepared by the reducing of nitric acid under the reaction of
electrolysis:
HNO3 + 6H  → NH2OH + 2H2O.
Hydroxilamine is a weaker base than ammonia:
electrolysis
NH2OH + HOH ⇄ NH3OH+ + OH−.
It can interact with acids giving corresponding stable salts:
NH2OH + HCl → [NH3OH]+Cl−.
Hydroxilamine displays oxidizing properties in acidic medium, and it displays
reducing properties in basic medium:
2NH2OH + I2 + 2KOH → N2↑ + 2KI + 4H2O,
2NH2OH + 4FeSO4 + 3H2SO4 → 2Fe2(SO4)3 + (NH4)2SO4 + 2H2O.
NITROGEN OXIDES
Nitrogen forms the following oxides: nitrogen (I) oxide N2O, nitrogen (II)
oxide NO, nitrogen (III) oxide N2O3, nitrogen (IV) oxide NO2, and nitrogen (V)
oxide N2O5. The oxides of nitrogen and some of their physical properties are listed
in Table 3.3.
Table 3.3 Properties of Nitrogen Oxides
Oxidation
State
+5
+4
Formula
Name
N2O5
N2O4
Dinitrogen pentaoxide
Dinitrogen tetraoxide
+4
NO2
Nitrogen dioxide
+3
N2O3
Dinitrogen trioxide
+2
+1
NO
N2O
Nitrogen monoxide
Dinitrogen monoxide
* Decomposes
42
Physical State at
25 °C
White solid
White solid
Brown gas
Deep blue liquid
(–80°C)
Colorless gas
Colorless gas
Tm , °C
Tb , °C
30
47*
–11.2
21.2
–102
–90.8
–163.7
–*
–88.5
N2O
DINITROGEN MONOXIDE, N2O
When solid ammonium nitrate is gently heated, it melts and decomposes to give
dinitrogen monoxide, N2O:
∆
NH4NO3 
→ N2O + 2H2O, ∆H° = –37 kJ.
In this oxidation-reduction reaction, nitrogen in the –3 oxidation state in NH3 is
oxidized by nitrogen in the +5 state in the nitrate ion, giving N2O in which nitrogen
is in the +1 oxidation state. Since the reaction forms 3 moles of gaseous products
and it is exothermic, it can occur with explosive violence if the solid would be
heated too strongly. Dinitrogen monoxide may be described by two resonance
structures:
So N2O has a linear structure.
Dinitrogen monoxide is a rather unreactive substance.
It decomposes to nitrogen gas and oxygen when heated up to 700 0C:
2N2O → 2N2 + O2.
Therefore N2O is an oxidizing agent for those compounds which can react with
hydrogen:
N2O + H2 → N2 + H2O.
It is a colorless gas with a pleasant smell and a sweet taste. When breathed in
small amounts, it acts as a mild intoxicant. Because of this property, it has been
called laughing gas. In large amounts N2O acts as a general anesthetic and is
sometimes used for this purpose in dental and other minor surgery. Having a useful
property of being rather soluble in fats dinitrogen monoxide has also been
exploited in making self-whipping cream. Cream is packaged with N2O under the
pressure to increase its solubility. When the pressure is released, some of N2O
escapes forming tiny bubbles, which produce whipped cream.
NITROGEN MONOXIDE, NO
NO can be called nitrogen(II) oxide also. It is a colorless, reactive gas which is
only slightly soluble in water.
Nitrogen combines with oxygen slowly and at very high temperatures to give
NO:
∆ 2NO ∆H°= 180 kJ.
N + O →
2
2
The reaction is endothermic.
Nitrogen monoxide is prepared on a large scale in the first step of the Ostwald
process for the manufacture of nitric acid. Ammonia and oxygen are passed over
platinum - rhodium wire gauze which should be heated to approximately 900 °C by
an electric current. The hot wire gauze catalyzes the oxidation of ammonia to
nitrogen monoxide:
∆ → 4NO + 6H O ∆H° = –906 kJ.
4NH + 5O Catalyst,


3
2
2
43
Once the reaction has started, it supplies the heat needed to keep the catalyst at
the operating temperature because it is strongly exothermic.
Nitrogen monoxide can be conveniently prepared in the laboratory by the
reaction of dilute nitric acid with copper:
3Cu + 8HNO3(dil) → 3Cu(NO3)2 + 2NO↑ + 4H2O .
Because the nitrogen in NO is in a low (+2) oxidation state, NO behaves as a
good reducing agent. With oxygen it forms nitrogen dioxide, NO2:
2NO + O2 → 2NO2.
Nitrogen monoxide reduces ozone to oxygen and is oxidized to NO2,
NO + O3 → O2 + NO2,
and at high temperatures it reduces CO2 to CO:
NO + CO2 → CO + NO2.
With chlorine it forms nitrosyl chloride, NOCl, an orange-yellow gas that
contains nitrogen in the +3 state:
2NO + Cl2 → 2NOCl,
When bubbled into aqueous solutions of oxidizing agents, NO is usually oxidized
to nitrate ion. For example, it reduces iodine to iodide ion:
2NO + 3I2 + 4H2O → 2HNO3 + 6HI,
and MnO4– to Mn2+:
10NO + 6KMnO4 + 9H2SO4 → 10HNO3 + 6MnSO4 + 3K2SO4 + 4H2O.
Because there are stable lower oxidation states of nitrogen, NO will react also
as an oxidizing agent. For example, it is reduced to nitrogen by hydrogen, which is
oxidized to water, and by phosphorus, which is oxidized to P4O6:
2NO + 2H2 → N2 + 2H2O,
P4 + 6NO → P4O6 + 3N2.
DINITROGEN TRIOXIDE, N2O3
Nitrogen monoxide and nitrogen dioxide combine at low temperature by
pairing their odd electrons and give dinitrogen trioxide, N2O3:
NO + NO2 → N2O3.
Figure 3.2. Structure of Dinitrogen Trioxide.
(a) dimensions; (b) resonance structures.
It may be prepared also by decomposition of nitrous acid:
44
2HNO2 ⇄ N2O3 + H2O.
In the laboratory N2O3 can be obtained by the reaction between an alkalis metal
nitrite with an acid:
2NaNO2 + 2H2SO4 → 2NaHSO4 + 2HNO2;
2HNO2 ⇄ N2O3 + H2O.
This deep blue liquid decomposes completely at room temperature to NO and
NO2:
N2O3 →NO + NO2.
N2O3 is an acidic oxide and readily reacts with alkalis giving corresponding
nitrites – salts of nitrous acid:
N2O3 + 2NaOH → 2NaNO2 + H2O.
NITROGEN DIOXIDE, NO2, AND DINITROGEN TETRAOXIDE, N2O4
The product of the reaction between NO and oxygen is nitrogen dioxide, NO2:
2NO + O2 → 2NO2.
It is a red-brown gas. It may be called also nitrogen(IV) oxide. Two molecules
of NO2 combine with each other to form colorless dinitrogen tetraoxide, N2O4. At
ordinary temperatures the gas is an equilibrium mixture of NO2 and N2O4:
2NO2 ⇄ N2O4 ∆H= –57 kJ.
Since the formation of N2O4 is an exothermic reaction, the equilibrium shifts to
the right with the temperature decreasing, and more N2O4 is formed. At 21 °C the
gas condenses to a deep red-brown liquid, which is a mixture of NO2 and N2O4 in
equilibrium. When the temperature decreases, the color of the liquid mixture
becomes less intense. Finally, at –11 °C it freezes to a white solid that consists
entirely of N2O4 molecules.
A Lewis structure for N2O4 may be written by combining two NO2 structures in
which the odd electron is on nitrogen. All four resonance structures can be
obtained in this way (see Figure 3.3).
Figure 3.3. Structure of Dinitrogen Tetraoxide. Dinitrogen tetraoxide is a planar
molecule with the dimensions shown in (a). It can be represented by the four
resonance structures (b).
45
Nitrogen dioxide can be conveniently made in the laboratory by reducing
concentrated nitric acid with copper:
Cu + 4HNO3(conc.) → Cu(NO3)2 + 2NO2↑ + 2H2O.
It can be produced also by heating certain metal nitrates. For example,
∆ 2PbO + 4NO ↑ + O ↑.
2Pb(NO ) →
3 2
2
2
Since nitrogen is in a high (+4) oxidation state, NO2 is a good oxidizing agent.
For example, it oxidizes SO2 to SO3:
SO2 + NO2 → SO3 + NO.
It oxidizes CO to CO2:
4CO + 2NO2 → 4CO2 + N2
and I– to I2:
8KI + 2NO2 + 4H2O → 4I2 + N2 + 8KOH.
When NO2 is dissolving in water it auto-oxidizes:
2NO2 + H2O → HNO3 + HNO2.
Nitrates and nitrites are forming when nitrogen (IV) oxide reacts with alkali:
2NO2 + 2NaOH → NaNO3 + NaNO2 + H2O.
But since it can be oxidized to the +5 oxidation state it can act also as a reducing
agent and, for example, it will reduce MnO4– to Mn2+ in aqueous solution:
10NO2 + 2KMnO4 + 3H2SO4 + 2H2O → 10HNO3 + 2MnSO4 + K2SO4.
When NO2 is dissolving in water in the presence of the oxygen excess nitric
acid forms:
4NO2 + O2 + 2H2O → 4HNO3.
DINITROGEN PENTAOXIDE, N2O5
The oxide N2O5 is made by dehydrating of 100 % nitric acid, HNO3, with phosphorus(V) oxide, P2O5:
2HNO3 + P2O5 → 2N2O5 + 2HPO3.
Its structure is shown in Figure 3.4. Since it is formed by the removal of water
from nitric acid, dinitrogen pentaoxide may be described as the anhydride of nitric
acid:
2HNO3 → N2O5 + H2O (combines with P2O5).
Figure 3.4. Structure of Dinitrogen Pentaoxide.
(a) in the gas state, dinitrogen pentaoxide consists of N2O5 molecules. One of
four equivalent resonance structures is shown here.
(b) in the solid state, dinitrogen pentaoxide is an ionic compound composed
of NO2+ and NO3– ions.
46
Dinitrogen pentaoxide is a white volatile solid which decomposes slowly at
room temperature to NO and O2:
2N2O5 → 4NO2 + O2.
It reacts readily with water to give nitric acid:
N2O5 + H2O → 2HNO3.
Dinitrogen pentaoxide is an interesting example of a substance that has
different structures in the solid and gaseous states. In the gas phase dinitrogen
pentaoxide consists of covalent N2O5 molecules. In contrast, the white solid is an
ionic compound composed of nitronium ions, NO2+, and nitrate ions, NO3–.
Therefore it might be called nitronium nitrate.
Oxoacids of Nitrogen
Following are the important oxoacids formed by nitrogen: Hyponitrous acid,
H2N2O2, or HNO, Nitrous acid, HNO2, Nitric acid, HNO3, and Pernitric acid,
HNO4.
NITROUS ACID, HNO2
The Lewis structure of nitrous acid is
Nitrous acid is an unstable acid and is known only in solution.
A solution of nitrous acid may be prepared by the action of water on its
anhydride, N2O3:
N2O3 + H2O → 2HNO2,
or by barium nitrite treatment with calculated quantity of dilute sulfuric acid. The
precipitate of barium sulfate so formed is filtered off and the filtrate contains
nitrous acid. Since the acid is very unstable, the reaction is carried out at the
temperature of freezing mixtures:
Ba(NO2)2 + H2SO4 → BaSO4↓ + 2HNO2.
Nitrous acid is a weak acid. Aqueous solution of nitrous acid is pale blue due to
the presence of nitrogen trioxide. This colour fades when standing for sometime.
The acid is unstable and even in cold solution undergoes auto-oxidation:
3HNO2 → 2NO + HNO3 + H2O.
At higher temperature it decomposes into nitric oxide and nitrogen peroxide:
2HNO2 → N2O3 + H2O;
N2O3 ⇄ NO + NO2.
It acts as a good oxidizing agent. For example, it oxidizes hydrogen sulfide to
sulfur, liberates iodine from potassium iodine, oxidizes sulfur dioxide to sulfuric
acid, Fe2+salts to Fe3+ salts:
2HNO2 + H2S → 2NO + S + 2H2O;
2KI + 2HNO2 → 2KOH + I2 + 2NO;
2HNO2 + SO2 → 2NO + H2SO4;
2FeSO4 + H2SO4+ 2HNO2 → Fe2(SO4)3 + 2NO + 2H2O.
47
Nitrous acid can be easily oxidized to nitric acid and thus reduces strong
oxidizing agents. For example, it reduces bromine and iodine to hydrobromic acid
and hydroiodic acids, respectively:
HNO2 + I2 + H2O → 2HI + HNO3;
HNO2 + Br2 + H2O → 2HBr + HNO3.
Acidified potassium permanganate is reduced to manganese sulfate, and
acidified potassium dichromate is reduced to chromium sulfate (green color):
2KMnO4 + 3H2SO4 + 5HNO2 → K2SO4 + 2MnSO4 + 5HNO3 + 3H2O,
K2Cr2O7 + 4H2SO4 + 3HNO2 → K2SO4 + Cr2(SO4)3 + 3HNO3 + 4H2O.
It reduces hydrogen peroxide also:
H2O2 + HNO2 → HNO3 + H2O.
Nitrous acid reacts with alkalis forming nitrites (salts of nitrous acid). Nitrites
are much more stable than nitrous acid:
NaOH + HNO2 → NaNO2 + H2O;
KOH + HNO2 → KNO2 + H2O.
Nitrites may be prepared also by passing nitrogen trioxide through the solution
of an alkali:
N2O3 + 2NaOH → 2NaNO2 + H2O,
or alkali nitrites are prepared by heating the alkali nitrates with lead:
KNO3 + Pb → KNO2 + PbO.
The nitrites are stable substances. They are soluble in water, excepting silver
nitrite which dissolves in hot water. Except the alkali nitrites, they decompose
when heated. With dilute acids, in the cold, they give nitrous acid:
NaNO2 + HCl → NaCl + HNO2.
At ordinary temperatures, reddish brown fumes are obtained by the reaction
with concentrated acids:
2NaNO2 + 2HCl → 2NaCl + NO + NO2 + H2O.
Since they are readily oxidized to nitric acid and nitrate ion, nitrous acid and
the nitrites are good reducing agents:
NO2– + H2O → NO3– + 2e– + 2H+
Figure 3.5 Structure of the Nitrite Ion.
Nitrous acid is used in organic chemistry especially in preparation of diazo
compounds which are the basic substances in the manufacture of the aniline dyes.
It is used also as an oxidizing as well as a reducing agent in analytical chemistry.
NITRIC ACID, HNO3
In the laboratory nitric acid is prepared by heating concentrated sulfuric acid
with an alkali nitrite in a glass retort:
48
NaNO3 + H2SO4 (conc.) → NaHSO4 + HNO3.
Nitric acid is prepared on a large scale by heating sodium nitrate with
concentrated sulfuric acid:
∆ 2HNO + Na SO .
2NaNO + H SO (conc.) →
3
2
4
3
2
4
The mixture is placed in a cast iron retort and subjected to distillation at about
200 0C. The vapours of nitric acid from the pot are conducted to a trap and then
passed through water-cooled silica pipes. The acid obtained in this way is
concentrated by distilling the dilute acid with concentrated sulfuric acid.
Commercial nitric acid contains hydrochloric acid, sulfuric acid, iodine, iron salts
and nitrogen pentoxide as impurities. A current of dry air passed through the liquid
removes N2O5. Further purification is carried out by distilling the acid under
reduced pressure.
Nitric acid may be prepared also from air (Birkland-Eyde process). Nitrogen
and oxygen of the air combine forming nitric oxide when air is passed through an
electric arc:
N2 + O2 ⇄ 2NO.
The reaction is endothermic and is favoured by a high temperature. At 600 0C the
nitric oxide begins to combine with oxygen forming nitrogen peroxide:
2NO + O2 ⇄ 2NO2.
The oxide is then absorbed in water trickling over quartz stone in towers about
70 ft high and 20 ft in diameter. The nitric acid in the tower reaches strength of 30
to 40 per cent. The nitric acid is formed according to the equation:
2NO2 + H2O → HNO2 + HNO3
in the case of cold dilute solution. In the hot state, the nitrous acid undergoes autooxidation to nitric acid, and the nitric oxide formed again takes up oxygen and
forms nitrogen dioxide:
3HNO2 → HNO3 + 2NO + H2O.
Ostwald’s process of nitric acid preparation is based on catalytic oxidation of
ammonia. In this method a mixture of ammonia gas with air in the ratio 1:8 is
passed through a catalyst chamber packed with platinum gauze maintained at the
temperature of 803 0C. The ammonia oxidizes under these conditions to nitrogen
monoxide:
4NH3 + 5O2 → 4NO + 6H2O.
The nitrogen monoxide so formed is led to an oxidizing tower where it combines
with oxygen of the air forming nitrogen dioxide, NO2:
2NO + O2 → 2NO2.
The resulting gas passed up a tower packed with earthen balls down with a spray of
water trickles. Nitrogen dioxide reacts with water in the presence of oxygen
forming nitric acid:
3NO2 + H2O → 2HNO3 + NO.
The acid obtained is dilute and is concentrated by distillation under the reduced
pressure.
49
The dilute HNO3 can be concentrated to 68% by distillation when constant
boiling mixture is formed. More concentrated acid needed for certain uses is
produced by distillation with conctntrated H2SO4 which holds back the water. The
distillate is 98% HNO3.
A recently developed method for the direct production of concentrated HNO3
involves the reaction between the liquid NO2 and water (or dilute HNO3) in the
presence of O2 at 50 atm. pressure and a temperature of 75 0C. The product is 98 %
HNO3:
4NO2 + O2 + 5H2O → 4HNO3.
Fuming nitric acid is made by distilling concentrated HNO3 with a little
amount of starch. The starch reduces some acid to NO2 which dissolves in the
remaining acid forming fuming nitric acid. This acid is yellow in color due to the
presence of dissolved nitrogen oxides in it. Fuming nitric acid is a much more
powerful oxidizing agent and nitrating agent than concentrated HNO3.
Properties of nitric acid
Pure HNO3 is colorless liquid of specific gravity 1.56 (at 0 0C). It boils at 86
0
C. It is soluble in water in all proportions. The acid is highly corrosive and
produces painful blisters when it comes in contact with skin.
The nitric acid molecule can be represented by two resonance structures:
Figure 3.6. Dimensions of HNO3 molecule
HNO3 is colorless but often gradually turns yellow. Yellow color produces due
to the fact that HNO3, when exposed to light, undergoes decomposition and forms
NO2 which then dissolves in HNO3 and colors it in yellow:
hν
4HNO3 → 4NO2 + O2 + 2H2O.
At a room temperature the pure acid undergoes decomposition giving N2O2 and
O2:
room temperature
2HNO3     → N2O5 + H2O
At heating HNO3 decomposes to give O2 and brown fumes of NO2:
∆
4HNO3 
→ 4NO2 + O2 + 2H2O.
HNO3 is a very strong monobasic acid. Therefore it reacts with basic oxides,
carbonates, bicarbonates, sulfites and hydroxides, forming corresponding nitrates:
CaO + 2HNO3 → Ca(NO3)2 + H2O,
50
Na2CO3 + 2HNO3 → 2NaNO3 + H2O + CO2,
NaHCO3 + HNO3 → NaNO3 + H2O + CO2,
NaOH + HNO3 → NaNO3 + H2O,
Na2SO3 + 2HNO3 → 2NaNO3 + H2O + SO2.
Being a strong acid, HNO3 is completely ionized in water:
HNO3 → H+ + NO3−.
HNO3 is a very strong oxidizing agent and can be reduced to NO2, NH4NO3 (or
NH4+ ions), N2O, NO, N2 or NH2OH depending on the nature of the substance
being oxidized, concentration of HNO3 and the temperature applied.
When non-metals are treated with HNO3 they are oxidized to their
corresponding highest oxy-acids and HNO3 is reduced to NO2. Thus, for example:
P + 5HNO3 → H3PO4 + 5NO2 + H2O;
C + 4HNO3 → H2CO3 + 4NO2 + H2O;
S + 6HNO3 → H2SO4 + 6NO2 + 2H2O;
conc. and hot
I2 + 10HNO3 → 3HIO3 + 10NO2 + 4H2O.
Concentrated HNO3 can oxidize all metals (except gold and platinum).
Oxidation of metals which lie above hydrogen in the electrochemical series
• Mg and Mn liberate H2 when treated with dilute HNO3 (1-2 %):
Me (Me = Mg, Mn) + 2HNO3 → Me(NO3)2 + H2↑.
• Zn, Fe, Sn and Pb are oxidated to corresponding nitrates when treated with
concentrated and cold nitric acid (70%). HNO3 is redused to NO2:
Zn + 4HNO3(conc.) → Zn(NO3)2 + 2NO2 + 2H2O.
When they are interacting with dilute nitric acid (20%), HNO3 is reduced to
N2O:
4Zn + 10HNO3(dil.) → 4Zn(NO3)2 + N2O + 5H2O.
When Zn reacts with very dilute nitric acid (6%), HNO3 is reduced to NH4NO3:
4Zn + 10HNO3(very dil.) → 4Zn(NO3)2 +NH4NO3 + 3H2O.
Highly concentrated HNO3 renders Fe, Cr, and Al passive (inert).
Oxidation of metals which lie below hydrogen in the electrochemical series
• cold and dilute HNO3 evolves N2O:
4 Cu + 10HNO3(dil.) → N2O + 4Cu(NO3)2 + 5H2O;
• cold and moderately concentrated HNO3 evolves NO:
3Cu + 8HNO3(mod.conc.) → 3Cu(NO3)2 + 2NO + 4H2O;
• cold and concentrated HNO3 evolves NO2:
Cu + 4HNO3(conc.) → Cu(NO3)2 + 2NO2 + 2H2O;
• hot and concentrated HNO3 evolves N2:
5Cu + 12HNO3(hot and conc.) → 5Cu(NO3)2 + N2 + 6H2O.
The mixture of one volume of nitric acid and three volumes of hydrochloric
acid is called “aqua regia” and is a stronger oxidizing agent:
HNO3 + 3HCl → NO + 2H2O + 3Cl (atomic chlorine evolves).
“Aqua regia” can dissolve gold and platinum:
51
Au + HNO3 + 4HCl → H[AuCl4] + NO + 2H2O.
Reaction of NO3− ion identification
• in the alkali medium a metal nitrate may be reduced to ammonia with Zn or
Al:
4Zn + NaNO3 + 7NaOH + 6H2O → 4Na2[Zn(OH)4] + NH3↑.
The releasing of ammonia may be detected by the lathmus paper turning
blue.
• Iron (II) sulfate in the acidic medium is oxidized with nitric acid to
iron (III) sulfate. Nitric acid is reduced to NO which forms with the excess
of FeSO4 the complex compound of brown color:
2NaNO3 + 6FeSO4 + 4H2SO4 → 3Fe2(SO4)3 + 2NO + 4H2O,
NO + FeSO4 → [Fe(NO)]SO4.
Salts of nitric acid are called nitrates. They have crystal structure and readily
dissolve in water.
Nitrates decompose at heating. The alkali metals nitrates turn to nitrites when
heated and oxygen gas evolves:
2KNO3 → 2KNO2 + O2↑.
Nitrates of metals which lie below hydrogen in the electrochemical series
ingdecompose give corresponding metals:
2AgNO3 → 2Ag + NO2↑ + O2↑.
All another metals nitrates decompose forming oxides:
2Cu(NO3)2 → 2CuO + 4NO2↑ + O2↑,
2Pb(NO3)2 → 2PbO + 4NO2↑ + O2↑.
Phosphorus and its Compounds
Phosphorus has several allotropes. White phosphorus is obtained when
phosphorus vapor is condensed. It is a colorless crystalline solid, with a melting
point of 44 °C, but it rapidly becomes white and opaque unless stored under
nitrogen in the dark. White phosphorus is very toxic. The vapor causes decay of
cartilage and bones, particularly of the nose and jaw.
White phosphorus reacts with oxygen at room temperature, emitting a bluish
green light. Some of the energy produced in the reaction is emitted as
electromagnetic radiation. This phenomenon is known as chemiluminescence,
which is the emission of energy released by a reaction as light rather than as heat.
At temperatures above 40 °C the oxidation becomes quite rapid, and phosphorus
ignites. In order to avoid the danger of fire, white phosphorus is stored away from
contact with the air; usually, it is stored under water, in which it is insoluble. It is
soluble, however, in certain other solvents such as carbon disulfide, CS2. White
phosphorus consists of tetraatomic P4 molecules that have a tetrahedral structure
(Figure 3.7). Each of the phosphorus atoms has a pyramidal AX3E geometry. As the
angles at the corners of a face of a tetrahedron are only 60° it is thought that the P–
52
P bonds are bent, so that the angles between the electron pairs in the valence shell
of phosphorus are considerably larger than 60°.
Figure 3.7 The Tetrahedral P4 Molecule.
In this ball-and-stick model the angles between the bonds at each phosphorus atom are considerably
larger than 60° so that the bonds are bent
There are several other allotropes of phosphorus; the best known are red
phosphorus and black phosphorus. Red phosphorus can be obtained by heating
white phosphorus in the absence of air at atmospheric pressure. Black phosphorus
is obtained by heating white or red phosphorus under a very high pressure. Both
allotropes have much higher melting points and boiling points than white
phosphorus. For example, red phosphorus sublimes (changes directly to vapor) at
417 °C and only melts under pressure at 540 °C. Neither red nor black phosphorus
ignites spontaneously in the air, and they are much less reactive and less poisonous
than white phosphorus. These differences result from the differences in their
structures. Red and black phosphorus do not consist of P4 molecules but have
polymeric network structures. The structure of black phosphorus is shown in
Figure 3.8. The structure of red phosphorus is more complicated. To melt and
vaporize a network structure covalent bonds are required to be broken, so the
melting point and boiling point are high. The vapor obtained from both red and
black phosphorus consists of P4 molecules; when the vapor is condensed, white
phosphorus is obtained. Red phosphorus is an essential component of the striking
surface on a box of safety matches.
Figure 3.8 The Structure of Black Phosphorus.
Each P atom has a pyramidal AX3E geometry, but the atoms are not connected so as to form P4
molecules but instead form a corrugated sheet of atoms, that is, a two-dimensional network or giant
two-dimensional molecule. These sheets of atoms are then stacked one upon another in the crystal.
When black phosphorus is vaporized, this structure is broken up and P4 molecules are formed.
Preparation of Phosphorus
The main sources of phosphorus are minerals apatites Ca5Hal(PO4)3 (Hal –
F, Cl or OH) and phosphorites Ca3(PO4)2.
53
The process of phosphorus preparation is carried out in special furnaces at a
temperature of electric arc and consists of two stages:
Ca3(PO4)2 + 8C → Ca3P2 + 8CO;
3Ca3(PO4)2 + 5Ca3P2 → 24CaO + 16P.
To increase the reaction rate SiO2 should be added to the mixture of reactants.
SiO2 reduces the melting point of phosphorite and can bind calcium oxide:
CaO + SiO2 → CaSiO3.
The summary reaction of phosphorus preparation is:
10C + 6SiO2 + 2Ca3(PO4)2 → 10CO + 6CaSiO3 + P4.
Chemical Properties of Phosphorus and its Compounds
Phosphorus has the following ground state configuration:
3s
3p
3d
If an electron is transferred from the 3s-orbital to a 3d-orbital, the electron
configuration becomes the following one:
3s
3p
3d
This valence state has five unpaired electrons. Therefore phosphorus in this
state can form five covalent bonds, as in PCl5 and P4O10.
Phosphorus can be easily oxidizes with oxygen, halogens and another strong
oxidizing agents.
When phosphorus burns in the oxygen atmosphere, phosphorus(V) oxide
forms:
4P + 5O2 → 2P2O5.
When phosphorous is treated with nitric acid phosphorus oxidizes to
PO43- ion:
3P + 5HNO3 + 2H2O → 3H3PO4 + 5NO↑.
In the reactions with halogens phosphorus may give phosphorus(III) and
phosphorus(V) halides:
2P + 3Cl2 → 2PCl3 ;
PCl3 + Cl2 → PCl5 .
In the reactions of phosphorus with metals corresponding phosphides form:
3Mg + 2P → Mg3P2.
In aqueous solutions phosphides hydrolyze:
Ca3P2 + 6H2O → 3Ca(OH)2 + PH3.
Phosphorus is able to undergo self-oxidation reaction:
+1
4P + 3KOH + 3H2O
−3
t

→
3KH2 P O2 + P H3.
PHOSPHINE, PH3
Phosphine is the best known hydride of phosphorus. It is a colourless gas with
54
the unpleasant garlic-like odour. It is poisonous. Phosphine liquefies at
–87 0C
0
and solidified at –134 C. It is very slightly soluble in water.
In the laboratory phosphine may be prepared by boiling white phosphorus with
aqueous alkalies:
P4 + 3NaOH + 3H2O → PH3 + 3NaH2PO2 (Sodium hypophosphite).
Phosphine obtained in such a way is impure. It is passed into an aqueous
solution of HI where phosphorium iodide is formed. Phosphorium iodide obtained
is then heated with KOH or NaOH and pure phosphine is obtained:
PH3 + HI → PH4I,
PH4I + NaOH → PH3 + NaI + H2O.
Other methods of phosphine preparing are:
- by heating phosphorus acid, H3PO3:
∆
4H3PO3 
→ 3H3PO4 + PH3;
- by treating phosphides with water:
Ca3P2 + 6H2O → 3Ca(OH)2 + 2PH3.
Phosphine dissociates at about 450 0C and gives red phosphorus:
4PH3 → P4 + 6H2.
When heated to 150 0C, it burns and produces phosphoric acid:
PH3 + 2O2 → H3PO4.
Phosphine burns in chlorine spontaneously forming PCl3 and PCl5:
PH3 + 3Cl2 → PCl3 + 3HCl,
PH3 + 4Cl2 → PCl5 + 3HCl.
From certain metallic salts solutions, phosphine gives precipitates of metallic
phosphides:
3AgNO3 + PH3 → Ag3P + 3HNO3,
3CuSO4 + 2PH3 → Cu3P2 + 3H2SO4.
Phosphine is neutral to litmus but it reacts with acids forming phosphonium
salts comparable to ammonium salts:
PH3 + HCl → PH4Cl (Phosphonium chloride),
PH3 + HBr → PH4Br (Phosphonium bromide).
The phosphine molecule has a pyramidal AX3E structure like ammonia.
Phosphine is much less soluble in water than ammonia, and it is an exceedingly
weak base. In other words, the equilibrium
450 0 C
PH3(g) + H2O(l) ⇄ PH4+(aq) + OH–(aq)
Is shifted far to the left. Since PH3 is an exceedingly weak base, the phosphonium
ion, PH4+, is almost a strong acid and therefore it reacts with water almost
quantitatively liberating gaseous phosphine:
PH4+(aq) + H2O(l) → H3O+(aq) + PH3(g).
55
Phosphine
Pyramidal AX3E geometry
Phosphonium ion
Tetrahedral AX4 geometry
PHOSPHORUS TRIOXIDE, P2O3 or P4O6
This oxide is the anhydride of phosphorus acid, H3PO3. It is a snow white waxlike solid and has garlic smell. It melts at 23.6 0C and boils at 17.4 0C.
It is prepared by the combustion of yellow or white phosphorus in a limited
quantity of air:
P4 + 3O2 → 2P2O3.
It may be prepared also by the treating phosphorus vapours with N2O at low
pressure at 600 0C:
Low pressure
P4 (Vapours) + 6N2O
→
600 0 C
2P2O3 + 6N2.
P2O3 obtained by both these methods contains some quantity of phosphorus as
impurity. This impurity can be removed by repeated crystallization from carbon
disulfide solution.
When heated at about 400 0C, P2O3 dissociates giving phosphorus dioxide and
red phosphorus:
2P2O3 → 3PO2 + P.
When P2O3 is shaken with ice-water, it gives phosphorus acid:
P2O3 + 3H2O → 2H3PO3.
With hot water this reaction is violent and phosphoric acid with some
phosphine are produced:
2P2O3 + 6H2O → 3H3PO4 + PH3.
PHOSPHORUS TRICHLORIDE, PCl3
PCl3 is a colourless liquid which fumes strongly in air. Its density is 1.57. It
boils at 76 0C.
The phosphorus trichloride is obtained by passing dry chlorine over red
phosphorus which is heated gently in a retort over a water bath. The trichloride
formed distils over and is collected in a receiver cooled in a freezing mixture:
P4 + 6Cl2 → 4PCl3.
A CaCl2 tube is attached to the receiver in order to protect the PCl3 from the
influence of atmospheric moisture. It is purified by distilling it over white
phosphorus.
PCl3 reacts violently with water forming phosphorus acid:
PCl3 + 3H2O → H3PO3 + 3HCl.
Concentrated sulfuric acid reacts with PCl3 forming chlorosulfonic and
metaphosphoric acids:
56
PCl3 + 2H2SO4 → HSO3·Cl + HPO3 + 3HCl.
Chlorosulfonic acid
PCl3 reacts with chlorine forming phosphorus pentachloride:
PCl3 + Cl2 → PCl5.
It reacts with sulfur trioxide SO3 forming phosphorus chloride and sulfur
dioxide:
SO3 + PCl3 → POCl3 + SO2.
When PCl3 reacts with thionyl chloride SOCl2 phosphoryl chloride,
thiophosphoril chloride and phosphorus pentachloride are formed:
SOCl2 + 3PCl3 → POCl3 + PSCl3 + PCl5.
DIPHOSPHORUS TETROXIDE, P2O4 or P4O8
P2O4 is a colourless crystalline solid which sublimes in vacuum at 180 0C.
It is formed by heating P2O3 in a sealed tube at 440 0C or with P2O5 in a sealed
tube at 290 0C:
4P2O3 → 3P2O4 + 2P (red),
P2O3 + P2O5 → 2P2O4.
Under P2O4 treatment with water a mixture of H3PO3 and H3PO4 is obtained:
P2O4 + 3H2O → H3PO3 + H3PO4.
P2O4 can be considered to be a mixed anhydride of H3PO3 and H3PO4.
DIPHOSPHORUS PENTOXIDE, P2O5 or P4O10
P2O5 is a white solid with a garlic smell due to presence of phosphorous
trioxide traces. Pure pentoxide has no smell. Its vapour density corresponds to the
formula P4O10 but it is generally represented by the formula P2O5. It exhibits strong
phosphorescence especially at low temperatures.
It is obtained when white phosphorus burns in excess of air:
P4 + 5O2 → 2P2O5.
It is obtained also when white phosphorus is burned in CO2 at 1000 0C:
P4 + 10CO2 → 2P2O5 + 10CO.
It can be purified by sublimation in oxygen at about 360 0C.
Phosphorous pentoxide gives different acids with different quantities of water,
e.g.,
P2O5 + H2O → 2HPO3 (Metaphosphoric acid),
P2O5 + 2H2O → H4P2O7 (Pyrophosphoric acid),
P2O5 + 3H2O → 2H3PO4 (Orthophosphoric acid).
The main product is definitely orthophosphoric acid.
P2O5 is a dehydrating agent. It removes water from various compounds like
H2SO4, HNO3 etc.:
H2SO4 + P2O5 → 2HPO3 + SO3,
2HNO3 + P2O5 → 2HPO3 + 2NO2 + 1/2 O2.
When reacting with alkalies it gives metaphosphates:
P2O5 + 2NaOH → 2NaPO3 + H2O.
Sodium metaphosphate
57
PHOSPHORUS PENTACHLORIDE, PCl5
PCl5 is an almost colorless crystalline solid with a pungent odour when pure. At
heating, it sublimes below 100 0C and can be melted (m.p. 148 0C) only by heating
under pressure.
Phosphorus pentachloride is prepared by phosphorus trichloride treatment with
dry chlorine (excess):
PCl3 + Cl2 → PCl5.
Phosphorus trichloride is slowly dropped from dropping funnel into a flask
cooled in freezing mixture, while a current of chlorine is passed in. At the end the
excess of chlorine is swept out with a current of carbon dioxide and the flask is
corked. PCl5 which is a yellowish white powder is formed in the flask. Pure PCl5 is,
however, a white powder.
Phosphorus pentachloride dissociates at heating into phosphorus trichloride and
free chlorine:
PCl5 ⇄ PCl3 + Cl2.
It combines with H2O violently forming phosphorus oxychloride, POCl3 (with
unsufficient water) and phosphoric acid, H3PO4 (with excess of water):
PCl5 + H2O → POCl3 + 2HCl,
PCl5 + 4H2O → H3PO4 + 5HCl.
It reacts with compounds containing hydroxyl group and replaces the these
groups by chlorine atoms. For example:
C2H5OH + PCl5 → C2H5Cl + POCl3 + HCl,
Ethyl alcohol
Ethyl chloride
CH3COOH + PCl5 → CH3COCl + POCl3 + HCl,
Acetic acid
Acetyl chloride
SO2(OH)2 + 5PCl5 → SO2Cl2 + 2POCl3 + 2HCl
Sulfuric acid
Sulfurylchloride
Concentrated sulfuric acid forms chlorosulfonic acid with PCl5:
H2SO4 + PCl5 → POCl3 + HCl + ClSO3H.
Oxoacids of phosphorus
Table 3.4 Oxoacids of Phosphorus
Oxidation
State
Name
Formula
Structure
+1
Hypophosphorus
acid
H3PO2
O
O
P H
H
+3
Orthophosphorus
acid
H3PO3
58
+5
Orthophosphoric
acid
H3PO4
+5
Pyrophosphoric
acid
H4P2O7
+5
Triphosphoric acid
H5P3O10
+5
Metaphosphoric
acid
(HPO3)n
HYPOPHOSPHORUS ACID, H3PO2
When phosphorus is boiled with barium hydroxide solution, hypophosphite
formes, which is filtered out:
3Ba(OH)2 + 2P4 + 6H2O → 3Ba(H2PO2)2 + 2PH3.
If sulfuric acid in calculated amounts is added to this barium salt, and then the
solution is filtered, free hypophosphorus acid comes down in the filtrate:
Ba(H2PO2)2 + H2SO4 → BaSO4 + 2H3PO2.
When evaporated, the filtrate leaves the syrupy acid, and when carefully cooled
below 0 0C, it can be crystallized out also.
When strongly heated, the acid decomposes to phosphine:
2H3PO2 → PH3 + H3PO4.
It is a reducing agent. The potential equation of the reduction is:
H3PO2 + 2H2O → H3PO4 + 4H.
Thus, it precipitates silver and gold from their solutions:
4AgNO3 + 2H2O + H3PO2 → 4Ag + 4HNO3 + H3PO4.
Mercuric chloride is reduced to calomel, and finally, to metallic mercury:
4HgCl2 + H3PO2 + 2H2O → 2Hg2Cl2 + 4HCl + H3PO4,
2Hg2Cl2 + H3PO2 + 2H2O → 4Hg + H3PO4 + 4HCl.
H3PO2 can be oxidized by chlorine or iodine:
2Cl2 + H3PO2 + 2H2O → 4HCl + H3PO4.
Hydrogen reduces it to phosphine and atmospheric oxygen oxidizes it to
phosphorus acid. With copper salts it forms cupric hydride, CuH2, or cuprous
hydride, CuH:
CuSO4 + H3PO2 + 2H2O → H3PO4 + H2SO4 + CuH2.
59
These hydrides evolve hydrogen at heating with hydrochloric acid.
Hypophosphorus acid salts are known as hypophosphites. They are soluble in
water as well as in alcohol, and they possess also reducing properties given above.
PHOSPHORUS ACID, H3PO3
Phosphorous acid is a white solid with a melting point of 74 °C. Phosphorus
acid is soluble in water and its solution behaves as a diprotic acid.
Phosphorous acid
It is prepared by adding phosphorus trichloride to water. Since the reaction is
very violent, the addition should be done slowly:
PCl3 + 3H2O → H3PO3 + 3HCl.
Another method of phosphorus acid preparing is by adding PCl3 to anhydrous
oxalic acid. In this case the reaction is not violent and both products are gases:
PCl3 + 3
C OOH
H3PO3 + 3CO + 3CO2 + 3HCl.
C OOH
When heated it gives orthophosphoric acid and phosphine:
4H3PO3 → 3H3PO4 + PH3.
Phosphorus acid is a strong reducing agent. Thus, it reduces mercuric chloride
to mercurous chloride, copper sulfate is reduced to metallic copper and silver
nitrate is reduced to metallic silver:
2HgCl2 + H3PO3 + H2O → Hg2Сl2 + 2HCl + H3PO4,
CuSO4 + H3PO3 + H2O → H3PO4 + Cu + H2SO4,
2AgNO3 + H3PO3 + H2O → 2Ag + 2HNO3 + H3PO4.
It decolourises acidified potassium permanganate, it reduces iodine to
hydroiodic acid and sulfur dioxide to sulfur:
2KMnO4 + 3H2SO4 + 5H3PO3 → K2SO4 + 2MnSO4 + 5H3PO4 + 3H2O,
I2 + H3PO3 + H2O → H3PO4 + 2HI,
2H2PO3 + SO2 → 2H3PO4 + S.
200 0 C
ORTHOPHOSPHORIC ACID, H3PO4
Orthophosphoric acid forms hard, colorless (glassy) rhombic crystals. The
crystals deliquesce in air and form a syrupy liquid. Commercial acid is 82.98 %.
It is soluble in water in all proportions. Its density is 1.87 and it melts at 42.3 0C
(anhydrous).
Orthophosphoric acid is formed when a solution of phosphorus pentoxide in
water is boiled:
P2O5 + 3H2O → 2H3PO4.
60
When red phosphorus is heated with concentrated nitric acid in a flask
containing a reflux condenser orthophosphoric acid is produced:
P4 + 20HNO3 → 4H3PO4 + 20NO2 + 4H2O.
H3PO4 is manufactured from bone ash. Orthophosphoric acid and calcium
sulfate are produced when bone ash which is obtained by the destructive
distillation of bones is treated with concentrated sulfuric acid at heating.
Ca3(PO4)2 + 3H2SO4 → 3CaSO4 + 2H3PO4.
Bone ash
Calcium sulfate being insoluble is filtered off. The filtrate which contains
phosphoric acid is concentrated.
Orthophosphoric acid and its anions all have a tetrahedral AX4 geometry.
Tetrahedral AX4 geometry
Orthophosphoric acid is triprotic acid and forms three types of salts with
alkalis:
H3PO4 + NaOH → NaH2PO4 + H2O;
Sodium dihydrogenphosphate
H3PO4 + 2NaOH → Na2HPO4 + 2H2O;
Sodium hydrogenphosphate
H3PO4 + 3NaOH → Na3PO4 + 3H2O.
Sodium phosphate
Phosphoric
acid
Dihydrogen
phosphate ion
Hydrogen phosphate
ion
Phosphate
ion
When heated to 250 0C, it gives pyrophosphoric acid, H4P2O7:
2H3PO4 → H4P2O7 + H2O.
When heated to 320 0C, it loses one molecule of water giving metaphosphoric
acid, HPO3:
250 0 C
H4P2O7 → 2HPO3 + H2O.
When strongly heated (to red heat), it gives phosphorus pentoxide:
320 0 C
600 − 700 0 C
2HPO3  → P2O5 + H2O.
Orthophosphoric acid liberates hydrobromic and hydroiodic acids from
bromides and iodides respectively:
3NaBr + H3PO4 → Na3PO4 + 3HBr,
3NaI + H3PO4 → Na3PO4 + 3HI.
61
Silver nitrate gives a yellow precipitate of silver orthophosphate with
orthophosphoric acid:
3AgNO3 + H3PO4 → Ag3PO4↓ + 3HNO3.
Barium chloride gives a white precipitate of barium orthophosphate in neutral
or slightly alkaline solutions:
3BaCl2 + 2H3PO4 → Ba3(PO4)2↓ + 6HCl.
Orthophosphoric acid reacts with magnesium salts in the presence of
ammonium chloride and hydroxide giving a white precipitate of ammonium
magnesium phosphate:
H3PO4 + MgCl2 + 3NH3 + 6H2O → NH4MgPO4·6H2O + 2NH4Cl.
PYROPHOSPHORIC (DIPHOSPHORIC) ACID, H4P2O7
Diphosphoric acid is a white crystalline mass with melting point of 61 0C. It is
readily soluble in water.
It is prepared by prolonged heating of orthophosphoric acid at 200-300 0C:
2H3PO4 → H4P2O7 + H2O.
It is obtained also by gently heating a mixture of phosphorus oxy-chloride and
orthophosphoric acid:
5H3PO4 + POCl3 → 3H4P2O7 + 3HCl.
It is a tetra-protic acid and forms two types of salts with alkalies:
2NaOH + H4P2O7 → Na2H2P2O7 + 2H2O;
Sodium dihydrogenpyrophosphate
4NaOH + H4P2O7 → Na4P2O7 + 4H2O.
Sodium pyrophosphate
It gives a white precipitate of silver pyrophosphate:
4AgNO3 + H4P2O7 → Ag4P2O7 + 4HNO3.
Pyrophosphoric acid dissolves in cold water and is slowly converted to
orthophosphoric acid but in hot water and essentially in the presence of nitric acid,
it is easily converted to orthophosphoric acid:
H4P2O7 + H2O → 2H3PO4.
METAPHOSPHORIC ACID, HPO3
It is a hard, sticky mass. It dissolves in water with a cracking noise. It sublimes
at heating.
When (NH4)3PO4 or H3PO4 is heated to about 320 0C, metaphosphoric acid is
formed:
(NH4)3PO4 → HPO3 + 3NH3 + H2O;
H3PO4 → HPO3 + H2O.
It is obtained also by dissolving phosphorus pentoxide in a limited quantity of
water:
62
P2O5 + H2O → 2HPO3.
The formula HPO3 is the empirical formula of metaphosphoric acid, which in
fact is polymeric. It has many different forms depending on the degree of
polymerization. Thus, the general formula is (HPO3)n, where n = 3, 4, 5, 6, ... .
These molecules are rings, such as (HPO3)3 and (HPO3)4, or very long chains,
(HPO3)n. Salts of the metaphosphoric acids, metaphosphates, are used extensively
as water softeners.
Metaphosphates
Metaphosphoric acid gives a white precipitate with silver nitrate solution
forming silver metaphosphate, and it gives also a white precipitate with barium
chloride in neutral or alkaline solutions:
AgNO3 + HPO3 → AgPO3↓ + HNO3;
BaCl2 + 2HPO3 → Ba(PO3)2↓ + 2HCl.
The precipitate of barium phosphate is insoluble in acetic acid.
Aqueous solution of metaphosphoric acid is gradually converted to
orthophosphoric acid at prolonged standing:
HPO3 + H2O → H3PO4.
The alkali metaphosphate NaPO3 forms mixed phosphates when heated with
metal oxides. The mixed phosphates of heavy metals have characteristic colours:
NaPO3 + CuO → NaCuPO4.
Metaphosphoric acid or its salts are not present in the form of simple molecules
but are often polymerized.
Arsenic, Antimony, and Bismuth
Arsenic, symbol As, antimony, Sb, and bithmuth, Bi belong to VA group also.
Arsenic exists in the form of two allotropic modifications: metallic arsenic,
which is a stable metallic look like solid at STP, and yellow arsenic which is not
stable. Metallic arsenic sublimates at heating up.
Antimony has two allotropies also. One is crystalline metallic look like solid,
it’s very brittle. Yellow allotropy of antimony is unstable like that of arsenic.
Bismuth is a glittery reddish-white metal. It’s brittle solid and a currency
conductor.
63
Preparation of the Elements
Arsenic, antimony and bismuth are prepared by their oxides reduction.
Charcoal is commonly used as a reducing agent:
2FeAs + 3O2 → As2O3 + Fe2O3,
2Sb2S3 + 9O2 → 2Sb2O3 + 6SO2↑,
2Bi2S3 + 9O2 → 2Bi2O3 + 6SO2↑.
As2O3 + 3C → 2As + 3CO↑,
Sb2O3 + 3C → 2Sb + 3CO↑,
Bi2O3 + 3C → 2Bi + 3CO↑.
Chemical Properties of the Elements and their Compounds
The arsenic family elements may have –3, 0, +3 and +5 oxidation numbers in
their compounds. The +5oxidation number is the most typical for Sb.
Arsenic and antimony can be oxidized to their corresponding compounds with
the oxidation numbers of these elements +3 and +5. Bismuth is usually oxidized to
Bi+3 compounds, but it can be also oxidized to Bi+5 compounds when it reacts with
a very strong reducing agent:
4As + 3O2 → 2As2O3;
2Sb + 5Cl2 → 2SbCl5;
2Bi + 3S → Bi2S3;
4Bi + 3O2 → 2Bi2O3.
The elements react with alkali and alkaline earth metals giving salts like
compounds (arsenides, antimonides and bismutides) in which they display the
oxidation number of –3:
6Ca + As4 → 2Ca3As2.
These compounds may be transformed into the corresponding hydrides by
treatment with dilute acids:
Mg3As2 + 6HCl → 3MgCl2 + 2AsH3.
Arsenic, antimony and bismuth react with acids in different ways. Arsenic
reacts with hydrochloric acid in the presence of oxygen gas giving arsenic (III)
chloride:
4As + 3O2 + 12HCl → 4AsCl3 + 6H2O.
Antimony does not interact with dilute acids. But it can dissolve in
concentrated hydrochloric acid:
2Sb + 6HCl → 2SbCl3 + 3H2.
Bismuth does not react with dilute hydrochloric and sulfuric acids.
Arsenic, antimony and bismuth can be oxidized by concentrated sulfuric acid
forming corresponding compounds with the oxidation number of +3:
2As + 3H2SO4 → 2H3AsO3 + 3SO2;
2Sb + 3H2SO4 → 2H3SbO3 + 3SO2;
2Bi + 6H2SO4 → Bi2(SO4)3 + 3SO2 + 6H2O.
Concentrated nitric and chloric acids oxidize arsenic and antimony to their
compounds with the oxidation number of +5:
melting
64
2As + 5HClO + 3H2O → 2H3AsO4 + 5HCl;
3As + 5HNO3 + 2H2O → 3H3AsO4 + 5NO;
3Sb + 5HNO3 → 3HSbO3 + 5NO + H2O.
Bismuth is passive upon the treatment with concentrated nitric acid, but it gives
nitrate with dilute one:
Bi + 4HNO3(dil.) → Bi(NO3)3 + NO + 2H2O.
HYDRIDES
The elements form corresponding hydrides with hydrogen. Arsenic hydride
arsin AsH3 and antimony hydride stibin SbH3 are prepared by reacting of their
oxides with acids in the presence of zinc metal:
As2O3 + 6Zn + 12HCl → 6ZnCl2 + 2AsH3 + 3H2O;
Sb2O3 + 6Zn + 12HCl → 6ZnCl2 + 2SbH3 + 3H2O.
Hydride of bismuth, bismuthin BiH3, is prepared by dissolving of a bismuth
alloy with magnesium in hydrochloric acid:
Mg3Bi2 + 6HCl → 3MgCl2 + 2BiH3.
The hydrides of As and Sb can be obtained under the same reaction.
The oxidation number of the elements in their hydrides is –3. All these hydrides
are unstable.
Arsin AsH3 decomposes at 300-400 0C:
2AsH3 ⇄ 3H2↑ + 2As.
The specific test (the Marsh’ reaction) which gives the possibility to detect the
presence of the elements in a substance is based on the ability of these elements
hydrides to decompose at heating up. In this method arsenic is reduced to AsH3
from its compounds. When arsin is then passed through a glass tube filled with a
hydrogen gas and strongly heated, a black-brown precipitate of arsenic metal,
“arsenic mirror”, is formed:
Mg3As2 + 6HCl → 3MgCl2 + 2AsH3;
2AsH3 → 2As + 3H2↑.
SbH3 decomposes at too much low temperature while BiH3 decomposes just at
room temperature.
The hydrides of these elements are strong reducing agents. If silver nitrate
AgNO3 is reacting with arsin, a silver metal forms (Gutcein’s reaction):
AsH3 + 6AgNO3 + 3H2O → H3AsO3 + 6Ag + 6HNO3.
HALIDES
The elements display the oxidation numbers of +3 and +5 in their halides.
Trihalides of arsenic have the similar properties as the phosphorus trihalides.
SbCl3 can dissolve in a small amount of water giving a transparent solution
which gives a precipitate of oxo-chlorides SbOCl and Sb4O5Cl2. Sb3+ ions do not
exist in a solution.
Bismuth trichloride can be reversibly hydrolized to BiOCl:
BiCl3 + H2O ⇄ BiOCl + 2HCl.
65
Trihalides of the elements are able to form complex compounds:
AsI3 + KI → K[AsI4];
SbF3 + KF → K[SbF4].
For all these elements their pentaflourides exist and also a pentachloride for
antimony.
OXIDES
Arsenic, antimony and bismuth can form oxides with the oxidation numbers of
+3 (E2O3) and +5 (E2O5).
Oxidation number +3. Oxides of the elements readily form under the reaction
of these elements with oxygen:
4E + 3O2 → 2E2O3 (E = As, Sb, Bi).
As2O3 has an acidic nature, Sb2O3 has a basic nature, and Bi2O3 displays strong
basic properties. Thus, arsenic trioxide can react with water, alkalis and
hydrohalogens acids:
As2O3 + 3H2O → 2H3AsO3;
As2O3 + 2KOH + 3H2O → 2K[As(OH)4];
Potassium tetrahydroxoarsenite (III)
As2O3 + 8HCl → 2HAsCl4 + 3H2O.
Sb2O3 almost does not dissolve in water, but it can interact with hydrochloric
acid and alkalis:
Sb2O3 + 8HCl → 2H3SbO3 + 3H2O;
Sb2O3 + 2KOH + 3H2O → 2K[Sb(OH)4].
Potassium hexahydroxoantimonite (III)
Bi2O3 does not dissolve in water, it does not react with alkalis but it can readily
interact with acids giving salts:
Bi2O3 + 6HNO3 → 2Bi(NO3)3 + 3H2O.
Oxidation number +5. As2O5 and Sb2O5 are solid substances. Arsenic (V) oxide
can be prepared by arsenic acid dehydration:
2H3AsO4 → As2O5 + 3H2O.
It is impossible to get pure Sb2O5 because it can partially lose oxygen and
transforms into Sb2O4.
As2O5 readily dissolves in water forming arsenic acid:
As2O5 + 3H2O → 2H3AsO4.
Sb2O5 does not dissolve in water but it dissolves in alkalis forming
hexahydroxoantimonates:
Sb2O5 + 2NaOH + 5H2O → 2Na[Sb(OH)6].
When As2O5 and Sb2O5 are melted with alkalis oxoarsenates (V) and
oxoantimonates (V) form with the general formulas of Me+1EO3, Me3+3EO4 and
Me2+2E2O7.
Tetraoxoarsenates Me+1As4 and hexahydroxoantimonates Me+1[Sb(OH)6] are
commonly precipitated from aqueous solutions.
HYDROXIDES
66
Oxidation number +3. Hydroxides of the elements As(OH)3, Sb(OH)3, and
Bi(OH)3 correspond to their oxides As2O3, Sb2O3 and Bi2O3, respectively.
As(OH)3 has the amphoteric properties but acidic ones prevail. It is not educed
in a free state. In aqueous solution it behaves as an acid referred to as arsenous
acid:
H3AsO3 ⇄ HAsO2 + H2O.
As(OH)3 is prepared by the reaction of arsenic trioxide As2O3 with water.
Arsenous acid salts are called arsenites and can be obtained by arsenic trioxide
treating with an alkali solution:
As2O3 + 6KOH → 2K3AsO3 + 3H2O.
Arsenic (III) compounds are able to act as the reducing agents and they can be
oxidized to As (V) compounds. Thus, in a slightly basic medium arsenites may be
oxidized to arsenates by the reaction with free iodine:
K3AsO3 + I2 + 2KOH → K3AsO4 + 2KI + H2O.
Sb(OH)3 also has the amphoteric properties but basic ones prevail. Bi(OH)3
displays only basic properties. Antimony and bismuth hydroxides do not dissolve
in water. They are commonly prepared by the reactions of alkalis solutions with
corresponding salts:
SbСl3 + 3NaOH → Sb(OH)3↓ + 3NaCl;
Bi(NO3)3 + 3NaOH → Bi(OH)3↓ + 3NaNO3.
Antimony(III) and bismuth (III) salts readily hydrolyze giving basic salts:
SbCl3 + 2H2O ⇄ Sb(OH)2Cl + 2HCl;
Bi(NO3)3 + 2H2O ⇄ Bi(OH)2NO3 + 2HNO3.
Antimony (III) compounds are too much weaker reducing agents than arsenic
ones. The oxidation of bismuth (III) compounds is possible in strong basic medium
and with very strong oxidizing agents only:
BiCl3 + Cl2 + 6KOH → KbiO3 + 5KCl + 3H2O.
OXOACIDS
Oxidation number +5. Arsenic acid H3AsO4 corresponds to arsenic (V) oxide
As2O5. It is a weak triprotic acid which can form salts of three types: arsenates
Na3AsO4, dihydroarsenates NaH2AsO4, and hydroarsenates Na2HAsO4.
Antimonic acid H3SbO4 corresponds to antimony (V) oxide Sb2O5. This acid
exists only in aqueous solutions. An attempt to educe it in a pure form leads to
precipitates of unfixed composition Sb2O5·nH2O formation.
Antimonic acid salts (antimonates) correspond to hexahydroantimonic acid
H[Sb(OH)6] which is the hydrated form of antimonic acid HsbO3·3H2O.
It is also impossible to educe the bismuthic acid in the pure state. Bi(V)
compounds are prepared by reacting of strong oxidizing agents with Bi(OH)3 in
concentrated NaOH solution. The composition of the formed compounds
corresponds to the formula Me+1BiO3.
Arsenic and antimonic acids display oxidizing ability in acidic medium only,
and Bi+5 compounds display their oxidizing properties both in acidic and basic
67
mediums:
K3AsO4 + 2KI + H2SO4 → K3AsO3 + I2 + K2SO4 + H2O,
Sb2O5 + 10HCl ⇄ 2SbCl3 + 2Cl2 + 5H2O,
4MnSO4+10KBiO3 + 14H2SO4 → 4KMnO4 + 5Bi2(SO4)3 + 3K2SO4 + 14H2O.
Uses
Most of the nitrogen used in the chemical industry is obtained by the fractional
distillation of liquid air. It is then used to synthesize of ammonia. A wide variety of
important chemical products, including fertilizers, nitric acid, urea, hydrazine, and
amines are prepared from ammonia produced in this manner. In addition, ammonia
compounds are used in the preparation of nitrous oxide (N2O) a colorless gas
commonly known as “laughing gas”. Mixed with oxygen, nitrous oxide is used as
an anesthetic for some types of surgery.
Used as a coolant, liquid nitrogen has found widespread application in the field
of cryogenics. With the recent advent of ceramic materials that become
superconductive at the boiling point of nitrogen, the use of nitrogen as a coolant is
increasing.
Phosphorus is used mainly in the manufacture of safety matches, pyrotechnics,
incendiary shells, smoke bombs, tracer bullets, and smoke screens since it burns
spontaneously and gives dense white fumes. It is used as fertilisers. Phosphates are
used in the production of special glasses, such as those used for sodium lamps.
Bone-ash (calcium phosphate) is used in manufacturing of fine chinaware and for
production of monocalcium phosphate used in baking powder. It is important in the
production of steels, phosphor bronzes, which are the alloys of copper, tin and
phosphorus, and many other products. Na3PO4 is important as a cleaning agent, as
a water softener, and for preventing boiler scale and corrosion of pipes and boiler
tubes. It’s used for pesticides manufacture. Phosphides are used in production of
Holme’s Signals for the ocean ships. Certain compounds of phosphorus, especially
the hypophosphites are used extensively in medicines as tonics. Radioactive
phosphorus (P32) is used in the treatment of leukemia and other blood diseases. Red
phosphorus is used in the preparation of HBr and HI acids in the laboratory.
Phosphates are important to metabolism in both plants and animals. Bones
contain calcium phosphate, Ca3(PO4)2, and the first step in the oxidation of glucose
in the body is formation of a phosphate ester. To provide cattle with phosphate,
dicalcium phosphate, CaHPO4·2H2O, is used as a food supplement. Primary
calcium phosphate, Ca(H2PO4)2, is an ingredient of plant fertilizers.
Phosphorus is a key component of biological molecules such as DNA and
RNA. Phosphorus is a component of bones, and teeth, and many other compounds
required for life. Chronic poisoning of people working unprotected with white
phosphorus leads to necrosis of the jaw ("phossy-jaw").
Arsenic is used in large quantities in the manufacture of glass to eliminate a
green color caused by impurities of iron compounds. A typical charge in a glass
furnace contains 0.5 percent of arsenic trioxide. Arsenic is sometimes added to lead
68
to harden it and is used also in the manufacture of such military poison gases as
lewisite and adamsite. Until the introduction of penicillin, arsenic was of great
importance in the treatment of syphilis. In other medicinal uses, it has been
displaced by sulfa drugs or antibiotics. Lead arsenate, calcium arsenate, and Paris
green are used extensively as insecticides. Certain arsenic compounds, such as
gallium arsenide (GaAs), are used as semiconductors. GaAs is used also as a laser
material. Arsenic disulfide (As2S2), also known as red orpiment and ruby arsenic,
is used as a pigment in the manufacture of fireworks and paints.
Arsenic is poisonous in doses significantly larger than 65 mg (1 grain), and the
poisoning can arise from a single large dose or from repeated small doses, as, for
example, inhalation of arsenical gases or dust. On the other hand, some persons
from the mountains of southern Austria, notably the so-called “arsenic eaters”,
have found that arsenic has a tonic effect and have built up a tolerance for it, so
that they can ingest each day an amount that would normally be a fatal dose. This
tolerance, however, does not protect them against the same amount of arsenic
administered hypodermically.
A reliable test that can detect the presence of minute amounts of arsenic is often
important, because arsenic is a violent poison, yet it is widely used and therefore is
a frequent contaminant. The Marsh test, named for its inventor, the English
chemist James Marsh, supplies a simple method for detecting traces of arsenic so
minute that they would escape discovery in ordinary analysis. The substance to be
tested is placed in a hydrogen generator and any arsenic present is converted to
arsine, (AsH3), which mixes with the evolved hydrogen. If the stream of hydrogen
is heated while it passes through a glass tube, the arsine decomposes, and metallic
arsenic is deposited in the tube. Minute amounts cause an appreciable stain; as little
as 0.1 mg of arsenic or antimony can be detected by using the Marsh test.
When cooling, liquid antimony has the exceptional property of expanding as it
solidifies (water is one of the few other substances with this same property). Thus
it will fill in the crevices of a mold and yield castings of exceptionally sharp
outlines. For this reason, it is used in making type metal; it is also a constituent of
many other alloys, such as Britannia metal, pewter, Babbitt metal (consists of 89 %
tin, 7 % antimony, and 4 % copper) and antimonial lead.
Among important compounds of antimony are tartar emetic, a double tartrate of
antimony and potassium used as a medicinal agent; red antimony sulfide, used on
safety matches and in vulcanizing rubber; glass of antimony, a mixture of
antimony sulfide and oxide, used as a yellow pigment in glass and porcelain; and
butter of antimony, antimony trichloride, used for bronzing steel, as a mordant in
dyeing, and as a caustic in medicine.
Basic bismuth nitrate is also known as bismuth oxynitrate, or bismuthyl nitrate,
and may be used in medicine and cosmetics. Other names for basic bismuth nitrate
are pearl white and Spanish white.
Bismuth expands at solidifying; this unusual property makes it useful for
castings. Some of its alloys have unusually low melting points. One of the most
69
strongly diamagnetic (difficult to magnetize) of all substances, bismuth tends to
turn at right angles to a magnetic field. It is a poor conductor of heat and
electricity, and its electrical resistance further increases in a magnetic field. Due to
this property, it is used in instruments for measuring the strength of such fields.
Bismuth is opaque to X rays and can be used in fluoroscopy.
Bismuth has no biological role. However it has been used for some time as a
medicine (tripotassium dicitratobismuthate) for treatment of stomach upsets. In
combination with antibiotics it is used now for treatment of some stomach ulcers.
Also it is found to be in haemorrhoid creams such as Anusol cream and
Hemocaneas as bismuth oxide and in Anusol ointment as bismuth subgallate.
4. Group VIA. Oxygen, Sulfur, Selenium, Tellurium and
Polonium
There are five chemical elements in VIA group of the periodic table: oxygen O,
sulfur S, selenium Se, tellurium Te and polonium Po.
The atoms of these elements have the following configurations:
– 1s22s22p4
[He]2s22p4
8O
2 2
6 2
4
– 1s 2s 2p 3s 3p
[Ne]3s22p4
16S
2 2
6 2
6
10 2
4
– 1s 2s 2p 3s 3p 3d 4s 4p
[Ar] 4s24p4
34Se
2 2
6 2
6
10 2
6
10 2
4
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p
[Kr] 5s25p4
52Te
2 2
6 2
6
10 2
6
10 14 2
6
10 2
4
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s 6p
[Xe] 6s26p4
84Po
Table 1 lists some physical properties of group VIA elements.
Table 4.1. Properties of Group VIA Elements
Property
Electronic configuration
Melting point, °C
Boiling point, °C
Density, g/cm3
Electronegativity
(Pauling scale)
Covalent radius, Å
Ionic radius E2-, Å
Ionic radius E6+, Å
Oxygen
2
[He]2s 2p
-218.4
-182.96
1.429
Sulfur
4
2
[Ne]3s 3p
115.21
444.6
2.06
Selenium
4
2
[Ar]4s 4p
217.0
685.0
4.81
4
Tellurium
2
[Kr]5s 5p
452
1390
6.25
4
Polonium
[Xe]6s26p4
254
962
9.40
3.44
2.58
2.55
2.10
2.0
0.73
1.36
―
1.04
1.84
0.30
1.17
1.98
0.35
1.37
2.11
0.56
―
―
0.67
Physical Properties and Occurrence
Oxygen, symbol O, is a colorless, odorless, tasteless, slightly magnetic gaseous
element. Gaseous oxygen can be condensed to a pale blue liquid which is strongly
magnetic. Pale blue solid oxygen is produced by compressing the liquid. The
70
atomic weight of oxygen is 15.9994; at atmospheric pressure, the element boils at –
182.96 °C, melts at –218.4 °C, and has a density of 1.429 g/l at 0 °C.
Three structural forms of oxygen are known: ordinary oxygen, containing two
atoms per molecule, formula O2; ozone, containing three atoms per molecule,
formula O3; and a pale blue, nonmagnetic form, O4, containing four atoms per
molecule, which readily breaks down into ordinary oxygen. Ozone (Greek ozein,
“to smell”) is a pale blue, highly poisonous gas with a strong odor. Ozone boils at –
111.9 °C, melts at –192.5 °C, and has a specific gravity of 2.144 g/l.
Oxygen was discovered in 1774 by the British chemist Joseph Priestley and,
independently, by the Swedish chemist Carl Wilhelm Scheele; it was shown to be
an elemental gas by the French chemist Antoine Laurent Lavoisier in his classic
experiments on combustion.
On the Earth, oxygen is more abundant than any other element. Oxygen
composes 21 percents by volume or 23.15 percents by weight of the atmosphere;
85.8 percents by weight of the oceans (88.8 percents of pure water is oxygen); and,
as a constituent of most rocks and minerals, 46.7 percents by weight of the solid
crust of the earth. Oxygen comprises 60 percents of the human body. It is a
constituent of all living tissues. Almost all plants and animals, including all
humans, require oxygen, in the free or combined state, to maintain life.
Sulfur, symbol S, is a tasteless, odorless, light yellow nonmetallic element. Its
atomic number is 16, and its atomic weight is 32.064. When molten sulfur is
slowly cooled, its physical properties change in accordance with the temperature,
pressure, and method of crust formation. Thus sulfur exists in a variety of forms
called allotropic modifications, which consist of the liquids Sλ, and Sµ, and several
solid varieties, the most familiar of which are rhombic sulfur and monoclinic
sulfur. The most stable allotropy of the element is rhombic sulfur, a yellow,
crystalline solid with a density of 2.06 g/cm3 at 20 °C. Rhombic sulfur is slightly
soluble in alcohol and ether, moderately soluble in oils and extremely soluble in
carbon disulfide. When kept at temperatures above 94.5 °C but below 120 °C the
rhombic form changes into monoclinic sulfur consisting of elongated, transparent,
needlelike structures with a density of 1.96 g/cm3 at 20 °C. The temperature at
which rhombic and monoclinic sulfur are in equilibrium, 94.5 °C, is known as the
transition temperature. When ordinary rhombic sulfur is melted at 115.21 °C, it
forms the mobile, pale yellow liquid Sλ, which becomes dark and viscous at 160 °C
forming Sµ. If sulfur is heated almost to its boiling point of 444.6 °C and is then
poured rapidly into cold water, it does not have time to crystallize into the rhombic
or monoclinic state, but forms a transparent, sticky, elastic substance known as
amorphous, or plastic, sulfur, which consists for the most part of supercooled Sµ.
Sulfur also called brimstone, has been known since prehistoric times. It is
mentioned in the Bible and classical records. Due to its flammability, alchemists
regarded sulfur as essential in combustion.
Sulfur ranks 16th in abundance among the elements in the earth's crust and is
found widely distributed in both the free and combined states. In combination it
occurs in many important metallic sulfides, such as lead sulfide, or galena, PbS;
71
zinc blende, ZnS; copper pyrite, (Cu,Fe)S2; cinnabar, HgS; stibnite, Sb2S3; and
iron pyrite FeS2. It combines also with other elements in the form of sulfates such
as barite, BaSO4; celestite, SrSO4; and gypsum, CaSO4·2H2O; and it is present in
the molecules of many organic substances such as mustard, eggs, hair, proteins,
and oil of garlic.
Selenium, symbol Se, is a semimetallic element with an atomic number of 34.
Like sulfur, it exists in several allotropic forms: a brick-red powder; a brownishblack, glassy, amorphous mass called vitreous selenium; red monoclinic crystals of
specific gravity 4.5; and gray, lustrous crystals called gray selenium. Gray
selenium melts at 217 °C, boils at about 685 °C, and has a specific gravity of 4.81.
The atomic weight of selenium is 78.96.
Selenium was discovered in 1817 by the Swedish chemist Baron Jöns Jakob
Berzelius in a sulfuric acid residue. It was so called because it was found in
association with tellurium (Latin tellus, “earth”).
The element occurs in a few selenide minerals, the most common of which is
clausthalite, or lead selenide. It occurs also with free sulfur and in many sulfide
ores; it is generally obtained as a by-product in the refining of copper-sulfide ores.
The yield from by-product sources, however, is insufficient to supply the rapidly
increasing industrial demand for the element. The world's first deposit of
commercial-grade ore was discovered near Baggs, Wyoming, in 1955.
Tellurium (Latin tellus, “earth”), symbol Te, is a silver-white, brittle,
semimetallic element. The atomic number of tellurium is 52. Tellurium melts at
about 452 °C, boils at about 1390 °C, and has a specific gravity of 6.25. The
atomic weight of tellurium is 127.60.
Tellurium was first discovered in 1782 by the German scientist Franz Joseph
Möller von Reichenstein; it was recognized as an element and given its name in
1798 by the German chemist Martin Heinrich Klaproth.
Tellurium ranks about 78th in natural abundance among the elements in the
earth's crust. It occurs in the pure state or is found in combination with gold, silver,
copper, lead, and nickel in such minerals as sylvanite, petzite, and tetradymite.
Occasionally it is found in rocks as tellurite (or tellurium dioxide), TeO2. The slime
from lead and copper refineries and the flue dust from telluride-gold deposits are
the principal commercial sources. It is prepared also by reduction of telluric oxide,
forming a grayish-white, metallic powder.
Polonium, symbol Po, is a rare, radioactive metallic element. Its atomic
number is 84. Polonium melts at about 254 °C, boils at about 962 °C, and has a
specific gravity of 9.4.
The first element being discovered by means of its radioactivity, polonium was
found in pitchblende in 1898 by the French chemist Marie Curie, who named it for
her native country, Poland. Polonium is one of the elements in the uranium-radium
series of radioactive decay, the first member of which is uranium-238. Polonium
occurs in radium-containing ores and is found in isotopic forms with mass numbers
ranging from 192 to 218. Polonium 209 (also called radium-F), the only naturally
occurring isotope, has a half-life of 138 days.
72
Oxygen and its Compounds
Oxygen is a colorless, odorless and tasteless gas at normal conditions. One liter
of water may dissolve 0.04 g of oxygen at STP.
Preparation of Oxygen
In the laboratory oxygen can be prepared by decomposition of potassium
chlorate or potassium permanganate:
2KClO3 → 2KCl + 3O2↑;
2KMnO4 → K2MnO4 + MnO2 + O2↑.
It may be prepared also by heating the alkali metals nitrates:
t
2NaNO3 
→
2NaNO2 + O2↑.
In industry oxygen is prepared by liquefaction and fractional distillation of the
air. The method of oxygen preparation by the electrolysis of water is widely used
also.
Chemical Properties of Oxygen and its Compounds
The electronic structure of diatomic molecule of oxygen gas O2 is:
.
:O
.
O:
Oxygen atom has six electrons on its external electronic shell. So it displays
oxidizing properties only in its reactions with other elements (except fluorine).
Oxygen forms compounds with almost all chemical elements except helium,
neon, and argon. With all elements except halogens, gold and platinum it can react
directly.
Cesium burns spontaneously in the air just at room temperature:
4Cs + O2 → 2Cs2O.
Oxygen reacts with phosphorus when heated up to 60 °C, with sulfur – at the
temperature of 250 °C, with hydrogen – at 300 °C, with carbon (charcoal) – at 700800 °C.
The type of binding in oxides may be ionic one (MgO, Na2O), covalent (N2O5,
P2O5) and with predominately covalent binding (BeO, SiO2, B2O3).
Depending on nature of the element with the positive oxidation number oxides
can be divided into basic (Na2O, MgO), amphoteric (Al2O3) and acidic oxides
(SiO2, P2O5, SO3, Cl2O7). In the reactions with water basic oxides give alkalis,
amphoteric oxides – amphoteric hydroxides, and acidic oxides – acids:
Na2O + H2O → 2NaOH;
Al2O3 + 3H2O → 2Al(OH)3;
SO3 + H2O → H2SO4.
In the consequence of oxides Na2O – MgO – Al2O3 their ability to interact with
water decreases, and in the row P2O5 – SO3 – Cl2O7 this ability increases.
Oxygen forms superoxides MeO2 with potassium, rubidium and cesium and
peroxides Me2O2 with alkali metals and hydrogen (Н2О2).
73
Another type of oxygen compounds are hydroxides. Hydroxides of alkali and
alkaline earth metals are strong bases and they are completely ionized in water
giving the metal ion and hydroxide-ion:
Me+OH−(s) + nH2O ⇄ Me+(aq) + OH−(aq).
In case the bond E – O in a hydroxide has a covalent nature, such hydroxides
considered to be acids and they are ionized in aqueous solutions to give H3O+-ion:
EOH + nH2O ⇄ EO−(aq) + H3O+.
OZONE
Ozone or trioxygen is an allotrope of oxygen consisting of O3 molecules. It is a
pale blue gas with a characteristic odor. When an electric discharge is passed
through gaseous oxygen about 10% of the oxygen is converted to ozone:
3O2(g) → 2O3(g).
The smell of ozone is often noticed during electric storms and in the vicinity of
electric motors. It is an important constituent of the stratosphere.
Ozone condenses to a deep blue liquid at –112 °C. The liquid and even the gas,
if it is not at low pressure or diluted with an inert gas such as nitrogen, can
decompose explosively in a strongly exothermic reaction:
2O3 → 3O2
∆H° = –285 kJ.
Ozone is isoelectronic with the nitrite ion, NO2–, and is an angular AX2E
molecule (see Figure 4.1).
Figure 4.1 Structure of Ozone
We might have expected ozone to be described by the following Lewis
structure:
However, this structure implies bond angles of 60° and equal distances between
all the oxygen atoms, which is not in agreement with the experimentally determined structure (Figure 4.1). Since the covalent radius of oxygen is 66pm, the
length of an oxygen-oxygen single bond is predicted to be 132 pm. The observed
distance of 218 pm between two of the oxygen atoms in ozone is so much larger
than 132 pm that there can be no bond between these two oxygen atoms. Ozone is
represented therefore by the following two resonance structures:
74
Ozone is an extremely powerful oxidizing agent. The only common oxidizing
agent that is stronger is fluorine. Ozone can be used for destroying bacteria in
water by oxidation.
HYDROGEN PEROXIDE, PEROXIDES, AND SUPEROXIDES
Hydrogen peroxide, H2O2, can be thought to be derived from water by
replacement of a hydrogen atom by an OH group. Its structure is shown in Figure
4.2.
Figure 4.2 Structure of Hydrogen Peroxide.
(a) Lewis structure; (b) Dimensions. The two OH bonds in the H2O2
molecule are not in the same plane. In the solid state, there is an angle of
112° between them.
Although the angular geometry at each oxygen atom can be predicted by the
VSEPR model, there is no simple theory that enables one to predict the overall
shape of the molecule – that is, whether it will be planar or nonplanar. In the solid
state it is nonplanar (see Figure 4.2). The O—O bond (147 pm) is somewhat longer
than the single-bond distance of 132 pm predicted from the covalent radius of 66
pm, which suggests that the O–O bond is rather weak and therefore rather easily
broken to give two -OH radicals. Thus H2O2 is a reactive substance.
In the laboratory hydrogen peroxide can be prepared by stirring solid barium
peroxide, BaO2, with a cold aqueous solution of sulfuric acid. Insoluble barium
sulfate precipitates and can be filtered off. If stoichiometric amounts of BaO2 and
sulfuric acid are used, a pure aqueous solution of hydrogen peroxide can be
obtained:
BaO2(s) + H2SO4(aq) → BaSO4(s) + H2O2(aq).
Hydrogen peroxide is a colorless, viscous liquid which boils at 150 °C. Like
water, it is strongly associated by hydrogen bonding. When the pure liquid is
heated, it decomposes rapidly and even explosively in a disproportionation
reaction:
2H2O2(l) → 2H2O(l) + O2(g);
∆H° = –196 kJ.
Normally, it is used as an aqueous solution (for example, 30% by mass for use
in the laboratory and 3% by mass for pharmaceutical use). Such aqueous solutions
decompose very slowly at room temperature, but the decomposition is catalyzed by
many different substances, including Fe(II) salts, manganese dioxide, powdered
platinum, and blood. Hydrogen peroxide is a very weak acid in aqueous solution.
Hydrogen peroxide is an important industrial chemical which has many applications; for example, it is used as a bleaching agent for textiles and for wood pulp
75
and waste paper in paper making. Its usage for bleaching hair is well known. Its
bleaching action is possible due to its strong oxidizing properties.
The oxidation number of oxygen in hydrogen peroxide is –1. It is reduced to
H2O, in which oxygen has an oxidation number of –2. The corresponding halfreaction is:
H2O2(aq) + 2H+(aq) + 2e– → 2H2O(l),
E= +1.77 V.
The large positive value of the standard reduction potential shows that it is a
stronger oxidizing agent than NO3– or MnO4–. Hydrogen peroxide oxidizes Fe2+ to
Fe3+, I– to I2, SO2 (or SO32–) to SO42–, and PbS to PbSO4. The last reaction is used
to restore the original white color to old paintings in which white lead pigment,
Pb3(OH)2(CO3)2, has become converted to dark brown PbS in an urban atmosphere
containing H2S.
Since it can be oxidized to oxygen, in which the oxidation number of oxygen is
zero, H2O2 can behave also as a reducing agent:
H2O2(aq) → O2(g) + 2H+(aq) + 2e–.
For example, H2O2 will reduce MnO4– to Mn2+.
When sodium is heated in a limited supply of air, it forms the expected oxide
Na2O. But in an excess of air it forms the pale yellow peroxide Na2O2, which
contains the peroxide ion, O22–:
Barium behaves in the same way giving barium peroxide, BaO2.
Potassium, rubidium, and cesium give the expected oxides K2O, Rb2O, and
Cs2O in a limited amount of air or oxygen. But when they are heated in the excess
of oxygen, they form the oxides KO2, RbO2, and CsO2. These orange-red compounds contain the superoxide ion, O2–, which has 13 electrons and is a stable,
free-radical ion. It can be represented by two resonance structures,
or by one structure in which the odd electron is shared between the two oxygen
atoms, as in NO:
As the bond order increases from 1 in O22– to 1.5 in O2– and to 2 in the oxygen
molecule, O2, the bond length decreases correspondingly:
Potassium superoxide KO2 is used in self-contained breathing apparatus to
remove moisture and CO2 from the exhaled air and replace them by oxygen. The
reactions are:
2KO2(s) + 2H2O(l) → 2KOH + O2(g) + H2O2(l);
4KO2(s) + 2CO2(g) → 2K2CO3(s) + 3O2(g).
76
Sulfur
Sulfur can be obtained in several different forms. When sulfur is allowed to
crystallize at room temperature from a suitable solvent such as carbon disulfide, it
is obtained in the form of brilliant yellow crystals of orthorhombic sulfur. If
orthorhombic sulfur is heated to 95.5 0C, it changes to another crystalline form,
monoclinic sulfur. However, the rate of this transformation is quite slow, and it is
simpler to obtain monoclinic sulfur by cooling molten sulfur. Monoclinic sulfur
crystallizes from molten sulfur at 119.3 0C in the form of long needles. But
monoclinic sulfur is not stable at room temperature; it slowly changes back to
orthorhombic sulfur. Thus orthorhombic sulfur and monoclinic sulfur are two
allotropes of sulfur. Their names refer to the different structures of their crystals.
Both crystals contain S8 molecules, which consist of a zigzag ring of eight sulfur
atoms. Each sulfur atom has an angular AX2E2 geometry with a bond angle of 1080.
Another form of sulfur is obtained by molten sulfur quick cooling by pouring it
into cold water. A brown rubbery material known as plastic sulfur is obtained. It is
not stable and within a few hours it transformes back into crystalline orthorhombic
sulfur.
Plastic sulfur consists of very long chains of sulfur atoms rather than S8 rings.
When liquid sulfur is heated to about 160 0C, one of the S-S bonds in some of the
rings breaks, and the rings open forming chains of eight sulfur atoms.
S
S
S .
S
S
S
S
S
S
S
S
S
.
S
S
S
S
The sulfur atoms at each end of the chain have only seven valence electrons.
Therefore they have a strong tendency to attach an additional electron and are very
reactive. An S8 chain reacts with S8 ring causing its opening up and thereby
forming a 19-atom chain
S
S
.
S
S
S +
S
S
S
S
S
S
S
. S
S
S
S
.
S
S
S
S
S
S
S
S
S
S
S
S
S
S
S
.S
But the sulfur atoms at the ends of the chain have only seven electrons, so the
process continues, leading to the formation of very long chains of thousands of
atoms. When the temperature is increasing from 160 0C to 190 0C, the chains
become longer and more tangle, and the liquid becomes increasingly thick and
sticky, like maple syrup or molasses. It is described as being very viscous. Plastic
sulfur exhibits properties similar to those of rubber, which also consists of longchain molecules. Long-chain molecules tend to coil up into compact shapes. If
plastic sulfur or rubber is stretched, the coiled-up molecules straighten out a little,
and the material stretches. If the stretching force is removed, the molecules tend to
77
resume their coiled-up structure and the material contracts again (see Figure 4.3).
Coiled-up
molecules
Stretched chains
Figure 4.3 Plastic Sulfur.
Plastic sulfur is an amorphous (noncrystalline) allotrope of sulfur consisting of long chains of sulfur
atoms. These chains have an irregular, disordered arrangement rather than the regular arrangement
which is a characteristic of crystalline substances. The chains tend to coil up and form a tangled mass
When a force is applied, the molecules straighten out a little but coil up again when the force is
removed. So the solid can stretch and contract like rubber
There are still other, less important allotropes of sulfur that contain rings of six,
seven, twelve, and more sulfur atoms.
Preparation of Sulfur
Normally it is not necessary to make sulfur in the laboratory as it is so readily
available. It is found as the native element in nature and extracted by the Frasch
process. This is an interesting process since it means that sulfur can be extracted
from underground without mining it. In the Frasch process underground deposits of
sulfur are forced to the surface using superheated water and steam (160°C, 16
atmospheres, to melt the sulfur) and compressed air (25 atmospheres). This gives
molten sulfur which is allowed to cool in large basins. It is possible to reach sulfur
purity of 99.5%.
The process is energy intensive. Commercial success of this operation depends
upon suitable geological conditions as well as access to cheap water and energy.
Hydrogen sulphide, H2S, is an important impurity in natural gas which must be
removed before the gas is used. This is done by an absorption and regeneration
process to concentrate the H2S, followed by a catalytic oxidation (Claus process)
using porous catalysts such as Al2O3 or Fe2O3.
8H2S + 4O2 → S8 + 8H2O.
Over the years the Claus process has been improved and a modified process can
yield 98% recovery.
In the laboratory, sulfur can be purified by re-crystallization from solutions in
carbon disulphide, CS2. However the resulting crystals are contaminated with
solvent, H2S, and SO2. One good way to purify sulfur is to use a quartz heater (700
78
°C) immersed in liquid sulfur. Carbon impurities decompose forming volatile
materials of solid carbon, which coat the heater. After a week or so, finishing with
a distillation under vacuum, the result is sulfur with a carbon content of about
0.0009 %.
Chemical Properties of Sulfur and its Compounds
Sulfur can exhibit valences of 4 and 6 in addition to the expected valence of 2,
which it has in such compounds as H2S and SCl2. Indeed, sulfur forms many
compounds in which it exhibits these higher valences. These compounds include
the sulfite ion, SO32-, the sulfate ion, SO42-, sulfuric acid, H2SO4, and the fluorides
SF4 and SF6.
Sulfur has the following ground state electronic configuration:
3s
3p
↑↓
↑↓ ↑
↑
Therefore sulfur is expected to form two covalent bonds as it does in H2S.
However, the n=3 shell has five 3d-orbitals also which are not occupied in the
ground state of sulfur:
3s
3p
3d
↑
↑↓
↑↓ ↑
If an electron is transferred from a 3p orbital to one of the 3d orbital we obtain
the excited or valence state configuration:
3s
3p
3d
↑↓
↑
↑
↑
↑
In this valence state there are four electrons in singly occupied orbitals, that is, four
unpaired electrons. These electrons can be used to form four covalent bonds, as in
SO2 and SF4. There is also an unshared pair of electrons in a nonbonding orbital.
Thus SO2 has two double bonds and an unshared pair and is, therefore, an AX2E
type molecule, and it has the expected angular shape with a bond angle of 1200.
Transferring of both 3s electrons to 3d orbitals gives another valence state in which
there are six electrons in singly occupied orbitals:
3s
3p
3d
↑
↑
↑
↑
↑
↑
These six unpaired electrons can be used to form six covalent bonds, as in the
compounds SO3 and SF6. Sulfur trioxide is a planar AX3 molecule and sulfur
hexafluoride, SF6, is an AX6 octahedral molecule.
Sulfur burns in the air forming the gaseous dioxide sulfur(IV) oxide, SO2 at 360
0
C (in oxygen atmosphere – at 280 0C):
S8(s) + 8O2(g) → 8SO2(s).
Sulfur does not react with water under normal conditions.
79
When reacts with metals sulfur displays oxidizing properties giving sulfides in
which sulfur atoms have –2 oxidizing number:
Fe + S → FeS.
Sulfur does not react with hydrogen gas under normal conditions. The
reversible reaction with hydrogen occurs at heating up to 150-200 0C only:
H2 + S ⇄ H2S.
Sulfur behaves as a reducing agent in the reactions with strong oxidizing
agents.
Sulfur reacts with all the halogens at heating. Molten sulfur reacts with molten
chlorine forming disulfur dichloride, S2Cl2. This apparently smells dreadfully.
With excess chlorine and in the presence of a catalyst, such as FeCl3, SnCl4, etc., it
is possible to make a mixture containing an equilibrium mixture of red sulfur(II)
chloride, SCl2, and disulfur dichloride, S2Cl2:
S8 + 4Cl2 → 4S2Cl2(l) [orange],
S2Cl2(l) + Cl2 → 2SCl2(l) [dark red].
Sulfur does not react with dilute non-oxidizing acids.
Sulfur does not react with cold concentrated sulfuric acid, but it interacts with
molten H2SO4 giving sulfur dioxide and H2O:
2H2SO4 + S → 3SO2 + 2H2O.
Dilute nitric acid has no influence on sulfur and concentrated one can oxidize
sulfur to SO42-:
S + 6HNO3 → H2SO4 + 6NO2 + 2H2O.
Sulfur reacts with hot aqueous potassium hydroxide, KOH forming sulphide
and thiosulphate species, that is sulfur is able to be self-oxidized self-reduced:
3S(s) + 6KOH(aq) → 2K2S + K2S2O3 + 3H2O(l).
HYDROGEN SULFIDE AND METAL SULFIDES
Hydrogen sulfide. Unlike water, which is a liquid at room temperature, H2S is
a gas (melting point –85.6 °C, boiling point –60.7 °C). It has a powerful,
unpleasant odor, and it is very poisonous.
H2S is usually prepared by the reaction of dilute acids with iron(II) sulfide:
FeS + 2HCl → FeCl2 + H2S.
AX2E2 angular geometry of H2S
The Lewis structure of H2S is just like that of water. So it is an AX2E2 molecule
and it has the expected angular shape although the bond angle is smaller than in
H2O.
Hydrogen sulfide is soluble in water; it is a weak diprotic acid:
H2S + H2O ⇄ H3O+ + HS–
80
hydrogen sulfide ion;
HS– + H2O ⇄ H3O+ + S2–
sulfide ion.
In hydrogen sulfide sulfur is in its lowest oxidation state ( –2). It is therefore a
good reducing agent; it is usually oxidized to sulfur. The half-reaction of this
oxidation in aqueous solution is the follows:
H2S(aq) → S(s) + 2H+(aq) + 2e–.
If an aqueous solution of H2S is exposed to the air, the H2S is oxidized by the
oxygen of the air and a precipitate of sulfur formes slowly:
2H2S(aq) + O2(g) → 2S(s) + 2H2O(l).
The same reaction occurs when H2S burns in air; sulfur deposits on a cold
surface held near the flame. Some SO2 forms at the same time:
2H2S(g) + 3O2(g) → 2H2O(l) + 2SO2(g).
If H2S is passed into an aqueous solution of bromine, Br2, the red-brown color
of the Br2 disappears as it is reduced to Br–. At the same time a milky precipitate of
finely divided sulfur forms. If we add the equation for the reduction of bromine to
bromide ion:
Br2(aq) + 2e– → 2Br–(aq),
to the equation for the oxidation of H2S to sulfur:
H2S(aq) → S(s) + 2H+(aq) + 2e–,
we obtain the overall equation of the reaction:
H2S(g) + Br2(aq) → 2H+(aq) + 2Br–(aq) + S(s).
Chlorine and iodine can be reduced in the same way.
Hydrogen sulfide decomposes when heated:
H2S → H2 + S.
Hydrogen sulfide reacts vigorously with different oxidizing agents. Free sulfur
is formed when H2S reacts with nitric acid:
H2S + 2HNO3 → S↓ + 2NO2 + 2H2O.
A steam of hydrogen sulfide when passing through the potassium
permanganate solution turns it colorless:
2KMnO4 + 5H2S + 3H2SO4 → 5S↓+ 2MnSO4 + K2SO4 + 8H2O.
H2S reduces chromium from +6 to +3 oxidizing number in the reaction with
potassium dichromate:
K2Cr2O7 + 3H2S + 4H2SO4 → 3S↓+ Cr2(SO4)3 + K2SO4 + 7H2O.
Sulfides and polysulfides
The sulfides of the alkali and alkaline earth metals are colorless solids that are
soluble in water. The sulfides of most other metals are often colored and insoluble.
They are precipitated when H2S is passed through a solution of a salt of the metal:
Zn2+(aq) + H2S(g) → ZnS(s) + 2H+(aq);
Pb2+(aq) + H2S(g) → PbS(s) + 2H+(aq).
We can consider this reaction to occur in two stages. The first is the
ionization of H2S as a weak acid:
H2S(aq) ⇄ 2H+(aq) + S2–(aq),
81
followed by the combination of the metal cation with S2– forming the insoluble
sulfide:
Pb2+(aq) + S2–(aq) → PbS(s).
Removal of sulfide ion by the precipitation of the insoluble metal sulfide forces
the ionization of H2S to the right:
H2S(aq) → 2H+(aq)+S2–(aq) removed as insoluble metal sulfide
The equilibrium between H2S and S2– can be shifted also to the left by using an
excess of H+(aq) and allowing the H2S to release:
H2S(aq) ← 2H+(aq) + S2–(aq).
llowed to release
Excess of the acid
Thus small amounts of H2S can be prepared in the laboratory by adding a dilute
aqueous acid such as hydrochloric acid to a metal sulfide. For example:
FeS(s) + 2H+(aq) → Fe2+(aq) + H2S(g);
FeS(s) + 2HCl(aq) → FeCl2(aq) + H2S(g).
Many insoluble metal sulfides occur in nature and are important ores, for
example NiS, CuS, Ag2S, HgS, and PbS.
Sulfides of metals can be prepared also by the reaction of these metals
hydroxides with hydrogen sulfide:
2NaOH + H2S → Na2S + 2H2O.
Sulfides of alkali and alkaline earth metals are prepared by fuming of these
metals sulfates with charcoal:
Na2SO4 + 2C → Na2S + 2CO2.
As the sulfides of metals are salts of the weak acid they are considerably
hydrolyzed in aqueous solutions. Thus the aqueous solutions of these salts have the
basic medium:
Na2S + H2O → NaHS + NaOH;
2CaS + 2H2O → Ca(HS)2 + Ca(OH)2.
Aluminium sulfide hydrolyzes completely in water:
Al2S3 + 6H2O →2Al(OH)3↓ + 3H2S.
All metals sulfides are strong reducing agents:
3PbS + 8HNO3 → 3PbSO4 + 4H2O + 8NO.
Sulfides can react with oxygen at heating. S2-, SO2 or SO42- can form under
such reactions. The reduction of a metal may take place in some cases:
2PbS + 3O2 → 2PbO + 2SO2;
NiS + 2O2 → NiSO4.
A colorless solution of an alkali or alkaline earth metal sulfide such as Na2S can
dissolve considerable amounts of sulfur giving a bright yellow solution containing
a mixture of polysulfides:
S2– + S → S22–
Disulfide
S + 2S →
Trisulfide
2–
S + 3S →
2–
S32–
S42–
Tetrasulfide
Na2S + (n-1)S → Na2Sn.
Polysulfides have both reducing and oxidizing properties:
82
Na2S2 + SnS → Na2S + SnS2;
4FeS2 + 11O2 → 2Fe2O3 + 8SO2.
Polysulfides are able to be self-oxidized self-reduced:
2FeS2 → 2FeS + 2S.
OXIDES
From its position in the periodic table we expect sulfur to have a valence of 2 as
it does in many of its compounds. We would therefore expect it to form the oxide
SO. Indeed, this oxide is formed in the gas phase when an electric discharge is
passed through sulfur dioxide but, rather surprisingly, it is a very unstable and
reactive substance. The much more stable and better-known oxides of sulfur are
sulfur dioxide, SO2, and sulfur trioxide, SO3.
SULFUR DIOXIDE
When sulfur is heated in air, it ignites at 350 0C and burns with a blue flame
forming sulfur dioxide, SO2. Sulfur dioxide is a colorless gas with a pungent,
choking odor. SO2 forms also by heating metals sulfides:
4FeS2 + 11O2 → 2Fe2O3 + 8SO2.
It condenses to a liquid at –10 0C, at 20 0C it can be liquefied at a pressure of
about 3 atm. It is usually sold as a liquid under pressure in metal cylinders.
The sulfur atom in SO2 molecule has a pair of nonbonding electrons and four
shared pairs of electrons making a total of five pairs, or 10 electrons, in its valence
shell:
.. ..
..
.. .. ..
..
..
: or : O= S =O:
:
:
:
:
O
S
:
O
.
:S + 2 :O
.
.
Like ozone O3, the molecule of SO2 has a triangular structure:
..
S
O
O
Since the sulfur atom in SO2 is in intermediate oxidation state, it can be either
oxidized to the +6 state or to be reduced to the lower oxidation state. Thus SO2 can
act not only as a reducing agent but as an oxidizing agent also. For example, SO2 is
an oxidizing agent in the important reaction with H2S which is used for purification
of natural gas:
2H2S + SO2 → 3S↓ + 2H2O.
In this reaction sulfur in the –2 oxidation state in H2S is oxidized to sulfur in the 0
oxidation state sulfur by SO2, in which sulfur atom is in the +4 oxidation state. At
the same time the sulfur in SO2 is reduced to the 0 oxidation state.
In the presence of strong oxidizing agents SO2 acts as reducing agents.
SO2 can be also self-oxidized:
4SO2 + 8KOH → 2K2SO4 + K2S + 4H2O.
SULFUR TRIOXIDE
83
When sulfur dioxide is heated with oxygen in the presence of finely divided
platinum metal or vanadium pentoxide, V2O5, sulfur trioxide, SO3, forms:
2SO2(g) + O2(g) → 2SO3(g).
Sulfur trioxide condenses to a colorless liquid at 44.5 0C and freezes to
transparent crystals at 16.8 0C.
In the molecule SO3 sulfur uses all 6 of its valence electrons in the formation of
three double bonds to three oxygen atoms:
..
:O
S
O:
..
:O
..
Therefore it has six pairs, or 12 electrons, in its valence shell.
Gaseous sulfur trioxide consists partly of this molecule and partly of a trimeric
molecule, S3O9, which has the ring structure:
...
.O
..
O:
. .
. O.
S
S
:O
O:
..
.. : O
O:
.. S . .
.
:O
O:
..
..
in which each sulfur atom forms six bonds. The liquid consists mainly of S3O9
molecules.
Sulfur trioxide reacts with water giving sulfuric acid, H2SO4:
H2O(l) + SO3(l) → H2SO4(aq),
in which sulfur again forms six bonds:
H
..
O
..
..
:O
S
..
O:
:O:
H
In both (SO3)3 and H2SO4 the sulfur atom forms two double and two single
bonds and has no nonbonding pairs of electrons. Thus the sulfur atom has AX4 type
geometry and therefore the oxygen atoms have a tetrahedral arrangement around
the sulfur atom (Figure 4.4).
Figure 4.4 The Structures of H2SO4 and (SO3)3.
In both molecules there is an AX4 tetrahedral arrangement of four oxygen atoms around each sulfur
atom.
84
Oxoacids of Sulfur
SULFUROUS ACID AND SULFITES
In addition to sulfur dioxide, the most important compounds in which sulfur in
the +4 oxidation state are the metal sulfites. Sulfur dioxide is very soluble in water
(9.4 g L–1 at STP), and it has been generally assumed that sulfurous acid, H2SO3, is
formed in such solutions. The reaction would be:
H2O + SO2 ⇄ H2SO3.
In fact, there is no firm evidence that the molecule H2SO3 exists. Nevertheless, the
anions HSO3– and SO32– are formed in small amounts:
SO2(aq) + 2H2O ⇄ H3O+ + HSO3–;
HSO3– + H2O ⇄ H3O+ + SO32–.
Thus sulfur dioxide behaves as a weak diprotic acid in water, and a solution of
SO2 in water is usually called sulfurous acid.
The sulfite and hydrogen sulfite ions have the following Lewis structures:
and
They both have the expected pyramidal AX3E geometry:
All metals sulfites are soluble in water. Addition of an aqueous acid to a
solution of a sulfite causes the above reactions to proceed from right to left and
SO2 is formed:
SO32–(aq) + 2H+(aq) → SO2(aq) + H2O(l).
When the solution becomes saturated with SO2 it then releases into the gas
phase. The high concentration of H+ (H3O+) and the releasing of the SO2 into the
gas phase both shift the position of the above equilibrium to the right. Small
amounts of SO2 can be prepared in the laboratory in this way.
SO32– is good reducing agents. It is oxidized to sulfate ion, SO42–, by many
oxidizing agents, although, the reaction with O2 is very slow in the absence of a
catalyst such as V2O5. For example, an orange-red solution of Br2 in water is
decolorized by SO2 or by a sulfite solution, because Br2 is reduced to Br–:
Br2(aq) + SO2(aq) + 2H2O(l) → 2Br–(aq) + SO42–(aq) + 4H+(aq).
This reaction is simply the reverse of the reaction that occurs when concentrated
sulfuric acid oxidizes bromide ion to bromine. The above reaction is shifted to the
right because the reaction occurs in dilute aqueous solution and all the sulfuric acid
is ionized to SO42– and H+(aq) – so that there is no un-ionized H2SO4 which is
needed to oxidize Br– to Br2. The high concentration of SO2 also shifts the reaction
equilibrium to the right. Both Cl2 and I2 can be reduced in a similar way.
85
The reducing properties of an aqueous solution of sulfur dioxide make it a
useful bleaching agent. For example, colored compounds in wool, silk, paper, and
straw become colorless when reduced by sulfur dioxide.
Sulfites of active metals can be self-oxidized when heated up to 600 °C giving
the metal sulfate and sulfide:
4K2SO3 
→ 3K2SO4 + K2S.
In aqueous solutions soluble in water sulfites are readily hydrolyzed:
t
Na2SO3 + H2O ⇄ NaHSO3 + NaOH.
When SO2 is passed through the aqueous solutions of metals hydrosulfites the
metal pirosulfites form:
NaHSO3(aq) + SO2(g) → Na2S2O5(aq).
When an aqueous solution of a sulfite is boiled with sulfur, the thiosulfate ion,
S2O32–, forms:
S(s) + SO32–(aq) → S2O32–(aq).
The Lewis structure of this ion is the same as that of the sulfate ion, except that
a sulfur atom replaces one of the oxygen atoms. The replacement of an oxygen
atom by a sulfur atom in a compound is often denoted by the prefix thio-, as in
thiosulfate.
When dilute acid is added to an aqueous solution of a thiosulfate, it decomposes
to sulfur and sulfur dioxide:
S2032–(aq) + 2H30+(aq) → S(s) + SO2(g) + 3H2O(l).
This reaction is the reverse of the reaction by which the thiosulfate ion is formed.
Addition of acid converts the small equilibrium amount of SO32– to H2O and SO2
and thus shifts the equilibrium back to the left.
The two sulfur atoms in thiosulfates are in two ozidation states: +6 and –2. The
presence of sulfur in –2 oxidation state causes the reducing properties of
thiosulfates:
I2 + Na2S2O3 → 2NaI + Na2S4O6.
The most typical reaction for sodium thiosulfate is its reaction with silver
nitrate. Under the reaction the white precipitate of silver thiosulfate Ag2S2O3 forms
which then transforms into black precipitate of silver sulfide Ag2S:
Na2S2O3 + 2AgNO3 → Ag2S2O3↓ + 2NaNO3;
Ag2S2O3 → Ag2SO3↓ + S↓;
Ag2SO3 + S + H2O → Ag2S↓ + H2SO4.
SULFURIC ACID
Pure liquid H2SO4 is easily obtained. It is a colorless liquid with a melting point
of 10.4 °C. The 98 % H2SO4 that is produced industrially and is commonly used in
the laboratory is a mixture of water and sulfuric acid, but since the latter is in a
large excess, it is best regarded as a solution of water in sulfuric acid rather than as
a solution of sulfuric acid in water. It contains H2SO4 molecules and some H3O+
and HSO4–, but no un-ionized H2O molecules. Its properties are quite similar to
those of the 100% pure H2SO4 that consists of H2SO4 molecules only.
86
Industrial preparation of sulfuric acid: contact process
Sulfuric acid can be described as the world's most important industrial chemical;
it is produced in larger quantities than any other substance and has many uses. The
uses of sulfuric acid are so important and so varied that it has been said that a
country's sulfuric acid production is a good measure of its industrial development.
The annual U.S. production, which is the largest in the world, is approximately 40
million tons. About half of this amount is used for the production of phosphate
fertilizers. Other uses include the manufacture of paints, dyes, explosives,
detergents, and synthetic fibers.
The industrial preparation of sulfuric acid involves three reactions that we have
already encountered:
1. The burning of sulfur or a metal sulfide in the air forming SO2:
S + O2 → SO2;
CuS + O2 → Cu + SO2.
2. The oxidation of SO2 to SO3:
2SO2 + O2 Catalyst

→ 2SO3.
3. The combination of SO3 with water giving H2SO4:
SO3 + H2O → H2SO4.
The oxidation of SO2 to SO3 is very slow at ordinary temperatures. To increase
the reaction rate this oxidation is carried out at approximately 400 °C in the
presence of a catalyst, usually vanadium pentaoxide, V2O5. The process is known
as the contact process because the reaction takes place when SO2 and O2 molecules
come into contact on the surface of the solid V2O5 catalyst. The SO 3 is not allowed
to react directly with water, because this is a violent reaction that produces a dense
mist of H2SO4 droplets that is not easily handled. Instead, the gaseous SO3 is
absorbed in 98% H2SO4, and water is added at a controlled rate to keep the
concentration of H2SO4 at approximately 98%. This solution is the acid that is
normally sold as concentrated sulfuric acid. If no water is added, the concentration
of the sulfuric acid rises to 100%, and then a series of polysulfuric acids are
formed:
H2SO4 + SO3 → H2S2O7
Disulfuric acid
H2SO4 + 2SO3 → H2S3O10
Trisulfuric acid
87
A mixture of sulfuric acid and polysulfuric acids is often called oleum or
fuming sulfuric acid.
Properties and reactions of sulfuric acid
Solutions in water. Sulfuric acid is a strong acid and its reaction with water is
highly exothermic:
H2O(l) + H2SO4(l) → H3O+(aq) + HSO4–(aq).
Consequently, the dilution of concentrated sulfuric acid must be carried out with
care. If water is added to the concentrated acid, the heat of the reaction can be
sufficient to raise the temperature of the water to its boiling point and cause drops
of the acid to be thrown violently out of the container. The only safe procedure is
to add the concentrated acid slowly, with constant stirring, to a large amount of
cold water.
Since sulfuric acid has two OH groups, it has two ionizable hydrogen atoms and
it can donate a second proton to another water molecule:
HSO4– + H2O ⇄ H3O+ + SO42–.
The anions formed in these two reactions are the hydrogen sulfate ion, HSO4–,
and the sulfate ion, SO42–. The HSO4– ion is a weak acid and its reaction with water
is not complete. The Lewis structures for the hydrogen sulfate ion and the sulfate
ion are:
and
A solution of sulfuric acid in water contains H3O+, HSO4–, and SO42–. The acidic
properties of such a solution are those of H3O +, and therefore they are the same as
those of an aqueous solution of any strong acid such as hydrochloric acid.
Sulfuric acid as an oxidizing agent
Concentrated aqueous sulfuric acid is a strong oxidizing agent. For example:
• Oxidation of nonmetals: it will oxidize the bromide ion in solid NaBr to
bromine, Br2, and the iodide ion in solid sodium iodide to iodine, I2. It
oxidizes also carbon to CO2 at heating and sulfur to SO2:
2NaBr + 2H2SO4 → Br2 + SO2 + Na2SO4 + 2H2O;
2NaI + 2H2SO4 → I2 + SO2 + Na2SO4 + 2H2O;
C + 2H2SO4 → CO2 + 2SO2 + 2H2O;
S + 2H2SO4 → 3SO2 + 2H2O.
• Oxidation of metals:
- when H2SO4 reacts with not very active metals, SO2 releases:
Cu + 2H2SO4 → CuSO4 + SO2↑ + 2H2O;
- with active metals H2SO4 gives different products:
88
Zn + 2H2SO4 → ZnSO4 + SO2↑ + 2H2O;
Zn + 4H2SO4 → 3ZnSO4 + S↓ + 4H2O;
Zn + 5H2SO4 → 4ZnSO4 + H2S↑ + 4H2O.
Sulfuric acid as a dehydrating agent
Water is completely ionized in solution of sulfuric acid. It behaves as a strong
base:
H2O(l) + H2SO4(l) → H3O+ + HSO4–.
Concentrated 98%
This property makes sulfuric acid a very good dehydrating agent. Gases that do
not react with sulfuric acid, such as O2, N2, CO2, and SO2 can be dried by bubbling
them through concentrated sulfuric acid.
Hydrated salts, such as CuSO4·5H2O, lose their water of crystallization when
stored in a closed container (desiccator) with concentrated sulfuric acid. A small
amount of water vapor is in equilibrium with the hydrated salt, and as this vapor is
absorbed by the sulfuric acid, more of the salt gives up its crystallization water.
Thus slowly all the water is removed:
CuSO4·5H2O ⇄ CuSO4 + 5H2O,
Blue
White
Absorbed by sulfuric acid
(H2O + H2SO4 → H3O+ + HSO4–).
The copper sulfate, that results, is a white powder and is called anhydrous copper
sulfate. It can be reconverted simply to the blue hydrated salt by addition of water.
This reaction can be used as a test for water.
The tendency of sulfuric acid to combine with water is so strong that it will
remove hydrogen and oxygen in a 2:1 atomic ratio from many compounds that do
not contain water in a molecular form. For example, it removes hydrogen and
oxygen as water from carbohydrates and many other organic compounds, leaving
behind a charred residue of carbon. Wood, paper, starch, cotton, and sugar
(sucrose) are all dehydrated in this way:
C12H22O11(s) + 11H2SO4(l) → 12C(s) + 11H3O+ + 11HSO4–.
SULFATES
Sulfuric acid is a diprotic acid which reacts with bases forming two types of
salts, the hydrogen sulfates and the sulfates:
H2SO4(aq) + KOH(aq) → KHSO4(aq) + H2O(l);
KHSO4(aq) + KOH(aq) → K2SO4(aq) + H2O(l).
The overall reaction is
H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2H2O(l).
Metals hydrosulfates and sulfates are ionic compounds containing metal ions
and HSO4– or SO42–. They are generally soluble in water but the sulfates of Ca2+,
Sr2+, Ba2+, Pb2+, and Ag+ are insoluble (Table 5.9). Some insoluble sulfates – such
as CaSO4·2H2O (gypsum), BaSO4 (barite), and PbSO4 - occur as minerals. Barium
sulfate is very insoluble, and its formation as a white precipitate at addition of an
89
acidic aqueous solution of BaCl2 is often used as a test for the presence of sulfate
in a solution:
Ba2+(aq) + SO42–(aq) → BaSO4(s).
Common soluble sulfates include Na2SO4·10H2O, (NH4)2SO4, MgSO4·7H2O
(Epsom salts), CuSO4·5H2O, ZnSO4·7H2O, and some sulfates containing two
different positive ions, such as (NH4)2Fe(SO4)2·6H2O and KAl(SO4)2·12H2O
(alum), which are known as double salts.
The crystals that form when many salts crystallize from water often contain
water molecules and are therefore called hydrates. They have such formulas as
CuSO4·5H2O. These salts are said to be hydrated. The water that is contained in
hydrated salts is called water of crystallization.
Selenium and Tellurium
Selenium and tellurium exist in a few allotropies in free state. For selenium the
most stable is grey selenium. Its crystals are formed with Se8 zigzag chains. Plastic
selenium is a redish-brown powder containing long chains of selenium atoms Se8
which have the irregular arrangement. Red selenium consists of Se8 ring
molecules.
Tellurium has amorphous and hexagonal crystals allotropic modifications. Both
of them contain Te8 chain molecules.
Chemical Properties of the Elements and their Compounds
In the reactions with strong reducing agents, such as hydrogen and metals,
selenium and tellurium display oxidizing properties giving hydrogen or metal
selenides and tellurides:
2K + Se → K2Se;
2Na + Te → Na2Te.
In these compounds Se and Te are in –2 oxidation state like sulfur does.
Under normal conditions hydrogen selenide and telluride H2Se and H2Te are
colorless gases with a sharp odor. They are readily dissolved in water. The aqueous
solutions of H2Se and H2Te are weak acids but they are stronger ones than that of
H2S. They are also stronger reducing agents than H2S.
SELENOUS AND TELLUROUS ACIDS
Selenium and tellurium dioxides SeO2 and TeO2 are polymeric substances.
SeO2 easily dissolves in water forming selenous acid:
SeO2 + H2O → H2SeO3.
TeO2 is insoluble in water but it readily interacts with alkalis forming tellurites:
TeO2 + 2NaOH → Na2TeO3 + H2O.
H2SeO3 can be obtained in a free state as a solid. It can be dehydrated by
heating.
90
Free tellurous acid, H2TeO3, is unknown. Tellurites react with strong acid
forming mTeO2 ⋅ nH2O precipitate.
Selenous and tellurous acids are weak ones. Both these acids as well as the
dioxides are oxidizing agents:
SeO2 + 2SO2 → Se + 2SO3;
H2SeO3 + 2H2S → 2S + Se + 3H2O.
SELENIC AND TELLURIC ACIDS
Selenium trioxide SeO3 is a colourless solid. It reacts vigorously with water
giving selenic acid, H3SeO4:
SeO3 + H2O → H2SeO4.
SeO3 may be prepared by the reaction of SO3 with metals selenates:
Na2SeO4 + SO3 → Na2SO4 + SeO3.
Selenic acid is a white crystalline solid. Its aqueous solution is a strong acid.
When heated selenic acid easily decomposes:
2H2SeO4 → SeO2 + 2O2 + 2H2O.
H2SeO4 is a stronger oxidizing agent than sulfuric acid. It can release chlorine
gas when reacts with hydrochloric acid:
H2SeO4 + 2HCl → H2SeO3 + Cl2 + H2O.
Hot H2SeO4 dissolves silver and gold:
2Au + 6H2SeO4 → Au2(SeO4)3 + 3SeO2 + 6H2O.
Uses
Large amounts of oxygen are used in high-temperature welding torches, in
which a mixture of oxygen and another gas produces a flame of much higher
temperature than is obtained by burning gases in air. Oxygen is administered to
patients whose breathing is impaired and also to people in aircraft flying at high
altitudes, where the poor oxygen concentration cannot support normal respiration.
Oxygen-enriched air is used in open-hearth furnaces for steel manufacture.
Most of the oxygen produced is used to make a mixture of carbon monoxide
and hydrogen called synthesis gas, used for the synthesis of methanol and
ammonia. High-purity oxygen is used also in the metal-fabrication industries; in
liquid form it is of great importance as a propellant for guided missiles and rockets.
The most important usage of sulfur is in the manufacture of sulfur compounds,
such as sulfuric acid, sulfites, sulfates, and sulfur dioxide, all mentioned above.
Medicinally, it has assumed importance due to its widespread usage in sulfa drugs
(commonly known as sulfanilamides) and in many skin ointments. Sulfur is also
employed in the production of matches, vulcanized rubber, dyes, and gunpowder.
In a finely divided state and, frequently, mixed with lime, sulfur is used as a
fungicide on plants. The salt, sodium thiosulfate, Na2S2O3·5H2O, commonly called
hypo, is used in photography for “fixing” negatives and prints. When combined
with various inert mineral fillers, sulfur forms a special cement used to anchor
metal objects, such as railings and chains, in stone. Sulfuric acid is one of the most
91
important of all industrial chemicals because it is employed not only in the
manufacture of sulfur-containing molecules but also in the manufacture of
numerous other materials which do not contain sulfur themselves, such as
phosphoric acid.
Gray selenium conducts electricity; it is a better conductor of electricity in light
than in darkness, the conductivity varying directly with the intensity of light. It is
therefore used in many photoelectric devices. In the form of red selenium or as
sodium selenide the element is used to impart a scarlet red color to clear glass,
glazes, and enamels. It is also used to a great extent as a decolorizer of glass
because it neutralizes the greenish tint produced by iron (ferrous) compounds.
Small amounts of selenium are added to vulcanized rubber to increase its resistance
to abrasion. Sodium selenate is an insecticide used to combat insects that attack
cultivated plants, particularly chrysanthemums and carnations; the insecticide is
scattered around the roots and is carried by the sap throughout the plant. Selenium
sulfide is used in the treatment of dandruff, acne, eczema, seborrheic dermatitis,
and other skin diseases.
Tellurium is used in the manufacture of rectifiers and thermoelectric devices
and in semiconductor research. With other organic substances, it is employed as a
vulcanizing agent in the processing of natural and synthetic rubber; and in
antiknock compounds for gasoline. It is used also to impart a blue color to glass.
Colloidal tellurium is an insecticide, germicide, and fungicide.
Most polonium isotopes disintegrate by emitting alpha particles, the element is
a good source of pure alpha radiation. It is used also in nuclear research with such
elements as beryllium that emits neutrons when bombarded by alpha particles. In
printing and photography equipment, polonium is used in devices that ionize the
air to eliminate accumulation of electrostatic charges.
5. Group VIIA. Fluorine, Chlorine, Bromine, Iodine and
Astatine
There are five chemical elements in group VIIA of the periodic table: fluorine
F, chlorine Cl, bromine Br, iodine I and astatine At.
The atoms of these elements have the following configurations:
– 1s22s22p5
[He] 2s22p5
7F
2 2
6 2
5
– 1s 2s 2p 3s 3p
[Ne] 3s23p5
17Cl
2 2
6 2
6
10 2
5
– 1s 2s 2p 3s 3p 3d 4s 4p
[Ar] 4s24p5
35Br
2 2
6 2
6
10 2
6
10 2
5
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p
[Kr] 5s25p5
53I
2 2
6 2
6
10 2
6
10 14 2
6
10 2
5
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s 6p
[Xe] 6s26p5
85At
Table 5.1 lists some physical properties of group VIA nonmetals.
92
Table 5.1 Properties of Group VIA Elements
Property
Fluorine
5
Electronic configuration [He]2s 2p
Melting point, °C
-219.61
Boiling point, °C
-188.13
Density, g/cm3
1.51
Electronegativity
3.98
(Pauling scale)
0.71
Covalent radius, Å
+
1.33
Ionic radius E , Å

Ionic radius E−, Å
Chlorine
5
2
[Ne]3s 3p
-101.0
-34.05
1.41
Bromine
5
2
[Ar]4s 4p
-7.25
58.78
3.10
Iodine
5
2
[Kr]5s 5p
113.6
185.0
4.94
Astatine
5
[Xe]6s26p5
411.0
299.0

3.16
2.96
2.66
2.2
0.99
1.81
0.26
1.14
1.96
0.39
1.33
2.20
0.50

2.30
0.62
Physical Properties and Occurrence
Fluorine (Latin fluo, “flow”), symbol F, is a chemically reactive, poisonous
gaseous element. The atomic number of fluorine is 9. Fluorine is a pale, greenishyellow gas, slightly heavier than air, poisonous, corrosive, and of penetrating and
disagreeable odor. Its atomic weight is 18.998. Fluorine melts at –219.61 °C, boils
at –188.13 °C, and has a specific gravity of 1.51 in its liquid state at its boiling
point.
The element was first isolated in 1886 by the French chemist Henri Moissan.
Fluorine occurs naturally in the combined form as fluorite (calcium fluoride,
CaF2), cryolite (sodium aluminum fluoride, Na3AlF6), and apatite (calcium
fluoride phosphate, Ca5(PO4)3F). Fluorine occurs also as fluorides in seawater,
rivers, and mineral springs, in the stems of certain grasses, and in the bones and
teeth of animals. It is the 17th element in order of abundance in the crust of the
earth.
Chlorine, symbol Cl, is a greenish-yellow gaseous element. The atomic
number of chlorine is 17. At ordinary temperatures, chlorine is a greenish-yellow
gas that can be readily liquefied under pressure of 6.8 atmospheres, at 20 °C.
Chlorine melts at –101 °C, boils at –34.05 °C at one atmosphere pressure, and has
a specific gravity of 1.41 at –35 °C; the atomic weight of the element is 35.453.
Elementary chlorine was first isolated in 1774 by the Swedish chemist Carl
Wilhelm Scheele, who thought that the gas was a compound; only in 1810 the
British chemist Sir Humphry Davy proved that chlorine was an element and gave it
its present name.
Free chlorine does not occur in nature, but its compounds are common
minerals, and it is the 20th most abundant element in the earth's crust. It is found
mainly as rock salt (common salt, halite, NaCl), carnallite (KCl·MgCl2·6H2O), and
sylvinite (KCl·NaCl). Seawater usually contains from 0.8 to 3.5 % of NaCl.
Bromine, symbol Br, is a poisonous element that is a dark, reddish-brown
liquid at a room temperature. The atomic number of bromine is 35. It melts at –
93
7.25 °C, boils at 58.78 °C, and has a specific gravity of 3.10; the atomic weight of
the element is 79.90.
Bromine is so similar in its chemical properties to chlorine, with which it is
almost invariably associated, that it was not recognized as a separate element until
1826, when it was discovered and isolated by the French chemist Antoine Jerome
Balard.
Bromine is widely distributed in nature. It does not occur in nature as a free
element, but is found in bromide compounds together with chloride compounds.
The usual source is some natural brines and there are few minerals containing
largely of bromine.
Iodine, symbol I, is a chemically reactive element, a blue-black solid at room
temperature. Its atomic number is 53. The atomic weight of iodine is 126.905.
Unlike the lighter halogens, iodine is a crystalline solid at room temperature. The
lustrous, blue-black, soft substance sublimes when heated, giving off a violet vapor
with a stinging odor like that of chlorine. The vapor condenses again rapidly on a
cold surface. Iodine melts at 113.6 °C and boils at 185 °C.
Iodine was first isolated from seaweed residues in 1811 by Bernard Courtois, a
French manufacturer of saltpeter. The discovery was confirmed and announced by
the French chemists Charles Desormes and Nicholas Clement. The nature of the
element was further established in 1813 by the French chemist Joseph Louis GayLussac, who gave iodine its name also.
Iodine is a relatively rare element, ranking about 62nd in abundance on Earth,
but its compounds are widespread in seawater, soil, and rocks. Iodine is obtained
from brines and from Chilean nitrate ores in which it occurs as an impurity. To a
lesser extent, iodine is derived also from sea organisms, such as brown seaweeds,
that concentrate iodine in their tissues.
Astatine (Greek astatos, “unstable”), symbol At, is a radioactive element that
is the heaviest of the halogens. The atomic number of astatine is 85.
Originally called alabamine because of early researches of the element at
Alabama Polytechnic Institute, it was prepared in 1940 by bombarding bismuth
with high-energy alpha particles. The first isotope synthesized had an atomic
weight of 211 and a half-life of 7.2 hours. Subsequently, astatine-210 was
produced and found to have a half-life of about 8.3 hours. Isotopes of astatine with
mass numbers from 200 to 219 have been cataloged, some with half-lives
measured in fractions of a second.
Preparation of the Halogens
Since the halogens occur naturally as halide ions, they are prepared by
oxidation of halide ion to the free halogen. Although in principle Cl– could be
oxidized by fluorine, F2, this method is impractical because fluorine is expensive to
produce and reacts vigorously with water and many other substances. In the
industrial manufacture of chlorine, chloride ion is oxidized by directly removing
electrons in a process called electrolysis, in which an electric current is passed
94
through molten sodium chloride or a concentrated aqueous solution of sodium
chloride:
2Cl–(aq)  → Cl2 + 2e–.
In the laboratory chloride ion can be oxidized to gaseous chlorine, Cl2, by using
a strong oxidizing agent. A suitable oxidizing agent is manganese dioxide, MnO2.
The reaction is carried out in an acidic solution:
MnO2(s) + 2Cl–(aq) + 4H+(aq) → Mn2+(aq) + Cl2(g) + 2H2O(l).
Bromide ion is readily oxidized to bromine by chlorine. Since chlorine is
manufactured in very large amounts and is relatively inexpensive, it is used to
prepare bromine from the bromide ion present in subterranean brines:
Cl2(g) + 2Br–(aq) → Br2(l) + 2Cl–(aq).
Chlorine and steam are passed through the brine. The gaseous bromine that is
formed is carried off from the solution with steam and the excess of chlorine. It is
condensed from this mixture and purified by distillation. Iodine can be obtained
from an aqueous iodide solution by the same method.
In the laboratory both bromine and iodine are prepared by oxidizing bromides
and iodides with manganese dioxide, MnO2, in acidic solution:
MnO2 + 2KBr + 2H2SO4 → MnSO4 + Br2 + K2SO4 + 2H2O,
MnO2 + 2KI + 2H2SO4 → MnSO4 + I2 + K2SO4 + 2H2O.
Since the fluoride ion, F–, is so difficult to oxidize (it is a very weak reducing
agent), the only practical method of fluorine, F2, preparation is by electrolysis of
molten sodium fluoride or a solution of sodium fluoride in liquid hydrogen
fluoride:
2F–(l) → F2(g) + 2e–.
Despite the difficulties of handling fluorine due to its great reactivity, a large
plant can produce about 9 tons of fluorine, F2, a day.
electrolysis
Chemical Properties of Halogens and their Compounds
HALIDE OF NONMETALS IN THEIR HIGHER OXIDATION STATES
In addition to the –1 and +1 oxidation states that arise from the ground state
electron configuration, chlorine, bromine, and iodine have +3, +5, and +7
oxidation states, which correspond to valence state electron configurations in
which one, two, or three electrons have been promoted to d orbitals (see Figure
5.1).
3s
3p
↓↑
↓↑ ↓↑ ↑
↓↑
↓↑ ↑ ↑
Number of
unpaired
electrons
3d
↑
95
Example
compounds
Ground
state
1
ClF,
HOCl
Valence
state
3
ClF3,
HClO2
↓↑
↑ ↑ ↑
↑ ↑
↑
↑ ↑ ↑
↑ ↑ ↑
Valence
state
5
ClF5,
HClO3
Valence
7
HClO4
state
Figure 5.1 Ground State and Valence State Electron Configurations for
Chlorine
Fluorine is such a strong oxidizing agent that it can oxidize most nonmetals to
their highest oxidation states. Thus fluorine reacts with phosphorus giving PF3 and
PF5:
2P + 3F2 → 2PF3;
2P + 5F2 → 2PF5;
and with sulfur giving SF4 and SF6:
S + 2F2 → SF4;
S + 3F2 → SF6.
Fluorine oxidizes even the other halogens, chlorine, bromine, and iodine, to
higher oxidation states. It oxidizes chlorine and bromine to the +1, +3, and +5
oxidation states in ClF, ClF3, and ClF5, and BrF, BrF3, and BrF5; and iodine to IF3
and IF5, and to the +7 state in IF7. Chlorine reacts with phosphorus forming PCl3
and PCl5:
2P + 3Cl2 → 2PCl3;
2P + 5Cl2 → 2PCl5;
but since it is a weaker oxidizing agent than fluorine, it gives chlorides of sulfur
only in lower oxidation states, namely, S2Cl2 and SCl2:
2S + Cl2 → S2Cl2;
S2Cl2 + Cl2 → 2SCl2.
It oxidizes iodine to ICl and ICl3 but not to the +5 and +7 states. These
compounds of fluorine and chlorine with the other halogens are known as
interhalogen compounds.
The halogens combine readily with almost all the other nonmetals giving compounds that consist of covalent molecules.
HYDROGEN HALIDES
Hydrogen reacts with each of the halogens producing the hydrogen halides HF,
HCl, HBr, and HI. For example,
H2(g) + Cl2(g) → 2HCl(g).
The reaction between hydrogen and fluorine is violent under all conditions. The
reaction of hydrogen with chlorine is slow in the dark but becomes explosive in
sunlight or other bright light or at heating to 250 °C. Bromine reacts more slowly
with hydrogen, and iodine reacts still more slowly at a room temperature. A
temperature of at least 200 °C is necessary to obtain a reasonably rapid reaction
96
and in the industrial process of HBr and HI production by direct combination of the
elements a catalyst is used also. If a stream of hydrogen is ignited in air the flame
will continue to burn in an atmosphere of Cl2(g) or Br2(g), producing HCl or HBr.
Thus although all the halogens react with hydrogen forming a hydrogen halide the
reactivity of the halogens towards hydrogen decreases quite markedly from
fluorine to iodine.
The hydrogen halides are all colorless gases that have pungent odors, and all
are very soluble in water. Some of their properties are summarized in Table 5.2.
The bond lengths of the molecules increase steadily from HF to HI as the size of
the halogen atom increases. The melting points and boiling points increase from
HCl to HI as the intermolecular forces become stronger with increasing of atomic
number and increasing of the halogen atom size. Hydrogen fluoride, however, has
unexpectedly high melting and boiling points compared with the other halogen
halides.
Table 5.2 Some Properties of the Hydrogen Halides
Melting Point Boiling Point Bond Length Solubility in Water at 10 °C
(°C)
(°C)
(pm)
(g·L-1)
HF
–83
20
92
Miscible
HCl
–115
–85
127
780
HBr
–89
–67
141
2100
HI
–51
–35
161
2340
NONMETAL HALIDES
When chlorine is heated with carbon, carbon tetrachloride (tetrachloromethane)
is produced:
C(s) + 2Cl2(g) → CCl4(l).
When red phosphorus is gently heated in a slow stream of chlorine, liquid phosphorus trichloride, PCl3, forms:
2P(s) + 3Cl2(g) → 2PCl3(l).
If an excess of chlorine is used, pale yellow solid phosphorus pentachloride,
PCI 5, is produced also.
When a stream of dry chlorine is passed over molten sulfur, an orange-red
liquid forms. It is a mixture of two chlorides of sulfur: S2Cl2, disulfur dichloride,
which is a foul-smelling, orange liquid, and SCl2, sulfur dichloride, which is a deep
red liquid:
2S(s) + Cl2(g) → S2Cl2(l);
S(s) + Cl2(g) → SCl2(l).
Many other nonmetal halides such as PF3, CF4, CBr4, OF2, and ICl can be
obtained similarly. They all are covalent molecular compounds. The shapes of
some nonmetal halides and their Lewis structures are given in Figure 5.2. Table 5.3
summarizes the properties of some nonmetal chlorides. Most of them are gases or
liquids with low boiling points. The boiling points and melting points increase,
97
with increasing of atomic number, as we go down each group, so that a few of the
halides of the heavier elements are solids.
Triangular pyramidal
Tetrahedral
Angular
Figure 5.2 Lewis Structures and Molecular Shapes of Nonmetal Halides
Table 5.3 Properties of Some Nonmetals Chlorides
Period Chloride Melting Point(°C) Boiling Point (°C) Bond Length (pm)
2
CCl4
–23
77
176
NCl3l
–40
71
173
OCl2
–120
2
169
FC
–56
–100
163
3
SiCl4
–68
57
201
PCl3
–94
76
204
SCl2
–122
59
200
Cl2
–101
-35
198
4
GeCl4
–50
83
209
AsCl3
–16
130
216
SeCl2 a
—
—
—
BrCl
-66
5
214
5
SnCl4
-33
114
232
SbCl3
73
223
238
TeCl2
209
327
234
ICl
27
97
232
a
SeCl2 is an unstable compound that has never been isolated in the pure state.
Halides of group VI elements, such as Cl2O, are all angular molecules like H2O.
The halides of group V elements, such as PCl3 and NCl3, are all triangular
pyramidal AX3E molecules, they have the same shape as NH3. The halides of group
IV elements are tetrahedral AX4 molecules, with the group IV atom at the center
and a halogen atom at each of the four corners of the tetrahedron; CCl4 is an
example.
98
The nonmetal halides are typical covalent substances, consisting of covalent
molecules and having relatively low melting points and boiling points. Some of
their physical properties are summarized in Table 5.4.
Table 5.4 Halides of phosphorus, sulfur, and the halogens
Phosphorus
PF3
PF5
PCl3
PCI5
Sulfur
SF4
SF6
S2Cl2
SCl2
Chlorine
ClF
ClF3
ClF5
Bromine
BrF
BrF3
BrF5
Iodine
IF3
IF5
IF7
ICl
ICl3
Phisical condition
Colorless gas
Colorless gas
Colorless liquid
White solid
Boiling Point (°C)
Tb= –95 °C
Tb= –84 °C
Tb= 76 °C
Sublimes at 163 °C
Colorless gas
Colorless gas
Yellow liquid
Red liquid
Tb= –38 °C
Tb= –64 °C
Tb= 138 °C
Tb= 59 °C
Colorless gas
Colorless gas
Colorless gas
Tb= –100°C
Tb= 12 °C
Tb= –14°C
Pale brown gas
Yellow liquid
Colorless liquid
Tb= 20 °C
Tb= 126 °C
Tb= 41 °C
Yellow solid
Colorless liquid
Colorless gas
Red solid
Orange solid
Tb= 28 °C (decomposes)
Tb= 100 °C
Solid sublimes at 5 °C
Tb= 27 °C
Sublimes at 64 °C
Phosphorus pentachloride resembles AlCl3 as it is ionic in the solid state but
covalent in the gas phase. It consists of covalent PCl3 molecules in the gas phase
but is an ionic crystal composed of PCl4+ and PCl6+ ions in the solid state (Figure
5.3).
Figure 5.3 Structure of Phosphorus Pentachlonde
(a) In the gas phase phosphorus pentachloride consists of PCl5 molecules. In the solid state it is an ionic
crystal composed of (b) tetrahedral PCl4+ ions and (c) octahedral PCl6+ ions
99
Most nonmetal halides are rather reactive compounds. Almost all they react
very readily with water giving a hydrogen halide and either an oxoacid or an oxide;
for example,
PCl5(s) + 4H2O(l) → H3PO4(aq) + 5HCl(aq);
SF4(g) + 2H2O(l) → SO2(aq) + 4HF(aq);
ClF5(g) + 3H2O(l) → HClO3(aq) + 5HF(aq).
These reactions between nonmetal halides and water are all Lewis acid-base
reactions. The nonmetal halide is the Lewis acid and water is the Lewis base. The
reaction can be considered to occur in a number of steps, each of which is a Lewis
acid base reaction. The first step in the reaction of PCl5 with water involves the
donation of an unshared electron pair on the oxygen atom of a water molecule into
the valence shell of the phosphorus atom followed by the loss of an HCl molecule:
Repetition of this reaction four more times gives P(OH)5 which loses a water
molecule giving H3PO4.
SF5 is exceptional in that it is not attacked by water even at 500 °C. It appears
that six fluorine atoms completely fill all the space around the sulfur atom.
Therefore there is no space for a water molecule to approach close enough to the
sulfur atom to donate an electron pair into its valence shell. In its inertness to water
SF6 resembles CCl4 and CF4, two other nonmetal halides that do not react with
water. Four chlorine or fluorine atoms completely surround the smaller carbon
atom, leaving no space for attack by a water molecule.
Another property of many nonmetal halides is their ability either to lose or to
gain halide ions forming cations and anions; for example,
PCl5 → PCl4+ + Cl–;
PCl5 + Cl– → PCl6–.
The formation of the ionic solid PCl4+ and PCl6– from gaseous PCl5 molecules
can be regarded as a transfer of a chloride ion from one PCl5 molecule to another.
Other examples include:
BrF3(l) + KF(s) → KBrF4(s);
ICl3(s) + AlCl3(s) → ICI2+AlCl4–(s);
ICl3(s) + KCl(s) → KICl4(s);
ICl(s) + KCl(s) → KICl2(s).
Such anions as BrF4–, ICl4+, and ICl2– are called polyhalide ions. Other anions
of this type are I3–, which is formed by the addition of I– to I2:
KI + I2 → KI3,
and Br3–, which is formed by the similar reaction of Br– with Br2.
Carbon tetrachloride is the best known and most widely used of these nonmetal chlorides. It dissolves many of the other nonmetal halides and many other
substances such as Br2 and I2; therefore it is used as a solvent. Since it dissolves
100
grease, it has been used as a cleaning fluid in dry cleaning. But recently CCl4 has
been replaced by other solvents since it has been found to be carcinogenic.
METAL HALIDES
The halogens react with many metals forming halides. Some typical reactions
are the following:
2Na(s) + Cl2(g) → 2NaCl(s);
Ca(s) + F2(g) → CaF2(s);
Mg(s) + Cl2(g) → MgCl2(s).
The compounds formed in such reactions have quite different physical and
chemical properties from the halides of the nonmetals. Some of the physical
properties of the fluorides and chlorides of the alkali and alkaline-earth metals are
summarized in Table 5.5. They are all colorless crystalline solids with high melting
points.
The metal halides are ionic compounds consisting of oppositely charged ions
held together by electrostatic attraction. The formation of these ionic compounds
can be represented as follows:
Table 5.5 Some Properties of the Alkali and Alkaline Earth Metals Fluorides and
Chlorides
Melting Point
Boiling Point
Solubility in Water at 25
Compound
(°C)
(°C)
°C (g·L–1)
LiF
845
1681
13
NaF
995
1704
40.1
KF
856
1501
1020a
RbF
775
1408
1310
CsF
682
1250
3700a
LiCl
610
1382
850a
NaCl
808
1465
360
KCl
772
1407
350
RbCl
777
1381
940
CsCl
645
1300
1900
BeF2
—
800b
5500
MgF2
1263
2227
013
CaF2
1418
2500
0016
SrF2
1400
2460
012
BaF2
1320
2260
16
BeCl2
405
488
720a
MgCl2
714
1418
550
CaCl2
772
>1600
830a
101
SrCl2
875
1250
560a
BaCl2
1350
—
370a
a
Solubilities of hydrates KF 2H,O Csl H2O LiCl H,O BeCl2 4H,O, CaCl2
6H,O SrCl2 6H2O, BaCl2 2H2O
b
Sublimes (changes directly from solid to gas without forming a liquid)
The bonds in the metal halides are ionic, while those in the nonmetal halides
are covalent.
The halogens as oxidizing agents
All halogens are strong oxidizing agents, but their strength decreases in the
series F2 > Cl2 > Br2 > I2. When an aqueous solution of chlorine, Cl2, is added to a
colorless aqueous solution of sodium bromide, NaBr, the solution becomes orangered rapidly due to the formation of bromine, Br2, in the reaction:
Cl2 + 2Br– → 2Cl– + Br2.
This reaction is an oxidation-reduction reaction, in which bromide ion is
oxidized to bromine,
2Br– → Br2 + 2e–;
while chlorine is reduced to chloride ion:
Cl2 + 2e– → 2Cl–.
Similarly, if an aqueous solution of chlorine is added to a colorless aqueous
solution of an iodide, or if gaseous chlorine is passed into an iodide solution, the
solution becomes brown because of the formation of iodine, I2:
Cl2 + 2I– → 2CP +I2.
Chlorine oxidizes iodide ion to iodine. Iodine is insoluble in water, and the
color is brown due to the triiodide ion, I3–, formed by combination of an iodine
molecule and an iodide ion:
I2(S) + I–(aq) → I3–(aq).
When this solution is shaken with carbon tetrachloride, which is insoluble in
water, molecular iodine is removed from I3– and forms a violet solution in the
carbon tetrachloride.
If bromine is added to an aqueous solution of an iodide, iodine forms according
to the equation:
Br2 + 2I– → 2Br– + I2.
However, if bromine is added to an aqueous solution of a chloride, no reaction
is observed. In other words, the reaction:
Cl2 + 2Br– → 2Cl– + Br2
proceeds from left to right but not from right to left. Whereas chlorine can oxidize
bromide and iodide, bromine can oxidize iodide only. Iodine, the weakest
oxidizing agent among the halogens, oxidizes neither Br– nor Cl–.
Similar reactions occur between the halogens and the hydrogen halides in the
gas phase. For example,
Cl2(g) + 2HI(g) → I2(s) + 2HCl(g).
102
These simple experiments show that the order of oxidizing strengths is Cl2 >
Br2 > I2. Fluorine is a much stronger oxidizing agent than chlorine. Fluorine, for
example, oxidizes water to oxygen:
2F2(g) + 2H2O(l) →→ 4HF(aq) + O2(g).
In contrast, chlorine and bromine dissolve largely unchanged in water, and
iodine is insoluble. Fluorine reacts so vigorously with many substances that it can
be handled in the laboratory only by using special apparatus and taking suitable
precautions.
The order of oxidizing strengths F2 > Cl2 > Br2 > I2 is not unexpected when we
recall that the electronegativities of F, Cl, Br, and I decrease in the same sequence
also. The electronegativity of a halogen is a measure of the tendency of a halogen
atom to attract the electrons of a covalent bond and therefore to acquire a partial
negative charge. It is therefore not the same as the oxidizing strength of a halogen,
which is a measure of the tendency of the diatomic molecule X2 to acquire two
electrons and become two X– ions:
X2 + 2e– → 2X–.
Nevertheless, both the oxidizing strength and the electronegativity of the halogens are related to the tendency of halogen atoms to acquire electrons. Therefore,
not surprisingly, as the electronegativity of X decreases, the oxidizing strength of
X2 decreases also in the order F2 > Cl2 > Br2 > I2.
The ease with which halide ions can be oxidized increases in the order
F– <Cl– <Br– <I–,
which is the order of their strengths as reducing agents.
HYDROGALIDE ACIDS
While dissolving in water all hydrogen halides dissociate like acids giving
H3O+ ion and halide ion:
HHal + H2O ⇄ H3O+ + Hal−.
HCl, HBr and HI are strong acids, but HF is a weak acid because of a high
strenght of the bond between H and F atoms in its molecule.
Hydrofluoric acid can oxidize almost all metals except gold and platinum. But
the reaction of HF with metals occurs mainly on the surface of the metals. The
fluoride of a metal which forms on the metal’s surface prevents the further passing
of the reaction. Hydrofluoric acid destroys glass so it can not be stored in the glass
containers:
SiO2 + 4HF → SiF4 + 2H2O.
Hydrogen fluoride does not react with H2SO4.
HF may be prepared by the reacting of sulfuric acid with calcium fluoride:
CaF2 + H2SO4 → CaSO4 + 2HF.
Hydrogen chloride, HCl, is a colorless gas with a pungent, irritating odor. It can
be obtained by the reaction between hydrogen and chlorine either when a spark is
passed through a mixture of the two gases or when they are strongly heated:
H2(g) + Cl2(g) → 2HCl(g).
103
It can be prepared also by the reaction of sodium chloride with concentrated
sulfuric acid:
NaCl + H2SO4 → NaHSO4 + HCl.
At heating the second step of the reaction takes place:
NaCl + NaHSO4 → Na2SO4 + HCl.
HCl is very soluble in water and is usually encountered in the laboratory as an
aqueous solution called hydrochloric acid. Upon contact with gaseous ammonia,
hydrogen chloride forms white solid ammonium chloride, NH4Cl:
NH3(g) + HCl(g) → NH4Cl(s).
Hydrochloric acid as a strong acid reacts readily with many metals, their oxides
and hydroxides. HCl has no influence on those metals which stand below hydrogen
in the electrochemical series of metals.
Hydrogen bromide and hydrogen iodide are very similar to hydrogen chloride
in their chemical properties. Their aqueous solutions are one-basic acids.
HBr and HI are prepared by the hydrolysis of phosphorous bromide and
phosphorous iodide, correspondingly:
PBr3 + 3H2O → 3HBr + H3PO3;
PI3 + 3H2O → 3HI + H3PO3.
Upon contact with the atmospheric oxygen HBr and HI are easily oxidized
giving free bromine and iodine:
4HBr + O2 → 2Br2 + 2H2O;
4HI + O2 → 2I2 + 2H2O.
The halide ions Hal− except F− are the reducing agents. Their reducing ability
increases from Cl− to Br− and more over to I−:
MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O;
2KMnO4 + 10KI + 8H2SO4 → 2MnSO4 + 5I2 + 6K2SO4 + 8H2O;
6KBr + K2Cr2O7 + 7H2SO4 → 3Br2 + 4K2SO4 + Cr2(SO4)3 + 7H2O.
HBr reduces sulfuric acid to sulfur dioxide SO2, while HI – to hydrogen sulfide,
H2S:
2HBr + H2SO4 → Br2 + SO2↑ + 2H2O;
8HI + H2SO4 → 4I2 + H2S↑ + 4H2O.
Identification reactions of hydrogen halides acids and their salts
Silver nitrate in the presence of dilute nitric acid can precipitate silver halides
from the solutions of the hydrogen halides acids:
AgNO3 + KCl → AgCl↓ + KNO3;
AgNO3 + KBr → AgBr↓ + KNO3;
AgNO3 + KI → AgI↓ + KNO3.
AgF dissolves in water.
AgCl is a white cheese precipitate. It is insoluble in water and nitric acid but it
can be dissolved in the ammonia solution:
AgCl + 2NH4OH → [Ag(NH3)2]Cl + 2H2O.
When the solution formed is acidified the complex compound decomposes and the
AgCl precipitate forms again:
[Ag(NH3)2]Cl + 2HCl → AgCl↓+ 2NH4Cl.
104
AgBr is a pale yellow precipitate which is insoluble in water and nitric acid,
and it is insoluble in the ammonia solution as well.
AgI is a yellow-colored precipitate. It is insoluble neither in water and nitric
acid nor in the ammonia solution. But it dissolves readily in ammonia carbonate
solution.
Oxides and oxoacids of the halogens
The most important compounds containing the halogens in positive oxidation
states are the oxoacids. Oxoacids and oxoanions of the halogens are summarized in
Table 5.6.
Tabel 5.6. Oxoacids and oxoanions of the halogens
Oxidation Formula
O
xoacids
+1
HClO
+3
HCIO2
+5
HClO3
+7
HClO4
Name
Formula
Hypochlorous HBrO
Chlorous
Chloric
HBrO3
Perchloric
HBrO4
Name
Formula
Name
Hypobromous HIO
Hypoiodous
Bromic
Perbromic
lodic
Periodic
Paraperiodic
HIO3
HIO4
H5IO6
A
nions
+1
+3
+5
+7
ClO–
CIO2–
ClO3–
ClO4–
Hypochlorite BrO–
Chlorite
Chlorate
BrO3–
Perchlorate
BrO4–
Hypobromite IO–
Bromate
Perbromate
IO3–
IO4–
IO65–
Hypoiodite
lodate
Periodate
Paraperiodate
OXIDES OF HALOGENS
OF2 can be considered as a fluorine oxide in which fluorine is a more
electronegative element.
It may be prepared by passing fluorine gas through cooled 2% solution of
sodium hydroxide:
2F2 + 2NaOH → 2NaF + OF2 + H2O.
It is used as an oxidizing and fluoriding agent.
All oxides of chlorine, unlike those of phosphorus and sulfur, are somewhat
unstable and explosive. They include Cl2O, ClO2, ClO3, and Cl2O7. The structures
of some of these oxides are shown in Figure 5.4.
Dimensions
Resonance structures
105
Figure 5.4 Structures of Some Oxides of Chlorine
Dichloric oxide Cl2O is a highly endothermic compound, it is very unstable and
easily explodes at heating:
2Cl2O → 2Cl + O2.
Cl2O dissolves readily in water forming hypochlorous acid, HClO:
Cl2O + H2O → 2HClO.
Cl2O is a strong oxidizing agent:
S + 2Cl2O → SO2↑ +2Cl2↑.
It can be also self-oxidized:
3Cl2O + 6NaOH → 2NaClO3 + 4NaCl + 3H2O.
The molecule of dichloric oxide is a polar one and has a triangular structure.
Dichloric oxide is prepared by the reaction of cooled chlorine gas with mercury
(II) oxide:
HgO + 2Cl2 → Cl2O + HgCl2.
ClO2 is prepared by reducing NaClO3 with SO2 in acidic solution:
2NaClO3(aq) + SO2(g) + H2SO4(aq) → 2ClO2(g) + 2NaHSO4(aq).
Chlorine dioxide is a yellow-brown gas that condenses to a red liquid at
–
59 0C. The ClO2 molecule, which has 33 electrons, is an odd-electron molecule.
Like NO2, it is a stable free radical, and it shows even less tendency to dimerize
than NO2. It can be represented by the Lewis structures in Figure 5.4. They
indicate that the odd electron is delocalized over the chlorine and oxygen atoms. It
is an angular AX2E2 molecule.
Chlorine dioxide dissolves easily in water:
2ClO2 + H2O ⇄ HClO2 + HClO3.
Despite its instability and the fact that it can form explosive mixtures with air,
ClO2 is produced in large quantities. It is used as a bleaching agent for cellulose in
the pulp and paper industry and in the manufacture of white flour. It is used also
for water purification.
Chlorine (V) oxide is unknown.
Chloric trioxide, ClO3, is an unstable compound. It is a dark-red oil-like liquid
which can be prepared by the reacting of ozone with ClO2:
ClO2 + O3 → ClO3 + O2.
ClO3 decomposes steadily even under normal conditions:
4ClO3 → 2ClO2 + Cl2 + 4O2.
Dichloric heptoxide Cl2O7 is a colorless liquid with a boiling point of
+83
0
C.
It can be prepared under the reaction of P2O5 with perchloric acid:
106
P2O5 + 6HClO4 → 2H3PO4 + 3Cl2O7.
Cl2O7 is rather stable but it decomposes with explosive when it is heated over
120 0C.
It dissolves in water giving perchloric acid:
Cl2O7 + H2O → 2HClO4.
The following oxides of bromine and iodine are known: Br2O and I2O (the
oxidizing number of halides in these oxides is +1); BrO2 (oxidizing number is +4);
BrO3 (oxidizing number is +6), and I2O5 (oxidizing number is +5). Bromine and
iodine oxides with oxidizing numbers +3 and +7 are unknown.
I2O5 is of the most practical importance of all these oxides. It forms under a
long-time heating of HIO3:
2HIO3 ⇄ I2O5 + H2O.
I2O5 is a strong oxidizing agent:
5CO + I2O5 → I2 + 5CO2.
OXOACIDS OF CHLORINE
The oxoacids of chlorine and their salts are particularly important and have
several useful applications. There are oxoacids corresponding to each of the
positive oxidation states of chlorine (Table 5.6).
HYPOCHLOROUS ACID AND HYPOCHLORITES
Hypochlorous acid is a weak acid in water (Ka= 3.1x10–8). It exists in dilute
aqueous solution only. The pure acid, or even concentrated solutions, cannot be
made because the acid decomposes when dilute solutions are concentrated:
2HOCl(aq) + 2H2O(aq) → 2H3O+(aq) + 2Cl–(aq) + O2(g).
In this reaction hypochlorous acid, in which chlorine is in the +1 oxidation state
oxidizes water to oxygen and reduces itself to Cl– (oxidation number –1). Even
fairly dilute solutions of hypochlorous acid are not completely stable, they
decompose slowly, more rapidly in sunlight.
Hypochlorous acid is formed in equilibrium with molecular chlorine in
solutions of chlorine in water:
Cl2 + 2H2O → HOCl + H3O+ + Cl–.
About a third of the chlorine in a 0.1M solution is converted to HOCl. We can
increase the concentration of HOCl in the solution by removing the Cl– and thus
shifting the equilibrium to the right. Chloride ion can be removed by adding silver
oxide (Ag2O). The Ag+ ion reacts with Cl– forming insoluble AgCl which can be
removed by filtration and the O2– neutralizes the H3O+. The overall equation for
this reaction is:
2Cl2(g) + Ag2O(s) + H2O(l) → 2AgCl(s) + 2HOCl(aq).
Hypochlorites can be made by passing chlorine into a cold solution of a soluble
hydroxide:
Cl2 + 2KOH → KCl + KClO + H2O.
This is essentially the same as the reaction of Cl2 with water. The equilibrium
107
Cl2 + 2H2O → HCl + HOCl
is shifted to the right by removal of HOCl and H+ by reaction with the hydroxide
ion, which converts them to OCl– and H+, respectively. The reactions of chlorine
with water and with OH– are disproportionation reactions in which chlorine in the
zero oxidation state is converted to chlorine in the +1 and –1 oxidation states.
Hypochlorite solutions are prepared on a large scale by electrolyzing of a cold
sodium chloride aqueous solution. The solution is stirred vigorously to mix the
chlorine produced at the anode,
2Cl– → Cl2 + 2e–
with the hydroxide ion produced at the cathode:
2H2O + 2e– → H2 + 2OH–,
so that the reaction
Cl2 + 2OH– → Cl– + ClO– + H2O
takes place.
Hypochlorous acid and hypochlorites are strong oxidizing agents:
OCl–(aq) + 2H+(aq) + 2e– → Cl–(aq) + H2O(l).
Hypochlorites can be self-oxidized:
3KСlO → 2KCl + KClO3.
They are easily decompose and the oxygen gas releasing:
2KClO → 2KCl + O2.
Hypochlorites can act as reducing agents also:
3KClO + 4KMnO4 + 2H2O → 4MnO2 + 3KClO3 +4KOH.
Sodium hypochlorite NaOCI and calcium hypochlorite, Ca(OCl)2, have many
important uses. Calcium hypochlorite is known as bleaching powder, and solutions
of sodium hypochlorite are familiar as household bleach. Hypochlorites are used in
industry to bleach cotton and linen materials and paper pulp. They kill
microorganisms and therefore are used as disinfectants and for water purification.
CHLOROUS ACID AND CHLORITES
Chlorous acid, HClO2, is known in aqueous solution only, in which it behaves
as a weak acid (Ka = 1.1·10–2). Its salts are the chlorites. They are good oxidizing
agents but are not very important. The chlorite ion is an angular AX2E2 molecule. It
can be represented by two resonance structures (see Figure 5.5).
Resonance structures
Dimensions
Figure 5.5 Structure of the Chlorite Ion: An AX2E2 Molecule.
Сhlorous acid is more stable if compared with hypochlorous acid and it is a
weaker oxidizing agent.
It can be prepared from barium peroxide in two stages:
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BaO2 + 2ClO2 → Ba(ClO2)2 + O2;
Ba(ClO2)2 + H2SO4 → BaSO4↓ + 2HClO2.
But chlorous acid is still unstable and readily decomposes even at a room
temperature giving hypochlorous and chloric acids:
2HClO2 → HClO + HClO3.
Chlorites can display their oxidizing ability in acidic solutions only:
5HClO2 → 4ClO2 + Cl− + H+ + 2H2O.
In basic solutions ClO2- ion is stable even at the boiling point.
Chlorites are prepared by the reaction of ClO2 with the solution of a metal
hydroxide:
2ClO2 + 2OH- → ClO2− + ClO3− + H2O.
Chlorates self-oxidizes when heating:
3NaClO2 → 2NaClO3 + NaCl.
In the series of anions ClO2–, ClO3–, ClO4–, the bond orders obtained from the
resonance structures are 1.50, 1.67, and 1.75, and the bond lengths decrease from
157 to 148 and to 144 pm, respectively.
STRENGTHS OF CHLORINE OXOACIDS
The oxoacids of chlorine form an interesting series,
in which the nonbonding pairs are successively replaced by oxygen atoms. This
causes a steady increase in the acid strength from hypochlorous acid, which is the
weakest, to chlorous acid, which is still a weak acid, and to chloric acid and
perchloric acid, which both are strong acids. We have previously seen that sulfuric
acid, H2SO4, is stronger than sulfurous acid, H2SO3, and that nitric acid, HNO3, is
stronger than nitrous acid, HNO2. The increasing number of oxygen atoms in the
series HOCl, HOClO, HOClO2, and HOClO3 increases the effective
electronegativity of the chlorine atom by withdrawing electrons from the chlorine,
thereby increasing the ability of chlorine to withdraw electrons from the OH group.
If the formulas of oxoacids are written in the form of XOm(OH)n the strength of
the acid increases with the value of m (see Table 5.7). According to this
classification, perchloric acid is a stronger acid than sulfuric acid, nitric acid, or
chloric acid. However, all four acids are completely ionized in aqueous solution.
They all are strong acids, and we cannot distinguish between their strengths by
studying their solutions in water. Measurements on the ionization of these acids in
solvents that are less basic than water have shown that perchloric acid is indeed a
stronger acid than nitric acid or sulfuric acid.
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Table 5.7. Strengths of Some Oxoacids XOm(OH)n
Formula
Strength in water
Ka(molL-1)
X(OH)n
Very weak
<10–7
XO(OH)n
Weak 10–2 – 10–4
XO2(OH)n
Strong >1
ClOH
BrOH
IOH
Si(OH)4
B(OH)3
ClO(OH)
PO(OH)3
IO(OH)5
NO(OH)
HPO(OH)2
CO(OH)2
ClO2(OH)
SO2(OH)2
NO2(OH)
XO3(OH)n
Very strong
≫1
ClO3(OH)
CHLORIC ACID AND CHLORATES.
In aqueous solution hypochlorite ion decomposes slowly in two ways. As we
have stated, it decomposes to H3O+, Cl–, and O2 and it self-oxidizes self-reduces
also to chlorate ion, ClO–, and Cl–
3ClO– → 2Cl– + ClO3–.
The rate of this reaction increases if the solution is heated. Chlorates can
therefore be prepared by passing chlorine into a hot hydroxide solution. The
overall equation for the reaction is
3Cl2 + 6KOH → 5KCl + KClO3 + 3H2O.
In the commercial preparation a hot solution of potassium chloride is
electrolyzed with vigorous stirring so that the chlorine produced at the anode mixes
with the hydroxide produced at the cathode. Potassium chlorate has a solubility of
only 3 g in 100 g of water at 0 °C but potassium chloride has a solubility of 28 g in
100 g. The solubility of KClO3 is decreased further by the common-ion effect of
the potassium chloride produced in the reaction. Hence when the solution resulting
from the electrolysis is cooled, potassium chlorate crystallizes.
Chloric acid can be prepared in an aqueous solution by adding the equilibrium
quantity of sulfuric acid to barium chlorate:
Ba(ClO3)2 + H2SO4 → BaSO4 + 2HClO3.
Chloric acid, HClO3, is a strong acid in water, but neither the pure substance
nor concentrated aqueous solutions can be obtained. If solutions of HClO3 are
evaporated, the acid decomposes, sometimes explosively.
The chlorate ion can be described by three equivalent resonance structures. It is
an AX3E molecule and has the expected triangular pyramidal structure (Figure 5.7).
Resonance structures
Figure 5.7 Structure of the Chlorate Ion: An AX3E Molecule.
Chloric acid and chlorates are powerful oxidizing agents. Chlorates form
sensitive explosive mixtures with reducing agents such as carbon, sulfur, and many
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organic compounds, which are converted in exothermic reactions to gaseous
products such as CO2 and SO2:
2KClO3 + 3S → 3SO2↑ + 2KCl;
5KClO3 + 6P → 3P2O5 + 5KCl;
2KClO3 + 3c → 3CO2↑ + 2KCl.
Potassium chlorate finds extensive use in matches, fireworks, and various types
of explosives due to its strength as an oxidizing agent. Sodium chlorate destroys
microorganisms and bacteria in the soil by oxidation. Therefore it is used to kill
weeds and other vegetation.
When potassium chlorate is heated it decomposes in two ways:
(1) The reaction
2KClO3(s) → 2KCl(s) + 5O2(g)
is catalyzed by various substances; manganese dioxide is frequently used.
(2) In the absence of a catalyst, gently heating of potassium chlorate gives potassium perchlorate as the main product:
4KClO3(s) → 3KClO4(s) + KCl(s).
This reaction is an example of a self-oxidation self-reduction, in which potassium
chlorate may be said to oxidize itself. Although almost all perchlorates are rather
soluble in water, the solubility of potassium perchlorate is only 0.75 g in 100 g of
water at 0 °C; therefore it is easily separated from the soluble KCl.
PERCHLORIC ACID AND PERCHLORATES
Perchlorates are strong oxidizing agents but are somewhat less sensitive and
less dangerous than chlorates. Potassium perchlorate and ammonium perchlorate
mixed with carbon or various organic compounds are used as explosives and as
rocket fuels. The solid booster rockets of the space shuttle use a solid propellant
that consists of 70 % ammonium perchlorate, 16 % aluminum powder, and a
polymer known as PBAN (a polybutadiene-acrylic acid-acrylonitrile copolymer),
which binds the mixture into a fairly hard material that looks and feels like a hard
rubber eraser. This propellant is quite safe to prepare and handle. Perchlorates are
used extensively in fireworks and flares.
The perchlorate ion, ClO4–, can be represented by four equivalent resonance
structures. It is an AX4 species and therefore has a tetrahedral structure (see Figure
5.8). It is isoelectronic with the tetrahedral PO43– and SO42– ions. The observed
bond length (144 pm) is much smaller than the sum of the single bond, covalent
radii of chlorine and oxygen, which is 165 pm. This bond length is consistent with
the bond order of 1.75 implied by the resonance structures.
Dimensions
One of 4 equivalent resonance structures
Figure 5.8 Structure of the Perchlorate Ion: An AX4 Molecule
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Anhydrous perchloric acid, HClO4, can be prepared by distilling a mixture of
potassium perchlorate and concentrated sulfuric acid under reduced pressure:
KClO4 + H2SO4 → HClO4 + KHSO4.
It is a somewhat hazardous substance as it is a very powerful oxidizing agent
that explodes violently at contact with organic matter such as wood and cloth.
Perchloric acid is a strong acid in water and its dilute aqueous solutions are much
less reactive than the pure acid and are safe to handle. Aqueous solutions of
perchloric acid and perchlorates are potentially good oxidizing agents but the
reactions are usually quite slow at ordinary temperatures.
Although all the halogen oxoacids and their anions are good oxidizing agents,
their oxidizing ability and, particularly, the rates at which they react depend very
much on the conditions. Their oxidizing strengths do not simply increase with
increasing number of oxygen atoms.
OXOACIDS OF BROMINE AND IODINE
The oxoacids of bromine are HBrO, HBrO3, and HBrO4. They resemble the
corresponding chlorine oxoacids. Bromous acid, HBrO2, is not known.
Hypobromous acid, HBrO, can exist in dilute aqueous solutions only. Like
hypochlorous acid hypobromous acid may be prepared by the interaction of
bromine with water:
Br2 + H2O ⇄ HBr + HbrO.
HBrO is a weak acid and it acts as a oxidizing agent.
Hypobromites (salts of hypobromous acid) can be obtained by passing bromine
through a cold solution of a soluble hydroxide:
Br2 + 2NaOH → NaBrO + NaBr + H2O.
Hypobromites are unstable and exist in dilute solutions only.
Bromic acid, HBrO3, is also unstable and exists in dilute aqueous solutions. Its
chemical properties are very much similar to that of chloric acid.
Salts of bromic acid, bromates, can be prepared by passing bromine into a
boiling hydroxide solution:
3Br2 + 6KOH → KBrO3 + 5KBr + 3H2O.
Bromites does not display oxidizing ability in basic and neutral medium. But
they act as oxidizing agents in acidic solutions:
6FeSO4 + KBrO3 + 3H2SO4 → 3Fe2(SO4)3 + KBr + 3H2O.
If potassium bromide, for example, is used as a reducing agent, free bromine
forms:
5KBr + KBrO3 + 6HCl → 3Br2 + 6KCl + 3H2O.
Perbromic acid, HBrO4 is known in dilute aqueous solutions only. It is a
stronger oxidizing agent than perchloric acid. Salts of perbromic acid, perbromates,
can be prepared by oxidizing of bromates in basic solution:
NaBrO3 + F2 + 2NaOH → NaBrO4 + 2NaF + H2O.
The oxoacids of iodine are hypoiodous acid, HIO, iodic acid, HIO3, and the
periodic acids, HIO4 and H5IO6. Sodium iodate occurs naturally in small amounts
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in sodium nitrate, deposits of which are found in certain desert regions of Chile.
Iodine is obtained from sodium iodate by reduction with sulfur dioxide or with
sodium hydrogen sulfite, NaHSO3:
2IO3–(aq) + 5SO2(g) + 4H2O(l) → I2(s) + 8H+(aq) + 5SO42–(aq).
Iodine is insoluble in water but dissolves in a solution of hydroxide ion with the
formation of hypoiodite ion, IO–. This reaction is quite similar to the reaction of
Cl2 with OH–:
I2(s) + 2OH–(aq) → IO–(aq) + I–(aq) + H2O(l).
However, hypoiodite ion is considerably less stable than hypochlorite ion, and
it self-oxidizes self-reduces rather rapidly:
3IO–(aq) → IO3–(aq) + 2I–(aq).
This reaction is further accelerated by heating the solution. Thus when iodine is
dissolved in warm hydroxide solution, iodate and iodide form:
3I2(s) + 6OH–(aq) → 5I–(aq) + IO3–(aq) + 3H2O(l).
We have seen that chlorine reacts similarly giving chlorate and chloride in hot
NaOH solution.
Iodate ion, IO3–
Iodic acid, HIO3
Figure 5.9 Structures of Iodic Acid and the Iodate Ion.
(six equivalent resonance structures)
dimensions
dimensions
Paraperiodic acid, H5IO6
Paraperiodate ions, IO65–
(four equivalent dimensions resonance structures) Metaperiodate ion, IO4–
Figure 5.10 Structures of Periodic Acid and the Periodate Ion
Whereas chloric acid is unstable and cannot be isolated in a pure form, iodic
acid is a stable, white crystalline solid. It can be made by oxidizing iodine with
nitric acid:
I2(s) + 10HNO3(aq) → 2HIO3(aq) + 10NO2(g) + 4H2O(l).
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Like the chlorate and bromate ions, iodic acid and the iodate ion have a triangular pyramidal AX3E geometry (Figure 5.9).
Periodic acid and periodates differ from perchloric acid and perbromic acid as
they exist in two forms (see Figure 5.10). Periodic acid is a solid which, when
crystallized from water, has the formula H5IO6. This form is known as
paraperiodic acid. It is a weak acid (see Table 20.6). If it is heated to 80 °C under
vacuum, it loses water giving HIO4, which is known as metaperiodic acid:
H5IO6 → HIO4 + 2H2O.
Salts of both forms of the acid are known, such as NaIO4, Na2H3IO6, and
Ag5IO6.
Uses
Fluorine compounds have many applications. The chlorofluorocarbons,
odorless and nonpoisonous liquids or gases such as Freon, are used as a dispersing
agent in aerosol sprays and as a refrigerant. In 1974, however, some scientists
suggested that these chemicals destroy the earth's ozone layer when reached the
stratosphere. With confirmation of these findings by the late 1980s, the production
of these chemicals began to be phased out. Another chemical, Teflon, fluorine
plastic that is very resistant to most chemical action, is widely used to make such
products as motor gaskets and dashboard accessories in the automobile industry.
Teflon is also used as a coating on the inner surface of frying pans and other
kitchen utensils to reduce the need for fat in cooking. Many organic fluorine
compounds developed during World War II (1939–1945) showed extensive
commercial potential. For example, the liquid fluorinated hydrocarbons derived
from petroleum are useful as highly stable lubricating oils. Uranium hexafluoride,
the only volatile compound of uranium, is used in the gaseous diffusion process to
provide fuel for atomic power plants.
Drinking water containing excessive amounts of fluorides causes tooth enamel
to become brittle and to chip off, leaving a stained or mottled effect. The proper
proportion of fluorides in drinking water, however, has been found to greatly
reduce tooth decay.
Chlorine is used for bleaching paper pulp and other organic materials,
destroying germ life in water, and preparing bromine, tetraethyl lead, and other
important products. It has an irritating odor and in large concentration is
dangerous; it was the first substance used as a poison gas in World War I (1914–
1919).
Bromine has been used in the preparation of certain dyes and of dibromoethane
(commonly, ethylene bromide), a constituent of antiknock fluid for leaded
gasoline. Bromides are used in photographic compounds and in natural gas and oil
production.
Iodine is medicinally very important because it is an essential trace element,
present in a hormone of the thyroid gland that is involved in growth-controlling
and other metabolic functions. Without iodine, stunted growth and conditions such
as goiter can result. Thus in areas where iodine is not sufficiently abundant
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naturally, iodine-containing salt serves to make up the deficit. In medicine, iodinealcohol solutions and iodine complexes have been used as antiseptics and
disinfectants. Radioisotopes of iodine are used in medical and other fields of
research. More broadly, various iodine compounds find use in photography, the
making of dyes, and cloud-seeding operations. In chemistry, various iodine
compounds serve as strong oxidizing agents, among other uses.
Astatine is highly carcinogenic. Mammary and pituitary tumors have been
induced in laboratory animals by a single injection.
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Contents
Introduction .............................................................................................................. 4
1. Group IIIA. Boron, Aluminum, Gallium, Indium, and Thallium ......................... 5
Physical Properties and Occurrence ..................................................................... 5
Preparation of the Elements ................................................................................. 7
Chemical Properties of Boron and its Compounds............................................... 8
Chemical Properties of Aluminum and its Compounds ..................................... 13
Uses .................................................................................................................... 15
2. Group IVA. Carbon, Silicon, Germanium, Tin and Lead................................... 16
Physical Properties and Occurrence ................................................................... 17
Carbon and its Compounds ................................................................................ 19
Silicon and its Compounds ................................................................................. 25
Compounds of Lead ........................................................................................... 31
Uses .................................................................................................................... 34
3. Group VA. Nitrogen, Phosphorus, Arsenic, Antimony and Bismuth................. 35
Physical Properties and Occurrence ................................................................... 35
Nitrogen and its Compounds .............................................................................. 37
Phosphorus and its Compounds.......................................................................... 52
Arsenic, Antimony, and Bismuth ....................................................................... 63
Uses .................................................................................................................... 68
4. Group VIA. Oxygen, Sulfur, Selenium, Tellurium and Polonium ..................... 70
Physical Properties and Occurrence ................................................................... 70
Oxygen and its Compounds ............................................................................... 73
Sulfur .................................................................................................................. 77
Selenium and Tellurium ..................................................................................... 90
Uses .................................................................................................................... 91
5. Group VIIA. Fluorine, Chlorine, Bromine, Iodine and Astatine ........................ 92
Physical Properties and Occurrence ................................................................... 93
Preparation of the Halogens ............................................................................... 94
Chemical Properties of Halogens and their Compounds .................................... 95
Uses .................................................................................................................. 114
116