Download Module-2-s-and-d-elements - Львівський національний медичний

DANYLO HALYTSKY LVIV NATIONAL MEDICAL UNIVERSITY
DEPARTMENT of GENERAL, BIOINORGANIC, PHYSICAL and
COLLOIDAL CHEMISTRY
V.V. Ogurtsov, O.V. Klenina, O.I. Marshalok, I.I. Myrko
CHEMISTRY OF s- AND d-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.І. Маршалок, І.І. Мирко
ХІМІЯ s- І d-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.
3
Inorganic chemistry
Chemical elements are substances that cannot be decomposed or
broken into more elementary substances by ordinary chemical means.
Elements were at one time believed to be the fundamental substances but
they are known to consist of a number of different elementary particles
now. More than 100 chemical elements are known to exist in the universe,
although several of these, the so-called transuranium elements, have not
been found in nature, and can be produced artificially only.
Chemical elements are classified as metals and nonmetals. The atoms
of metals are electropositive and combine readily with the electronegative
atoms of the nonmetals. A group of elements called metalloids, which are
intermediate in properties between the metals and the nonmetals, is
sometimes considered a separate class. When the elements are arranged
in the order of their atomic numbers (a number proportional to the net
positive charge on the nucleus of an element atom), elements of similar
physical and chemical properties occur at specific intervals. These groups
of elements with similar physical and chemical properties are called
families, examples of which are the alkali metals, alkaline earth metals, rare
earth elements, halogens, and the noble gases.
When two atoms have the same atomic number but different atomic
mass, they are said to be isotopes. Many natural isotopes are known for
some elements, whereas other elements occur in one isotopic form only.
Hundreds of synthetic isotopes have been made. Some natural isotopes,
and many synthetic ones, are unstable.
1. Hydrogen
Hydrogen (Greek for “water former”), symbol H, is a reactive, colorless,
odorless, and tasteless gaseous element. The atomic mass of hydrogen is
1.00797. The atomic number of hydrogen is 1. The element is usually
classed in group 1 (or IA) of the periodic table. Electronic configuration is
1
1s .
Hydrogen was confused with other gases until the British chemist Henry
Cavendish demonstrated in 1766 that it was evolved by the action of
sulfuric acid on metals and also showed later that it was an independent
substance which combined with oxygen to form water. The British chemist
Joseph Priestley named the gas “inflammable air” in 1781, and the French
chemist Antoine Laurent Lavoisier renamed it into hydrogen.
4
Physical Properties and Occurrence
Like most gaseous elements, hydrogen is diatomic (its molecules
contain two atoms), but it dissociates into free atoms at high temperatures.
Hydrogen has a lower boiling point and melting point than any other
substance except helium; hydrogen melts at –259.2 °C and boils at
–252.77° C. At 0 °C and 1 atmosphere pressure (STP), hydrogen is a gas
with a density of 0.089 g/L. Liquid hydrogen, obtained firstly by the British
chemist Sir James Dewar in 1898, is colorless (but light blue in thick layers)
with the specific gravity 0.070. When allowed to evaporate rapidly under
reduced pressure, it freezes into a colorless solid.
Hydrogen is known to exist in three isotopic forms. The nucleus of each
atom of ordinary hydrogen is composed of one proton. Deuterium,
presenting in ordinary hydrogen to the extent of 0.02 percent, contains one
proton and one neutron in the nucleus of each atom and has an atomic
mass of two. Tritium, an unstable, radioactive isotope, contains one proton
and two neutrons in the nucleus of each atom, and has an atomic mass of
three.
Free hydrogen is found only in very small traces in the atmosphere, but
solar and stellar spectra show that it is abundant at the sun and other stars,
and is, in fact, the most common element in the Universe. In combination
with other elements it is widely distributed on the earth, where the most
important and abundant compound of hydrogen is water, H2O. It is a
component of all the constituents of living matter as well as of many
minerals. It forms an essential part of all hydrocarbons and a vast variety of
other organic substances. Most acids contain hydrogen; the distinguishing
characteristic of an acid is its dissociation, upon going into solution, to yield
hydrogen ions.
Preparation Methods
Hydrogen is prepared in the laboratory:
1. by the action of a dilute strong acid on metals, such as zinc:
Zn + H2SO4 → ZnSO4 + H2↑;
2. by reaction amphoteryc metals with a strong base, such as sodium
hydroxide:
2Al + 6NaOH + 6H2O → 2Na3[Al(OH)6]3 + 3H2↑,
Zn + 2NaOH + 2H2O → Na2[Zn(OH)4] + H2↑;
3. by electrolysis of water:
is
2H2O electroles


→ 2H2 + O2 .
Industrially, hydrogen is prepared from water and hydrocarbons. Until
recently the water-gas reaction was an important way of hydrogen
preparing. The water-gas reaction is an industrial process in which steam
5
is passed over red-hot coke giving a gaseous mixture of carbon monoxide
and hydrogen:
C + H2O(g) → CO + H2.
Now two general processes are used, both starting with natural gas or
petroleum hydrocarbons. The steam-reforming process is an industrial
preparation of hydrogen and carbon monoxide mixtures by the reaction of
steam and hydrocarbons at high temperature and pressure over a nickel
catalyst. For example,
CH4 + H2O(g) → CO + 3H2
C2H8 + 3H2O(g) → 3CO + 7H2 .
The second process involves the partial oxidation of hydrocarbons.
Natural gas, for example, is mixed with a limited supply of oxygen and
burned at elevated pressures:
2CH4 + O2 → 2CO + 4H2.
natural gas
As natural gas and petroleum become more expensive, the water-gas
reaction may become widely used again.
The gaseous mixture of carbon monoxide and hydrogen from these
reactions (called synthesis gas) is used in the preparation of methanol.
Hydrogen is obtained from this mixture free of carbon monoxide by means
of the water-gas shift reaction, an industrial process in which carbon
monoxide reacts with steam in the presence of a catalyst producing carbon
dioxide and hydrogen:
CO + H2O → CO2 + H2.
Carbon dioxide is removed by dissolving the gas in a basic solution to give
carbonate ion.
Reactions of Hydrogen
1. Hydrogen reacts with many nonmetals. In these reactions it derives
+
hydrogen-cation, H . Hydrogen combines with nitrogen in the presence
of a catalyst forming ammonia:
3H2 + N2 → 2NH3;
with sulfur forming hydrogen sulfide:
H2 + S → H2S;
with chlorine forming hydrogen chloride:
H2 + Cl2 → 2HCl;
and with oxygen forming water:
2H2 + O2 → 2H2O.
The reaction of oxygen and hydrogen takes place at room temperature
in the presence of a catalyst such as finely divided platinum only.
When hydrogen is mixed with the air or oxygen and ignited, the
mixture explodes.
6
2. Hydrogen also combines with some metals. Because of the small
electron affinity of the hydrogen atom (73 kJ/rnol, compared with 349
–
kJ/mol for Cl), ionic compounds containing the hydride ion, H , are
formed with metals of the lowest ionization energy only: the alkali
metals and the alkaline earth metals. Thus, sodium and calcium react
with hydrogen gas at moderate temperatures giving the hydrides:
2Na + H2 → 2NaH;
Ca + H2 → CaH2.
These ionic hydrides conduct electricity when molten, indicating the
presence of ions. Hydrogen liberates at the electrode connected to the
positive terminal of the battery, according to the electrode reaction:
–
–
2H → H2 + 2e .
Ionic hydrides react with water, giving hydrogen:
NaH + H2O → NaOH + H2↑;
CaH2 + 2H2O → Ca(OH)2 + 2H2↑.
According to the last reaction, calcium hydride is a convenient,
portable source of small quantities of hydrogen.
3. It acts as a reducing agent on metallic oxides, such as copper oxide,
removing the oxygen and leaving the metal in a free state:
CuO + H2 → Cu + H2O.
4. Hydrogen reacts with unsaturated organic compounds forming
corresponding saturated compounds:
CH2=CH2 + H2 → CH3–CH3 .
Uses
Enormous quantities of hydrogen are used in the manufacture of
ammonia and in the synthesis of methyl alcohol. It is an important
combustible constituent of fuel. Large amounts of hydrogen are required for
the hydrogenation of oils to produce edible fats, of coal to form synthetic
petroleum, and of petroleum oils to enrich the gasoline fraction.
Being the lightest in weight of all gases, hydrogen has been used for the
inflation of balloons and dirigibles. It ignites very easily, however, a small
spark causing it to burn, and several dirigibles, including the Hindenburg,
have been destroyed by hydrogen fires. Helium, which has 92 percent of
the lifting power of hydrogen and is not inflammable, is used whenever
possible. Hydrogen is usually stored in steel cylinders at pressures of 120
to 150 atmospheres. Hydrogen is used also in high-temperature torches for
cutting, melting, and welding metals.
Water
Water is a common name applied to the liquid state of the hydrogenoxygen compound H2O. The ancient philosophers regarded water as a
basic element typifying all liquid substances. Scientists did not discard that
7
view until the latter half of the 18th century. In 1781 the British chemist
Henry Cavendish synthesized water by detonating a mixture of hydrogen
and the air. However, the results of his experiments were not clearly
interpreted until two years later, when the French chemist Antoine Laurent
Lavoisier proved that water was not an element but a compound of oxygen
and hydrogen. In a scientific paper presented in 1804, the French chemist
Joseph Louis Gay-Lussac and the German naturalist Alexander von
Humboldt demonstrated together that water consisted of two volumes of
hydrogen to one of oxygen, as expressed by the present-day formula H2O.
Almost all the hydrogen in water has an atomic weight of 1. In 1932 the
American chemist Harold Clayton Urey discovered the presence in water of
a small amount (1 part in 6000) of so-called heavy water, or deuterium
oxide (D2O). In 1951 the American chemist Aristid Grosse discovered that
naturally occurring water contains also minute traces of tritium oxide (T2O).
Physical Properties and Occurrence of Water
Pure water is an odorless, tasteless liquid. It has a bluish tint, which may
be detected, however, in layers of considerable depth only. At standard
atmospheric pressure (760 mm of mercury) the freezing point of water is
0 °C and its boiling point is 100° C. Water attains its maximum density at a
temperature of 4 °C and expands upon freezing. Its physical properties are
used as standards to define the calorie and specific and latent heat and in
the metric system for the original definition of the unit of mass, the gram.
Water is the only substance that occurs at ordinary temperatures in all
three states of matter, that is, as a solid, a liquid, and a gas. As a solid, or
ice, it is found as glaciers and ice caps, on water surfaces in winter, as
snow, hail, and frost, and as clouds formed of ice crystals. It occurs in the
liquid state as rain clouds formed of water droplets, and on vegetation as
dew; in addition, it covers three-quarters of the surface of the earth in the
form of swamps, lakes, rivers, and oceans. As gas, or water vapor, it
occurs as fog, steam, and clouds. Atmospheric vapor is measured in terms
of relative humidity, which is the ratio of the vapor quantity actually
presents to the greatest amount possible at a given temperature.
Water occurs as moisture in the upper layers of the soil profile, in which
it is held to the particles of soil by capillary action. In this state, it is called
bound water and has different characteristics from free water. Under the
influence of gravity, water accumulates in rock interstices beneath the
surface of the earth as a vast groundwater reservoir supplying wells and
springs and sustaining the flow of some streams during periods of drought.
Chemical Properties of Water
1. Reactions of nonmetal oxides with water.
Water reacts with non-metal oxides giving oxoacids. For example,
8
SO3 + H2O → H2SO4;
CO2 + H2O → H2CO3;
P2O5 + 3H2O → 2H3PO4.
2. Reactions of nonmetal halides with water.
rd
Most of the 3 period and beyond nonmetals halides react with water
giving the corresponding oxoacid and hydrogen halides. For example,
PCl5 + 3H2O → H3PO4 + 5HCI;
SiCI4 + 3H2O → H2SiO2 + 4HCI.
Other halides that react in this way are PF3, PF5, PCl3, and SF4. Sulfur
hexafluoride, SF6, does not react with water and is an exception.
3. Reactions of metal oxides with water.
Water reacts with some metal oxides forming metal hydroxides. For
example,
Na2O + H2O → 2NaOH.
This is a acid-base reaction between the oxide ion, which is a strong
base, and water, which acts as an acid:
2–
–
O + H2O → 2OH .
However many metals oxides such as MgO and CuO are insoluble
and therefore they do not react with water.
4. Oxidation-reduction reactions.
Water is an oxidizing agent and a reducing agent also, although its
oxidizing, and particularly its reducing, properties are rather weak.
When it behaves as an oxidizing agent, that is, when it gains
electrons, it reduces to hydrogen:
–
–
2H2O + 2e → 2OH + H2.
We are familiar with its reaction with sodium and other reactive metals,
which are strong reducing agents:
2Na + 2H2O → 2NaOH + H2.
–
Similarly, water reacts with the hydride ion, H , which oxidizes to
elemental hydrogen:
–
–
H + H2O → H2 + OH
This reaction is an example of an acid-base reaction also, since the
hydride ion is a strong base that adds a proton giving hydrogen. Water
is a very weak reducing agent:
+
–
2H2O → 4H + O2 + 4e .
It reduces very strong oxidizing agents only such as fluorine:
2F2 + 2H2O → 4HF + O2.
Chlorine reacts in the same way with water, but the reaction is very
slow, although it is speeded up somewhat by strong light:
2Cl2 + 2H2O → 4HCl + O2.
The main reaction between chlorine and water is the formation of a
small equilibrium amount of hypochlorous acid, HOCl:
Cl2 + H2O → HOCl + HCl.
9
Water Purification
Suspended and dissolved impurities present in naturally occurring water
make it unsuitable for many purposes. Objectionable organic and inorganic
materials are removed by such methods as screening and sedimentation to
eliminate suspended materials; treatment with such compounds as
activated carbon to remove tastes and odors; filtration; and chlorination or
irradiation to kill infective microorganisms.
During the aeration, or the saturation of water with air, water is brought
into contact with air in such a manner as to produce maximum diffusion,
usually by spraying water into the air in fountains. Aeration removes odors
and taste caused by decomposing of organic matter, and also industrial
wastes such as phenols and volatile gases such as chlorine. It also
converts dissolved iron and manganese compounds into insoluble hydrated
oxides of the metals which then can be readily settled out.
Hardness of natural waters is caused largely by calcium and
magnesium salts and to a small extent by iron, aluminum, and other
metals. Hardness resulting from the bicarbonates of calcium and
magnesium is called temporary hardness and can be removed by boiling,
which also sterilizes the water
∆ CaCO ↓ + CO + H O.
Ca(HCO ) →
3 2
3
2
2
The residual hardness is known as noncarbonate, or permanent,
hardness. The methods of softening noncarbonate hardness include the
addition of sodium carbonate and lime (Ca(OH)2):
Ca(HCO3)2 + Na2CO3 → CaCO3↓ + 2NaHCO3;
Ca(HCO3)2 + Ca(OH)2 → 2CaCO3↓ + 2H2O;
and filtration through natural or artificial zeolites which absorb the
hardness producing metallic ions and release sodium ions to the water.
Sequestering agents in detergents serve to inactivation of the substances
that make water hard.
Iron, which causes an unpleasant taste of drinking water, can be
removed by aeration and sedimentation or by passing water through ironremoving zeolite filters, or the iron may be stabilized by addition of such
salts as polyphosphates. For use in laboratory applications, water is either
distilled or demineralized by passing it through ion-absorbing compounds.
Hydrogen Peroxide
Hydrogen Peroxide is a chemical compound of hydrogen and oxygen
with the formula H2O2. Pure, anhydrous hydrogen peroxide is a colorless,
syrupy liquid with a specific gravity of 1.44. It blisters the skin and has a
metallic taste. The liquid solidifies at –0.41 °C. Concentrated solutions are
unstable, and the pure liquid can explode violently if heated to a
temperature above 100 °C. It is soluble in water in all proportions, and the
10
usual commercial forms are a 3 % and a 30 % aqueous solutions. To retard
the decomposition of the peroxide into water and oxygen, organic
substances, such as acetanilide, are added to the solutions, and they are
kept in dark bottles at low temperature.
Hydrogen peroxide is manufactured in large amounts by the electrolysis
of aqueous solutions of sulfuric acid (or of potassium bisulfate or
ammonium bisulfate):
+
–
H2SO4 ⇄ H + HSO4
On cathode
On anode
+
–
–
–
2H + 2e → H2
2HSO4 – 2e → H2S2O8
+
–
2H2SO4 → H2S2O8 + 2H + 2e ;
H2S2O8 + 2H2O → 2H2SO4 + H2O2.
It is prepared also by the reaction of acid with other peroxides, such as
those of sodium and barium:
Na2O2 + H2SO4 → Na2SO4 + H2O2;
BaO2 + H2SO4 → BaSO4↓ + H2O2 .
Chemical Properties of Hydrogen Peroxide
1. Acid-base properties
Hydrogen peroxide is a week acid. In an aqueous solutions it
ionizes forming hydronium-ion and peroxid-ion:
+
2–
H2O2 + H2O ⇄ H3O + O2 .
2. Oxidation-reduction properties
Hydrogen peroxide acts as both an oxidizing and a reducing agent.
In acidic solution it is an oxidizing agent:
2KI + H2O2 + H2SO4→ I2 + K2SO4 + 2H2O.
However both in basic and in a neutral solutions hydrogen peroxide
can be an oxidizing agent:
Cr2(SO4)3 + 3H2O2 + 10NaOH → 2Na2CrO4 + 3Na2SO4 + 8H2SO4;
PbS + 4H2O2 → PbSO4 + 4H2O.
At the presence of oxidizing agent it exhibits reduction properties in
acidic, basic and neutral solutions:
Cl2 + H2O2 → 2HCl + O2↑;
Ag2O + H2O2 → H2O + O2↑ + 2Ag;
2KMnO4 + 5H2O2 + 3H2SO4 → 2MnSO4 + 5O2 + K2SO4 + 8H2O.
3. Decomposition of hydrogen peroxide
Light, heating and the heavy metals hardly accelerate the process of
hydrogen peroxide decomposition:
MnO
2 → 2H O + O ↑ .
2H2O2  

2
2
11
Uses
Its oxidizing properties are used in the bleaching of substances, such as
hair, ivory, feathers, and delicate fabrics, which would be destroyed by
other agents. It is used also medicinally, in the form of a 3 % aqueous
solution, as an antiseptic and throat wash. Hydrogen peroxide is used in
restoring the original colors on paintings that have darkened through the
conversion of the white lead used in the paintings to lead sulfide. The
hydrogen peroxide oxidizes the black lead sulfide to white lead sulfate. It is
used also as a source of oxygen in the fuel mixture for many rockets and
torpedoes.
Metals
Metals form a group of chemical elements that exhibit all or most of the
following physical qualities: they are solid at ordinary temperatures;
opaque, except in extremely thin films; good electrical and thermal
conductors; lustrous when polished; and have a crystalline structure in the
solid state. Metals and nonmetals are separated in the periodic table by a
diagonal line of elements. Elements to the left of this diagonal are metals,
and elements to the right are nonmetals. Elements that make up this
diagonal—boron, silicon, germanium, arsenic, antimony, tellurium,
polonium, and astatine—have both metallic and nonmetallic properties. The
common metallic elements include the following: aluminum, barium,
beryllium, bismuth, cadmium, calcium, cerium, chromium, cobalt, copper,
gold, iridium, iron, lead, lithium, magnesium, manganese, mercury,
molybdenum, nickel, osmium, palladium, platinum, potassium, radium,
rhodium, silver, sodium, tantalum, thallium, thorium, tin, titanium, tungsten,
uranium, vanadium, and zinc. Metallic elements can combine with one
another and with certain other elements, either as compounds, as
solutions, or as intimate mixtures. A substance composed of two or more
metals, or a substance composed of a metal and certain nonmetals such
as carbon are called alloys. Alloys of mercury with other metallic elements
are known as amalgams.
Within the general limits of the definition of a metal, the properties of
metals vary widely. Most metals are grayish in color, but bismuth is pinkish,
copper is red, and gold is yellow. Some metals display more than one
color, this phenomenon is called pleochroism. The melting points of metals
range from about -39° C for mercury to 3410° C for tungsten. Osmium and
3
iridium (specific gravity 22.6 g/cm ) are the most dense metals, and lithium
3
(specific gravity 0.53 g/cm ) has the least dense. The majority of metals
crystallize in the cubic system, but some crystallize in the hexagonal and
tetragonal systems. Bismuth has the lowest electrical conductivity of the
metallic elements, and silver has the highest at ordinary temperatures. The
conductivity of most metals can be lowered by alloying. All metals expand
12
when heated and contract when cooled, but certain alloys, such as
platinum and iridium alloys, have extremely low coefficients of expansion.
Physical Properties of Metals
Metals are generally very strong and resistant to different types of
stresses. Though there is considerable variation from one metal to the
another, in general metals are marked by such properties as hardness, the
resistance to surface deformation or abrasion; tensile strength, the
resistance to breakage; elasticity, the ability to return to the original shape
after deformation; malleability, the ability to be shaped by hammering;
fatigue resistance, the ability to resist repeated stresses; and ductility, the
ability to undergo deformation without breaking.
Chemical Properties of Metals
Typically metals have positive oxidation states in most of their
compounds, which means they tend to donate electrons to the atoms to
which they bond. Also, metals tend to form basic oxides. Typical
nonmetallic elements, such as nitrogen, sulfur, and chlorine, have negative
oxidation states in most of their compounds—meaning they tend to accept
electrons - and form acidic oxides.
Metals typically have low ionization potentials. This means that metals
react easily by losing electrons to form positive ions, or cations. Thus,
metals can form salts (for example, chlorides, sulfides, and carbonates) by
serving as reducing agents (electron donors).
Electron Structure of Metals
In early attempts to explain the electronic configurations of the metals,
scientists cited the characteristics of high thermal and electrical conductivity
to support a theory that metals consist of ionized atoms in which the free
electrons form a homogeneous sea of negative charge. The electrostatic
attraction between the positive metal ions and the free-moving and
homogeneous sea of electrons was thought to be responsible for the bonds
between the metal atoms. Free movement of the electrons was then
considered to be responsible for the high thermal and electrical
conductivities. The principal objection to this theory was that the metals
should have then higher specific heats than they do.
In 1928 the German physicist Arnold Sommerfeld proposed that the
electrons in metals exist in a quantized arrangement in which low energy
levels available to the electrons are almost fully occupied. In the same year
the Swiss-American physicist Felix Bloch and later the French physicist
Louis Brillouin used this idea of quantization in the currently accepted
“band” theory of bonding in metallic solids.
13
According to the band theory, any given metal atom has only a limited
number of valence electrons with it can use to bond to all of its nearest
neighbors. Extensive sharing of electrons among individual atoms is
therefore required. This sharing of electrons is accomplished through
overlap of equivalent-energy atomic orbitals on the metal atoms that are
immediately adjacent to one another. This overlap is delocalized
throughout the entire metal sample forming extensive orbitals that span the
entire solid rather than being a part of individual atoms. Each of these
orbitals lies at different energies because the atomic orbitals from which
they were constructed were at different energies from the very beginning.
The orbitals are equal in number to the number of individual atomic orbitals
that have been combined, each holds two electrons, and are filled in order
from the lowest to the highest energy until the number of available
electrons has been used up. Groups of electrons are then said to reside in
bands, which are collections of orbitals. Each band has a range of energy
values that the electrons must possess to be a part of that band; in some
metals, there are energy gaps between bands, meaning that there are
certain energies that the electrons cannot possess. The highest energy
band in a metal is not filled with electrons because metals characteristically
possess too few electrons to fill it. The high thermal electrical conductivities
of metals are explained then by the notion that electrons can be transferred
by absorption of thermal energy into these unfilled energy levels of the
band.
2. Group IA. Alkali Metals
Alkali Metals, series of six chemical elements in group IA of the periodic
table in order of their atomic numbers increasing, are lithium Li, sodium Na,
potassium K, rubidium Rb, cesium Cs, and francium Fr.
The atoms of these elements have the following configurations:
2
1
1
– 1s 2s
[He]2s
3Li
2
2
6
1
1
[Ne]3s
11Na – 1s 2s 2p 3s
2
2
6
2
6
1
1
– 1s 2s 2p 3s 3p 4s
[Ar]4s
19K
2
2
6
2
6
10
2
6
1
1
[Kr]5s
37Rb – 1s 2s 2p 3s 3p 3d 4s 4p 5s
2
2
6
2
6
10
2
6
10
2
6
1
1
[Xe]6s
55Cs – 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p 6s
2
2
6
2
6
10
2
6
10 14
2
6
10
2
6
1
1
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s 6p 7s [Rn]7s
87Fr
Table 2.1 lists some physical properties of the alkali metals.
Physical Properties and Occurrence of the Alkali Metals
Lithium, symbol Li, is a silvery white metallic element. The atomic
number of lithium is 3. Discovery of the element is generally credited to
Johann A. Arfvedson in 1817. Lithium melts at about 181 °C, boils at about
14
1342 °C, and has a specific gravity of 0.53. The atomic weight of lithium is
6.941.
th
Lithium ranks 35 in order of abundance of the elements in the crust of
the earth. It does not occur in nature in the free uncombined state but in the
form of compounds only, which are widely distributed.
Table 2.1. Properties of Group IA Elements
Property
Electron configuration
Melting point, °C
Boiling point, °C
3
Density, g/cm
Ionization energy,
kJ/mol
Electronegativity
(Pauling scale)
Standard potential
(volts),
+
–
0
M +e →M
Covalent radius, Å
Ionic radius, Å
Lithium
1
[He]2s
181
1347
0.53
Sodium
1
[Ne]3s
97.8
883
0.97
Potassium Rubidium
1
1
[Ar]4s
[Kr]5s
63.6
38.9
774
688
0.86
1.53
Cesium
1
[Xe]6s
28.4
678
1.88
520
496
419
403
376
1.0
0.9
0.8
0.8
0.7
–3.04
–2.71
–2.92
–2.92
–2.92
1.23
0.90
1.57
1.16
2.02
1.52
2.16
1.66
2.35
1.81
Sodium, symbol Na, is a silvery-white, extremely soft metallic element.
The atomic number of sodium is 11. It was discovered in 1807 by the
British chemist Sir Humphry Davy.
Elemental sodium is a metal which is soft enough to be cut with a knife.
It has a hardness of 0.4. It melts at about 98 °C, boils at about 883 °C, and
has a specific gravity of 0.97. The atomic mass of sodium is 22.99.
Sodium is found in nature in the combined state only. It occurs in the
ocean and in salt lakes as sodium chloride, NaCl, and less often as sodium
carbonate, Na2CO3, and sodium sulfate, Na2SO4.
Potassium, symbol K (from Latin kalium, “alkali”). The atomic number of
potassium is 19. Potassium was discovered and named in 1807 by the
British chemist Sir Humphry Davy. The metal is silvery white and can be
cut with a knife. It has a hardness of 0.5. Potassium exists in three natural
isotopic forms, with mass numbers of 39, 40, and 41. Potassium-40 is
radioactive and has a half-life of 1.28 milliard years. The most abundant
isotope is potassium-39. Several radioactive isotopes have been artificially
prepared. Potassium melts at about 63 °C, boils at about 760 °C, and has a
specific gravity of 0.86; the atomic mass of potassium is 39.098.
Potassium is found in nature in large quantities, ranking eighth in order
of abundance of the elements in the crust of the earth, in various minerals
such as sylvite (KCl), carnallite (KMgCl3·6H3O), feldspar (KalSi3O8) and
15
saltpeter (KNO3). Potassium is a constituent of all plant and animal tissue
as well as a vital constituent of fertile soil.
Rubidium (Latin rubidus, “red”), symbol Rb, has an atomic number of
37. Rubidium was discovered spectroscopically in 1860 by the German
chemist Robert Wilhelm Bunsen and the German physicist Gustav Robert
Kirchhoff, who named the element after the red lines prominent in its
spectrum. Metallic rubidium is silvery white and very soft. Rubidium melts
at about 39 °C, boils at about 686 °C, and has a specific gravity of 1.53; the
atomic weight of rubidium is 85.468.
th
It is a widely distributed element, ranking 16 in order of abundance of
the elements in the crust of the earth. It is not found in large deposits but
occurs in small amounts in certain mineral waters and in many minerals
usually associated with other alkali metals.
Cesium, symbol Cs, is a white, soft element. The atomic number of
cesium is 55. Cesium melts at about 28° C, boils at about 669° C, and has
a specific gravity of 1.88; its atomic weight is 132.905. Cesium was
discovered in 1860 by the German chemist Robert Wilhelm Bunsen and the
German physicist Gustav Robert Kirchhoff by means of spectroscopy.
th
Cesium ranks about 46 in natural abundance among the elements in
crustal rocks.
Francium exists in a radioactive form only.
Preparation of the Elements
Sodium, potassium and lithium metals are the most easily prepared by
electrolysis of their fused chlorides. For example,
electrolysis
2NaCl (l)   → 2Na + Cl2↑.
The other alkali metals, rubidium and cesium, are prepared by chemical
reduction of their salts also. For example, when molten cesium chloride is
heated at 700 °C to 800 °C with calcium metal at low pressure, cesium
vapor distills over:
2CsCl + Ca → CaCl2 + 2Cs(g).
Chemical Properties of the Alkali Metals
The alkali metals are the most reactive metallic elements, readily losing
1
the ns valence electron to form compounds in the +1 oxidation state.
These elements are very reactive. The metals react directly with many
other elements and compounds. For example, with:
Oxygen:
Sulfur:
4Li + O2 → 2Li2O (lithium oxide)
Na + S → Na2S (sodium sulfide)
2Na + O2 → Na2O2 (sodium peroxide) Halogens:
K + O2 → KO2 (potassium superoxide) 2Na + Cl2 → 2NaCl
Water:
Rb + O2 → RbO2
2K + 2H2O → 2KOH + H2↑
16
Cs + O2 → CsO2
Nitrogen:
6Li + N2 → 2Li3N (lithium nitride)
Hydrogen:
2Na + H2 → 2NaH (sodium hydride)
Oxides and Hydroxides of the Alkali Metals
Oxides of the alkali metals are a basic oxides and has all properties of
basic oxides:
1. Reactions with water are:
Na2O + H2O → 2NaOH;
Li2O + H2O → 2LiOH.
2. Reactions with acids are:
Na2O + 2HCl → 2NaCl + H2O;
K2O + H2SO4 → K2SO4 + H2O.
3. Reactions with acidic oxides are:
Li2O + CO2 → Li2CO3;
K2O + SO3 → K2SO4.
Hydroxides of the alkali metals are strong bases. They react with:
─ acids
NaOH + HCl → NaCl + H2O;
LiOH + H2SO4 → Li2SO4 + H2O.
─ acidic oxides
2LiOH + CO2 → Li2CO3;
2KOH + SO3 → K2SO4.
─ salts
2NaOH + CuCl2 → 2NaCl + Cu(OH)2↓;
2KOH + FeSO4 → K2SO4 + Fe(OH)2 ↓.
Uses
Sodium is used in the manufacture of tetraethyl lead and as a cooling
agent in nuclear reactors. The most important compound of sodium is
sodium chloride, known commonly as salt. Other important compounds are
sodium carbonate, known as washing soda, and sodium bicarbonate,
known as baking soda. Sodium tetraborate is known commonly as borax.
Sodium fluoride, NaF, is used as an antiseptic. Sodium peroxide, Na2O2, is
an important bleaching and oxidizing agent.
Sodium is a vital element and the human diet must contain a sensible
amount of sodium. The sodium cation is the main extracellular cation in
animals and is important for nerve function in animals. The importance of
sodium as salt in the diet was recognized well before sodium was
understood to be an element. This recognition formed the basis of biblical
time trading by the Romans in salt deposits lining the Dead Sea. Prolonged
17
sweating results in sodium ion loss in sweat and it is most important that
the sodium ion is replaced.
Potassium bromide (KBr) is used in photography, engraving, and
lithography, and in medicine as a sedative. Potassium iodide (KI) is used in
medicine for the treatment of rheumatism and overactivity of the thyroid
gland. Potassium nitrate (KNO3) is used in matches, explosives, and
fireworks, and in pickling of meat. It occurs naturally as saltpeter.
Potassium permanganate (KMnO4) is used as a disinfectant and germicide
and as an oxidizing agent in many important chemical reactions. Potassium
hydrogen tartrate (KHC4H4O6), commonly known as cream of tartar, is a
white solid used in baking powder and in medicine.
Potassium salts are essential for life. The potassium cation is the major
cation in intracellular fluids (sodium is the main extracellular cation). It is
essential for nerve and heart function. A normal diet containing reasonable
amounts of vegetables contains all the potassium required. Potassium is
important for plants. Potassium salts are extremely toxic when injected.
Vomiting (emesis) helps prevents toxic effects from ingestion of excess
amounts of potassium.
Lithium compounds are regarded as slightly toxic, and perhaps more
than the other Group IA elements. Lithium appears not to have a biological
role, which does not meant that lithium compounds do not have an affect.
Sometimes, lithium-based drugs such as lithium carbonate (Li2CO3) are
used to treat manic-depressive disorders in doses of around 0.5 g - 2 g
daily. Some side effects are known. Ingestion of large amounts of lithium
results in drowsiness, slurred speech, vomiting, and other symptoms.
Excess of lithium poisons the central nervous system.
Biological function. Sodium, potassium, and chlorine occur almost
entirely in the fluids and soft tissues of the body, sodium and chlorine being
found mainly in the body fluids, and potassium occuring mainly in the cells.
They serve a vital function in controlling osmotic pressures and acid-base
equilibrium. They also play important roles in water metabolism. Sodium is
the main monovalent ion of extracellular fluids; sodium ions constituting
93% of the ions (bases) found in the blood stream. Although the principal
role of sodium in the animal is connected with the regulation of osmotic
pressure and the maintenance of acid-base balance, sodium also has an
effect on muscle irritability, and plays a specific role in the absorption of
carbohydrate. Potassium is the major cation of intracellular fluid, and
regulates intracellular osmotic pressure and acid-base balance. Like
sodium, potassium has a stimulating effect on muscle irritability. Potassium
is also required for glycogen and protein synthesis, and the metabolic
breakdown of glucose.
18
3. Group IIA. Beryllium, Magnesium and Alkaline Earth
Metals
There are six chemical elements in group IIA of the periodic table:
beryllium Be, magnesium Mg, calcium Ca, strontium Sr, barium Ba and
radium Ra. Calcium, strontium and barium name alkaline earth metal.
The atoms of these elements have the following configurations:
2
2
2
– 1s 2s
[He]2s
4Be
2
2
6
2
2
[Ne]3s
12Mg – 1s 2s 2p 3s
2
2
6
2
6
2
2
[Ar]4s
20Ca – 1s 2s 2p 3s 3p 4s
2
2
6
2
6
10
2
6
2
2
Sr
–
1s
2s
2p
3s
3p
3d
4s
4p
5s
[Kr]5s
38
2
2
6
2
6
10
2
6
10
2
6
2
2
[Xe]6s
56Ba – 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p 6s
2
2
6
2
6
10
2
6
10 14
2
6
10
2
6
2
2
88Ra – 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s 6p 7s [Rn]7s
Table 3.1 lists some physical properties of IIA group metals.
Table 3.1. Properties of Group IIA Elements
Property
Beryllium
Magnesiu
m
2
[Ne]3s
649
1090
1.74
Calcium
Electron configuration
Melting point, °C
Boiling point, °C
Density, g/cm3
Ionization energy,
kJ/mol
Electronegativity
(Pauling scale)
Standard
potential
2+
–
(volts), M + 2e → M
Covalent radius, Å
Ionic radius, Å
[He]2s
1278
2970
1.85
899
738
590
549
503
1.5
1.2
1.0
1.0
0.9
–1.70
–2.38
–2.76
–2.89
–2.90
0.89
0.59
1.36
0.86
1.74
1.14
1.92
1.32
1.98
1.49
2
[Ar]4s
839
1484
1.54
2
Strontium
[Kr]5s
769
1384
2.6
2
Barium
[Xe]6s
725
1640
3.51
2
Physical Properties and Occurrence
Beryllium, symbol Be, is a gray, brittle metallic element, with an atomic
number of 4. It was discovered as an oxide, known as beryllia now, in 1797
by French chemist Louis Nicolas Vauquelin. The free element was firstly
isolated in 1828 independently by Friedrick Wuhler and Antonine Alexandre
Brutus Bussy. Its atomic weight is 9.012. Beryllium melts at about 1287 °C,
boils at about 3000 °C, and has a specific gravity of 1.85.
Magnesium, symbol Mg, is a silvery white metallic element. The atomic
number of magnesium is 12. The metal was isolated firstly by the British
chemist Sir Humphry Davy in 1808. Magnesium melts at about 649 °C,
19
boils at about 1107 °C, and has a specific gravity of 1.74; the atomic mass
of magnesium is 24.305.
Magnesium ranks sixth in natural abundance among elements in crustal
rocks. It occurs in nature in chemical combination with other elements only,
particularly as the minerals carnallite (KMgCl3 · 6H2O), dolomite (CaCO3 ·
MgCO3), and magnesite (MgCO3); in many rock-forming silicates; and as
salts, such as magnesium chloride, in ocean and saline-lake waters. It is an
essential constituent of animal and plant tissue.
Calcium, symbol Ca, is a silvery-white metallic element. The atomic
number of calcium is 20. The British chemist Sir Humphry Davy isolated
calcium in 1808 by means of electrolysis. Calcium melts at about 839 °C,
boils at about 1484 °C, and has a specific gravity of 1.54; its atomic mass is
40.08.
Calcium is fifth in abundance among the elements in the earth's crust,
but it is not found uncombined in nature. It occurs in many highly useful
compounds, such as calcium carbonate (CaCO3), of which calcite, marble,
limestone, and chalk are composed; calcium sulfate (CaSO4) in alabaster
or gypsum; calcium fluoride (CaF2) in fluorite; calcium phosphate
(Ca3(PO4)2) in rock phosphate; and in many silicates.
Strontium, symbol Sr, is a malleable, ductile metallic element. The
atomic number of strontium is 38. Metallic strontium was isolated firstly by
the British chemist Sir Humphry Davy in 1808. Strontium melts at about
769 °C boils at about 1384 °C, and has a specific gravity of 2.6. The atomic
weight of strontium is 87.62.
Strontium is never found in the elemental state, occurring mainly as
th
strontianite (SrCO3), and celestite (SrSO4). Strontium ranks about 15
among the elements in natural abundance in the earth's crust and is widely
distributed in small quantities.
Barium, symbol Ba, is a soft, silvery, metallic element. The atomic
number of barium is 56. Barium was recognized as an element firstly in
th
1808 by the English scientist Sir Humphry Davy. Barium is the 14 most
th
common element, making up 1/2000 of the crust of the earth. The atomic
weight of barium is 137.33. The element melts at about 725 °C, boils at
about 1640 °C, and has a specific gravity of 3.5.
Barium occurs in nature in the form of its compounds only. Its most
important compounds are the minerals barium sulfate and barium
carbonate (witherite, BaCO3).
Preparation of Elements
Metals of Group IIA, like the alkali metals, are prepared by electrolysis
of the molten halides (usually the chlorides) or by chemical reduction of
either the halides or the oxides.
Beryllium extraction from ores is complex. The mineral beryl,
[Be3Al2(SiO3)6] is the most important source of beryllium. It is roasted with
20
sodimu hexafluorosilicate, Na2SiF6, at 700 °C to form beryllium fluoride.
This is water soluble and the beryllium can be precipitated as the hydroxide
Be(OH)2 by adjustment of the pH to 12.
Pure beryllium can be obtained by electrolysis of beryllium chloride,
BeCl2, to which sodium chloride is added to increase the conductivity of the
molten salt:
electrolysis
BeCl2(l)   → Be + Cl2↑.
Another method involves the reduction of beryllium fluoride with
magnesium at 1300°C:
∆ MgF + Be.
BeF + Mg →
2
2
Magnesium can be prepared commercially by several processes and
normally is not made in the laboratory because of its ready availability.
There are massive amounts of magnesium in seawater. It can be
recovered as magnesium chloride, MgCl2 through reaction with calcium
oxide, CaO:
CaO + H2O → Ca(OH)2;
2+
2+
Mg + Ca(OH)2 → Mg(OH)2↓ + Ca ;
Mg(OH)2 + 2HCl → MgCl2 + 2H2O.
Electrolysis of hot molten MgCl2 affords magnesium as a liquid which is
poured off, and chlorine gas:
electrolysis
MgCl2(l)   → Mg + Cl2↑.
The other methods used to produce magnesium are non electrolytic and
involve dolomite, (MgCa(CO3)2), an important magnesium mineral. It is
"calcined" by heating to form calcined dolomite, MgO·CaO, which reacts
with ferrosilicon alloy:
∆ 2Mg + Ca SiO + Fe.
2[MgO·CaO] + FeSi →
2
4
The magnesium can be distilled out from this mixture of products.
Calcium is prepared by electrolysis of molten calcium chloride and by
reduction of calcium oxide by aluminum in a vacuum, where the calcium
produced distills off:
∆ 3Ca + Al O .
3CaO + 2Al →
2
3
Barium is produced by reduction of the oxide by aluminum also, and
although very small amounts of strontium are used commercially, it can be
produced by a similar process.
Chemical Properties of Beryllium and its Compounds
Beryllium reacts directly with many nonmetals. For example, with:
Oxygen:
Sulfur:
2Be + O2 → 2BeO (beryllium oxide) Be + S → BeS (beryllium sulfide)
Nitrogen:
Halogens:
3Be + N2 → Be3N2 (beryllium nitride) Be + Cl2 → BeCl (beryllium chloride)
21
Beryllium is a amphoteric metal and reacts with strong acids and bases:
Be + H2SO4 → BeSO4 + H2↑;
3Be + 8НNO3(diluted) → Be(NO3)2 + 2NO↑ + 4H2O;
Be + 2NaOH + 2H2O → Na2[Be(OH)4] + H2↑;
t
Be + 2NaOH 
→ Na2BeO2 + H2↑.
Concentrated nitric and sulfuric acids do not react with beryllium.
Beryllium oxide is a white, high-melting compound. It is prepared by
decomposition of beryllium hydroxide or beryllium carbonate:
∆ BeO + H O;
Be(OH)2 →
2
∆
BeCO → BeO + CO .
3
2
Beryllium oxide is a amphoteric oxide, which reacts with strong acids and
bases:
BeO + H2SO4 → BeSO4 + H2O;
BeO + 2NaOH + H2O → Na2[Be(OH)4].
At heating beryllium oxide react with acidic and basic oxides:
∆ BeSiO ;
BeO + SiO →
2
3
∆ Na BeO .
BeO + Na2O →
2
2
Beryllium hydroxide is insoluble in water. It is an amphoteric hydroxide which
reacts readily with strong acids and bases:
Be(OH)2 + H2SO4 → BeSO4 + 2H2O;
Be(OH)2 + 2NaOH → Na2[Be(OH)4].
Beryllium sulfide like its oxide has amphoteric properties, at heating it reacts
with acidic and basic sulfides:
∆ BeSiS ;
BeS + SiS →
2
3
∆ Na BeS .
BeS + Na2S →
2
2
2+
Be has a tendency to accept electrons due to its small size. As a
result, bonding in its compounds is highly covalent. For example, the
melting point of beryllium chloride (450 °C) is low compared with that of the
other chlorides of Group IIA elements (MgCl2 melts at 714 °C). Also, the
molten salt has a low electrical conductivity, so electrolysis of BeCl2
requires the addition of a salt such as NaCl. In the solid phase, beryllium
chloride has a polymeric structure (that is, covalent substance with an
infinitely repeating unit), in which each chlorine atom "bridges" two
beryllium atoms:
.
22
Chemical Properties of Magnesium and its Compounds
Magnesium reacts with nonmetals. For example, with:
oxygen
sulfur
2Мg + O2→2MgO (magnesium oxide) Mg + S→MgS (magnesium sulfide)
nitrogen
Halogens
3Mg + N2→Mg3N2 (magnesium nitride) Mg+Cl2→MgCl (magnesium chloride)
Magnesium reacts directly with many acids:
Mg + 2HCl → MgCl2 + H2↑;
Mg + H2SO4 → MgSO4 + H2↑;
4Mg + 10HNO3(dilute) → 4Mg(NO3)2 + NH4NO3 + 3H2O.
Magnesium oxide is a white, high-melting compound. It can be prepared
by decomposition of magnesium carbonate:
∆ MgO + CO .
MgCO →
3
2
Magnesium oxide is a basic oxide which reacts with acids and acidic
oxides:
MgO + H2SO4 → MgSO4 + H2O;
∆ MgCO .
MgO + CO →
2
3
Magnesium hydroxide is insoluble in water, it is a basic hydroxide which
reacts readily with acids and acidic oxides:
Mg(OH)2 + H2SO4 → MgSO4 + 2H2O;
Mg(OH)2 + SO3 → MgSO4 + H2O.
Chemical Properties of the Alkaline Earth Metals
The chemistry of the alkaline earth metals is relatively simple. Since two
2+
electrons can be easily removed from their atoms to form the M ions,
there is only one stable oxidation state +2 respectively. With the exception
of a very few compounds, all compounds of alkaline earth metals are ionic.
As we might expect from their low ionization energies, all these
elements are very reactive. The metals react directly with many other
elements and compounds. They react with the halogens giving ionic
halides:
Ca + Cl2 → CaCl2;
Sr + Cl2 → SrCl2;
Ba + Cl2 → BaCl2.
The alkaline earth metals react with hydrogen at heating giving ionic
hydrides:
Ca + H2 → CaH2;
Sr + H2 → SrH2;
Ba + H2 → BaH2.
Calcium, strontium, and barium react with nitrogen at heating, forming
3–
ionic nitrides containing the N ion:
23
∆ Ca N ;
3Ca + N2 →
3 2
∆
3Sr + N2 → Sr3N2;
∆ Ba N .
3Ba + N →
2
3
2
Since all alkaline earth metals react readily with oxygen, they must be
stored out of contact with the atmosphere. The reactions with oxygen are
more complicated than we might have expected. Calcium and strontium
give the expected oxides, CaO and SrO, in a pure state when they react
with excess of oxygen under ordinary conditions. Barium gives peroxide:
Ba + O2 → BaO2.
Oxides and Hydroxides of the Alkaline Earth Metals
Oxides of the alkaline earth metals are basic oxides and have all
properties of basic oxides.
1. Reactions with water are:
CaO + H2O → Ca(OH)2;
BaO + H2O → Ba(OH)2.
2. Reactions with acids are:
SrO + 2HCl → SrCl2 + H2O;
BaO + 2HNO3 → Ba(NO3)2 + H2O.
3. Reactions with acidic oxides are:
CaO + CO2 → CaCO3;
BaO + SO3 → BaSO4.
Hydroxides of the alkaline earth metals are strong bases. These react
with:
─ acids
Ca(OH)2 + 2HCl → CaCl2 + H2O;
─ acidic oxides
Sr(OH)2 + CO2 → SrCO3↓;
─ salts
Ca(OH)2 + CuCl2 → CaCl2 + Cu(OH)2↓.
Carbonates of the alkaline earth metals are minerals and are the most
important sources of these elements. Unlike the alkali metal carbonates
(other than Li2CO3), they decompose when heated, giving the oxides and
CO2:
∆ CaO + CO ↑ .
CaCO →
3
2
Uses and biological role
Although beryllium products are safe to use and handle, the fumes and
dust released during fabrication are highly toxic. Extreme care must be
24
taken to avoid breathing or ingesting even very small amounts. Specially
designed exhaust hoods are used by persons working with beryllium oxide.
Beryllium and its oxide are being utilized more and more in industry.
Besides its importance in aircraft and X-ray tubes, beryllium is used in
computers, lasers, televisions, oceanographic instruments, and personal
body armor.
Beryllium has no biological role. Beryllium metal dust can cause major
lung damage and beryllium salts are very toxic. Compounds containing
beryllium are very poisonous and should be handled by a professional
under controlled conditions only. One route for beryllium into the biosphere
is by way of industrial smoke.
Magnesium is an important element in both plant and animal life.
Chlorophylls are porphyrins based upon magnesium. Magnesium is
required for the activation of some enzymes. The adult daily requirement of
magnesium is about 0.3 g·day-1. Magnesium is an essential component of
bone, cartilage and the crustacean exoskeleton.
Magnesium is an activator of several key enzyme systems, including
kinases, (i.e. enzymes that catalyse the transfer of the terminal phosphate
of ATP to sugar or other acceptors), mutases (transphosphorylation
reactions), muscle ATPases, and the enzymes cholinesterase, alkaline
phosphatase, enolase, isocitric dehydrogenase, arginase (magnesium is a
component of the arginase molecule), deoxyribonuclease, and
glutaminase. Through its role in enzyme activation, magnesium (like
calcium) stimulates muscle and nerve irritability (contraction), is involved in
the regulation of intracellular acid-base balance, and plays an important
role in carbohydrate, protein and lipid metabolism.
Magnesium hydroxide used in medicine as the laxative “milk of
magnesia,” and in sugar refining; magnesium sulfate (MgSO4·7H2O), well
known as Epsom salt; and magnesium oxide (MgO), called burnt
magnesia, or magnesia are used as a heat-refractory and insulating
material, in cosmetics, as a filler in paper manufacture, and as a mild,
antacid laxative.
Calcium is present in the chemically combined state in lime (calcium
hydroxide), cement and mortar (as calcium hydroxide or a variety of
silicates of calcium), teeth and bones (as a calcium hydroxyphosphate),
and in many body fluids (as complex proteinaceous compounds) essential
to muscle contraction, the transmission of nerve impulses, and the clotting
of blood.Calcium is an essential component of bone, cartilage and the
crustacean exoskeleton. Calcium is essential for the normal clotting of
blood, by stimulating the release of thromboplastin from the blood platelets.
It is an activator for several key enzymes, including pancreatic lipase, acid
phosphatase, cholinesterase, ATPases, and succinic dehydrogenase.
Through its role in enzyme activation, calcium stimulates muscle
contraction (ie. promotes muscle tone and normal heart beat) and regulates
25
the transmission of nerve impulses from one cell to another through its
control over acetylcholine production. Calcium, in conjunction with
phospholipids, plays a key role in the regulation of the permeability of cell
membranes and consequently over the uptake of nutrients by the cell. It is
believed to be essential for the absorption of vitamin B12 from the gastrointestinal tract.
Strontia (strontium oxide), SrO, is used in recovering of sugar from beetsugar molasses. A radioactive isotope of the element, strontium-85, is used
in the detection of bone cancer. Strontium-90 is a dangerous radioactive
isotope found in the fallout that results from the detonation of some nuclear
weapons.
Strontium has no biological role. But chemically, strontium resembles
calcium and the human body does a poor job of distinguishing the two.
Therefore it is absorbed by the body and stored in the skeleton in places
90
where calcium should be. This happens also with radioactive Sr which
was produced by above-ground nuclear explosions in the 1950s.
Regrettably strontium-90 is widely spread in the environment.
Barium metal has few practical applications, although it is used
sometimes in coating electrical conductors in electronic apparatus and in
automobile ignition systems. Barium sulfate (BaSO4) is used as a filler for
rubber products, in paint, and in linoleum. Barium nitrate is used in
fireworks, and barium carbonate in rat poisons. A form of barium sulfate,
which is opaque to X-rays, is used for the X-ray examination of the
gastrointestinal tract.
Barium has no biological role. The British Pharmaceutical Codex from
1907 indicates that barium chloride ("barii chloridum", BaCl2·2H2O] has a
stimulant action on the heart and other muscles. It was said that it "raises
blood pressure by constricting the vessels and tends to empty the
intestines, bladder, and gall bladder". Its poisonous nature was pointed out
also. Barium sulfide (BaS) was used as a depilatory agent (removes hair).
Barium sulfate (BaSO4) is insoluble and used for body imaging (barium
meal).
All dissolvable barium compounds should be regarded as highly toxic
although initial evidence would appear to suggest the danger is limited.
Barium salts can damage the liver. The metal dust presents a fire and
explosion hazard.
4. Group VIB. Chromium, Molybdenum, Tungsten
Group VIB of the periodic table contains three chemical elements –
chromium Cr, molybdenum Mo, and tungsten W.
The atoms of these elements has the following configurations:
2
2
6
2
6
5
1
5
1
– 1s 2s 2p 3s 3p 3d 4s
[Ar]3d 2s
24Cr
26
42Mo
74W
2
2
6
2
6
10
2
6
5
1
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s
2
2
6
2
6
10
2
6
10 14
2
6
4
2
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s
5
1
[Kr]4d 5s
4
2
[Xe]5d 6s
The unexpected electron configuration of chromium (together with
5
1
4
2
molybdenum) 3d 4s , rather than 3d 4s , emphasizes that the 3d and 4s
levels are very close in energy and that alternative electron arrangements
can differ only very slightly in energy. When the 3d and 4s levels are
sufficiently close, the repulsion between the two electrons in the 4s orbital
can cause one of them to move into a 3d orbital.
Table 4.1 lists some physical properties of of group VIB metals.
Table 4.1. Properties of Group VIB Elements
Chromium
Molybdenum
Tungsten
Property
Electron configuration
5
[Ar]3d 2s
1
5
[Kr]4d 5s
1
4
[Xe]5d 6s
Melting point, °C
1857
2610
3390
Boiling point, °C
2672
4640
5680
7.2
10.2
19.5
Ionization energy, kJ/mol
652
685
770
Electronegativity (Pauling
scale)
1.6
1.3
1.4
1.27
1.37
1.4
0.52
0.65
0.65
Density, g/cm
3
o
Covalent radius, A
6+
o
Ionic radius (Me ), A
2
Physical Properties and Occurrence of the Elements
Chromium, symbol Cr, is a gray metallic element that can aquire a high
polish. The atomic number of chromium is 24. The element was discovered
in 1797 by the French chemist Louis Nicolas Vauquelin, who named it
chromium (Greek chroma, “color”) because of the variety of different colors
of its compounds.
st
Chromium is a common element; overall it ranks about 21 in natural
abundance among the elements in crystal rocks. Chromium has an atomic
weight of 51.996; the element melts at about 1857 °C, boils at about
2672 °C, and has a specific gravity of 7.2.
Chromium can replace a part of the aluminium or iron in many minerals,
imparting to them their unique colors. Many precious stones owe their color
to the presence of chromium compounds. Workable ores are rare,
however, chromite (FeCr2O4) being the only important one.
Molybdenum, symbol Mo, is a silvery white, tough, malleable metal. The
atomic number of molybdenum is 42. Molybdenum was discovered in 1778
by the Swedish chemist Carl Wilhelm Scheele. Molybdenum melts at about
27
2610 °C, boils at about 4640 °C, and has a specific gravity, or relative
density, of 10.2.
Molybdenum does not occur in free state in nature, but in the form of its
ores, the most important of which are molybdenite (molybdenum disulfide,
th
MoS2) and wulfenite (lead molybdate, PbMoO4). It ranks 56 in order of
abundance of the elements in the crust of the earth and is an important
trace element in soils, where it contributes to the growth of plants.
Tungsten, symbol W (from the earlier name, wolfram), is a metallic
element that has the highest melting point of any metal. The atomic number
of tungsten is 74. The atomic weight of tungsten is 183.85.
Some credit the Swedish chemist Carl Wilhelm Scheele with the
discovery of tungsten in 1781, while others name the Spanish brothers
Juan Jos and Fausto D'Elhuyar as its discoverers, in 1783. Pure tungsten
is silver-white in color and is ductile; the more easily obtained impure form
is steel-gray and is hard and brittle. Tungsten melts at about 3410 °C, boils
at about 5660 °C, and has a specific gravity of 19.5.
th
Tungsten ranks 57 in abundance among the elements in the earth’s
crust. It is never found free in nature, but occurs in combination with other
metals, notably in the minerals scheelite (calcium tungstanate, CaWO4) and
wolframite (iron and manganese tungstanate, FeMn(WO4)2), which are the
important tungsten ores.
Preparation of the Elements
The most commercially useful source of chromium is the chromite ore,
FeCr2O4. Oxidation of this ore by air in molten alkali gives sodium
chromate, Na2CrO4 in which the chromium is in the +6 oxidation state:
∆ 8Na CrO + 2Fe O + 8CO ↑.
4FeCr O + 7O + 8Na CO →
2
4
2
2
3
2
4
2
3
2
It is converted to the Cr(III) oxide Cr2O3 by extraction with water,
precipitation, and reduction with carbon:
∆ 4CO↑ + 2Na O + Cr O .
2Na CrO + 4C →
2
4
2
2
3
The oxide is further reduced with aluminium or silicon to form chromium
metal:
Cr2O3 + 2Al → Cr + Al2O3;
2Cr2O3 + 3Si → 4Cr + 3SiO2.
Molybdenum and tungsten are most easily prepared by reduction their
oxides with hydrogen:
∆ Mo + 3H O.
МоO + 3H →
3
2
2
Chemical Properties of Chromium
Oxidation states. The principal oxidation states of chromium are +2, +3
and +6. The +6 oxidation state is the maximum possible oxidation state for
chromium, because in this state it uses all six valence electrons for the
28
formation of bonds. Chromium(VI) compounds are generally strong
oxidizing agents and are readily reduced to the +3 state, which is the most
stable oxidation state of chromium. In the +2 state chromium is a reducing
agent and is oxidized readily to the +3 state.
Chromium metal reacts directly with chlorine, sulfur, nitrogen and
oxygen at heating:
∆ CrCl ;
∆ 2CrN;
Cr + Cl →
2Cr + N →
2
3
2
∆ Cr S ;
∆ 2Cr O .
4Cr + 3O2 →
2Cr + 3S →
2 3
2 3
2+
With acids chromium metal reacts releasing hydrogen gas and Cr ion:
Cr + 2HCl → CrCl2 + H2↑;
Cr + H2SO4 → CrSO4 + H2↑.
The +2 oxidation state
Chromium(II) oxide (CrO) is a basic oxides and has all properties of
basic oxides:
CrO + 2HCl → CrCl2 + H2O;
CrO + SO3 → CrSO4.
It reacts with oxygen at heating up:
4CrO + O2 → 2Cr2O3.
Chromium(II) hydroxide (Cr(OH2)) has basic properties only and in
reactions with acids forms salts of chromium(II):
Cr(OH)2 + 2HCl → CrCl2 + 2H2O;
Cr(OH)2 + SO3 → CrSO4 + H2O.
2+
2+
In aqueous solution Cr ion is hydrated, [Cr(H2O)6] , and is blue in
color.
The chromium(ll) ion is easily oxidized by O2 to chromium(lll) ion:
4CrCl2 + O2 + 4HCl → 4CrCl2 + 2H2O;
thus to prepare chromium(ll) salts, we must carry out this reaction in the
absence of air. Chromium(ll) ion is oxidized even by hydrogen ion:
2+
+
3+
2Сr + 2H → 2Cr +H2↑;
but the reaction is slow.
The +3 oxidation state
3+
The Cr ion, which has a violet color in aqueous solution, forms many
salts. These include CrCl3 · 6H2O, Cr2(SO4)3 · 18H2O, and chrome alum,
KCr(SO4)2 ·12H2O, which forms large, violet octahedral crystals. In these
3+
3+
salts Cr ion is hydrated by six water molecules giving the ion [Cr(H2O)6] .
When an aqueous solution of ammonia or sodium hydroxide is added to
a solution of a chromium(lll) salt, a gray-green precipitate of chromium(lll)
hydroxide (Cr(OH)3) is obtained:
CrCl3 + 3NaOH → Cr(OH)3↓ + 3NaCl.
29
If the hydroxide is strongly heated, it dehydrates giving chromium(lll)
oxide:
∆ Cr O + 3H O.
2Cr(OH) →
3
2
3
2
Chromium(lll) oxide can be prepared also conveniently in a rather
spectacular reaction by heating ammonium dichromate:
(NH4)2Cr2O7 → Cr2O3 + N2↑ + 4H2O.
In this reaction the dichromate ion oxidizes the ammonium ion to nitrogen
and reduces itself to Cr2O3.
Chromium(lll) oxide and hydroxide have amphoteric properties, they
react with strong acids and bases:
2Cr(OH)3 + 3H2SO4 → Cr2(SO4)3 + 6H2O;
Cr(OH)3 + 3NaOH → Na3[Cr(OH)6];
(sodium hexahydroxochromate(III))
∆ H O + 2NaCrO (sodium chromite).
Cr2O3 + 2NaOH →
2
2
3+
6+
In basic solutions Cr can be oxidized with oxidizing agents to Cr
(salts of chromic acid, H2CrO4):
2CrCl3 + 3Br2 + 16NaOH → 2Na2CrO4 + 6NaBr + 6NaCl + 8H2O.
The +6 oxidation state
In its highest oxidation state chromium has a high electronegativity and
behaves like a nonmetal. Chromium(VI) compounds are predominately
covalent.
Сhromic acid (H2CrO4) cannot be made, because when a yellow
2–
chromate solution (CrO4 ) is acidified, the solution becomes orange-red
2–
owing to the formation of the dichromate ion (Cr2O7 ):
2–
+
2–
2CrO4 + 2H3O → Cr2O7 + 3H2O.
We can consider this reaction as occurring by the protonation of the
–
chromate ion giving the hydrogen chromate ion (HCrO4 ) followed by
–
elimination of water from two HCrO4 ions:
O
O
O
O
O Cr OH + HO Cr O
O Cr O Cr O
+ H2O
O
O
O
O
This is an example of a condensation reaction. Potassium dichromate,
K2Cr2O7 crystallizes from solution as bright orange-red crystals. The
formation of the dichromate ion is reversible, and if base is added to a
dichromate solution, the color changes back to the yellow of the chromate
ion:
2–
–
2–
Cr2O7 + 2OH → 2CrO4 + H2O.
If potassium dichromate is dissolved in concentrated sulfuric acid, it
forms dichromic acid, which being dehydrated forms red solid chromium(VI)
oxide (chromium trioxide), CrO3:
30
Cr2O7 + 2H → H2Cr2O7;
H2Cr2O7 → 2CrO3 + H2O.
An acidic dichromate solution is often used in the laboratory as a strong
oxidizing agent:
2–
+
–
3+
Cr2O7 +14H + 6e → 2Cr + 7H2O.
For example,
K2Cr2O7 + 14HCl → 2CrCl3 + 3Cl2↑ + 2KCl + 7H2O ,
K2Cr2O7 + 3H2S + 4H2SO4 → Cr2(SO4)3 + 3S↓ + K2SO4 + 7H2O .
2–
+
H2Cr2O7 is the elementary representative of isopolyacids of chromium,
with which the common formula nCrO3·mH2O (where n > m) corresponds.
For example,
H2Cr3O10 – trichromic acid (n = 3, m = 1);
H2Cr4O13 – tetrachromic acid (n = 4, m = 1).
Chromium peroxides
When a solution of hydrogen peroxide is added to a solution of
potassium dichromate, it oxidizes to a blue solution of chromium peroxide
(CrO5):
K2Сr2O7 + 4H2O2 + H2SO4 → 2CrO5 + K2SO4 + 5H2O.
Chromium peroxide can be extracted easily from an aqueous solution
by ether.
Uses
More than a half of the chromium production goes into metallic products.
It is an ingredient in several important catalysts. The main usage of
chromium is to form alloys with iron, nickel, or cobalt. The addition of
chromium imparts hardness, strength, and corrosion resistance to the alloy.
In the stainless steels, chromium makes up 10 percent or more of the final
composition. Due to its hardness, an alloy of chromium, cobalt, and
tungsten is used for high-speed metal-cutting tools. When deposited
electrolytically, chromium provides a hard, corrosion-resistant, lustrous
finish. For this reason it is widely used as body trim on automobiles and
other vehicles. The extensive use of chromite as a refractory is based on its
high melting point, moderate thermal expansion, and the stability of its
crystalline structure. Lead chromate (PbCrO4), an insoluble solid is widely
used as a pigment called chrome yellow.
Chromium is an essential trace element and is of importance in a
glucose metabolism. It potentiates (enhances) insulin. Trivalent chromium
is an integral component of the glucose tolerance factor (GTF; a low
molecular weight compound with trivalent chromium coordinated to two
nicotinic acid molecules with the remaining coordinates protected by amino
acids) and acts as a cofactor for the hormone insulin. Apart from its vital
31
role in carbohydrate metabolism (i.e. glucose tolerance and glycogen
synthesis), trivalent chromium is also believed to play an important role in
cholesterol and amino acid metabolism.
All chromium compounds should be regarded as highly toxic.
Chromium(VI) compounds are highly toxic and carcinogenic. Chromium(III)
compounds are less toxic. Chromium compounds are important pollutants.
Molybdenum metal is used mainly in alloying steel. The alloy withstands
high temperatures and pressures and is very strong, making it useful for
structural work, aircraft parts, and forged automobile parts. Molybdenum
wire is used in electron tubes, and the metal serves as electrodes in glass
furnaces also. Molybdenum sulfide is used as a lubricant in environments
requiring high temperatures.
Molybdenum is a necessary element apparently for all species. Only
very small amounts are required. Molybdenum takes part in nitrogen
fixation enzymes and nitrate reduction enzymes. Molybdenum compounds
are encountered relatively rarely by most people. Unless known otherwise,
all molybdenum compounds should be regarded as highly toxic and as
teratogenic.
The principal uses of tungsten are as filaments in incandescent lamps,
as wires in electric furnaces, and in the production of hard, tenacious alloys
of steel. It is used also in the manufacture of spark plugs, electrical contact
points, and cutting tools, and as a target in X-ray tubes.
Tungsten has a limited biological role. A number of enzymes
(oxidoreductases) employ tungsten in a way related to molybdenum, (using
tungsten-protein complex). Tungsten metal does not normally cause
problems but all tungsten compounds should be regarded as highly toxic.
The metal dust presents a fire and explosion hazard.
5. Group VIIB. Manganese, Technetium, Rhenium
Group VIIB of the periodic table contains three chemical elements –
manganese Mn, technetium Tc, rhenium Re.
The atoms of these elements have the following configurations:
25Mn
2
2
6
2
6
5
2
– 1s 2s 2p 3s 3p 3d 4s
2
2
6
2
6
10
2
6
5
2
Tc
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s
43
2
2
6
2
6
10
2
6
10 14
2
6
5
2
75Re – 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s
5
2
[Ar]3d 2s
5
2
[Kr]4d 5s
5
2
[Xe]5d 6s
Table 5.1 lists some physical properties of group VIIB chemical
elements.
32
Table 5.1. Properties of Group VIIB Elements
Manganese
Technetium
Rhenium
Property
Electron configuration
Melting point, °C
Boiling point, °C
3
Density, g/cm
Covalent radius, Å
7+
Ionic radius (Me ), Å
5
[Ar]3d 2s
1245
1962
7.2
1.30
0.46
2
5
[Kr]4d 5s
2200
4567
11.5
1.36
0.56
2
5
[Xe]5d 6s
3390
5600
20.53
1.37
0.56
2
Physical Properties and Occurrence of the Elements
Manganese, symbol Mn, is a silvery white, brittle metallic element. The
atomic number of manganese is 25. Manganese was first distinguished as
an element in 1774 by the Swedish chemist Johan Gottlieb Gahn.
Manganese melts at about 1245 °C, boils at about 1962 °C, and has a
specific gravity of 7.2; the atomic weight of manganese is 54.938.
The metal does not occur in the free state, except in meteors, but is
widely distributed over the world in the form of ores, such as, pyrolusite
(manganese dioxide, MnO2) and franklinite (mineral containing the oxides
of iron, manganese, and zinc, with the formula (Fe, Zn, Mn)(Fe, Mn)2O4). It
th
ranks about 12 in abundance among elements in the earth's crust.
Technetium, symbol Tc, is a radioactive metallic element, the first
element to be created artificially. The atomic number of technetium is 43.
Technetium melts at about 2200 °C, boils at about 4567° C, and has a
specific gravity of 11.5.
In 1937 Emilio Segru and Carlo Perrier created technetium by
bombarding molybdenum targets with deuterons. Since technetium is not a
part of the decay series of any naturally radioactive element, scientists had
thought that technetium does not occur in nature. In 1988, however, minute
quantities of it were detected in ore from a deep molybdenum mine in
Colorado. Isotopes ranging in mass number from 90 to 111 are known; the
most common isotope has a mass number of 98.
Rhenium, symbol Re, is a silvery white, metallic element. The atomic
number of rhenium is 75. Rhenium metal is very hard; with the exception of
tungsten, it is the least fusible of all common metals. Overall, it ranks about
th
79 in natural abundance among elements in crystal rocks. Rhenium melts
at about 3180 °C , and has a specific gravity of 20.53. The atomic weight of
rhenium is 186.207.
The existence of rhenium and the similarity of its chemical properties to
those of the element manganese were predicted in 1871 by the Russian
chemist Dmitry Ivanovich Mendeleyev, who named it dvi-manganese.
Rhenium was discovered in 1925 by the German chemists Walter Karl
Noddack and Ida Eva Tacke Noddack in the ores tantalite, wolframite
33
((Fe,Mn)WO4), and columbite (mineral oxide of niobium, tantalum, iron, and
manganese (Fe,Mn)(Nb,Ta)2O6) by means of X-ray spectrographic
analysis, and it was found later in larger quantities in molybdenite
(molybdenum disulfide, MoS2).
Preparation of Manganese
Pure manganese is obtained by igniting pyrolusite (manganese dioxide,
MnO2) with aluminum powder:
3MnO2 + 4Al → 3Mn + 2Al2O3;
or by electrolyzing manganese sulfate:
2+
MnSO4 ⇄ Mn
On cathode:
2+
–
Mn + 2e → Mn
2–
+ SO4
On anode:
–
+
2H2O – 4e → 4H + O2
is
2MnSO4 + 2H2O electrolys


→ 2Mn + O2↑ + 2H2SO4.
Chemical Properties of Manganese
Oxidation states. The principal oxidation states of manganese are +2,
+3, +4,+6 and +7. The +7 oxidation state is the maximum possible
oxidation state for manganese, since in this state it uses all seven valence
electrons for the formation of bonds. Generally Manganese(VII) compounds
are strong oxidizing agents and are readily reduced to the +2 state, or to +4
state, which is the most stable oxidation state of manganese.
Manganese metal reacts directly with chlorine, nitrogen, oxygen and
water at heatig:
∆ MnCl ;
∆ Mn N ;
3Mn + N →
Mn + Cl →
2
2
2
3
2
∆ MnO ;
∆ Mn(OH) + H ↑.
Mn + 2O2 →
Mn + 2H2O →
2
2
2
With dilute acids (except nitric acid) manganese metal reacts releasing
2+
hydrogen gas and Mn ion:
Mn + 2HCl → MnCl2 + H2↑;
Mn + H2SO4 → MnSO4 + H2↑.
2+
Concentrated sulfuric and nitric acids oxidize manganese to Mn at
heating:
∆ 3Mn(NO ) + 2NO↑ + 4H O;
3Mn + 8HNO →
3
3 2
2
∆ MnSO + SO ↑ + 2H O.
Mn + 2H2SO4(conc.) →
4
2
2
The +2 oxidation state
Manganese(II) oxide (MnO) is a basic oxide and has all properties of
basic oxides:
MnO + 2HCl → MnCl2 + H2O;
MnO + SO3 → MnSO4.
34
Manganese(II) hydroxide (Mn(OH2)) has only basic properties and in
reactions with acids formes salts of manganese (II):
Mn(OH)2 + 2HCl → MnCl2 + 2H2O;
Mn(OH)2 + CO2 → MnCO3 + H2O.
Manganese(II) hydroxide is oxidized easily by O2 to manganese(IV)
hydroxide:
2Mn(OH)2 + O2 + 2H2O → 2Mn(OH)4 .
2+
6+
In basic solutions Mn can be oxidized with oxidizing agents to Mn
(salts of manganic acid, H2MnO4):
∆ 3K MnO + 6KCl + 3K SO + 6H O.
3MnSO + 2KClO + 12KOH →
4
3
2
2+
4
2
4
2
In acidic solutions Mn can be oxidized with strong oxidizing agents to
7+
Mn (permanganic acid, HMnO4):
2MnSO4 + 5PbO2 + 6HNO3 → 2HMnO4 + 3Pb(NO3)2 + 2PbSO4 + 2H2O.
2+
This reaction is the identification reaction of Mn .
The +4 oxidation state
Manganese dioxide (MnO2) reacts with strong bases and basic oxides
forming manganites (salts of manganous acid, H2MnO3):
∆ Na MnO ;
MnO + 2NaOH →
2
2
3
∆ CaMnO .
MnO2 + CaO →
3
Manganese dioxide can be oxidizing as well as reducing agents. If
4+
6+
7+
MnO2 reacts with oxidizing agents Mn it oxidized to Mn or Mn :
3MnO2 + KClO3 + 6KOH → 3K2MnO4 + KCl + 3H2O;
2MnO2 + 3PbO2 + 6HNO3 → 2HMnO4 + 3Pb(NO3)2 + 2H2O.
4+
2+
If MnO2 reacts with reducing agents Mn it reduce to Mn :
MnO2 + 4HCl → MnCl2 + Cl2↑ + 2H2O.
Manganese(IV) hydroxide (Mn(OH)4) has amphoteric properties in
reactions with acids and strong bases:
Mn(OH)4 + 4HCl → MnCl4 + 2H2O;
Mn(OH)4 + 4NaOH → 4H2O + Na4MnO4 (sodium orthomanganite).
The +6 oxidation state
2–
The manganate ion (MnO4 ) forms when MnO2 is heated with oxidizing
agents and strong bases:
2MnO2 + 4KOH + O2 → 2K2MnO4 + 2H2O.
Solutions containing manganates have a green color.
The manganate ion is stable in a basic solution only. In neutral and
–
acidic solutions it gives permanganates (MnO4 ) and manganese dioxide.
In this reaction the +6 oxidation state is self-oxidized and self-reduced to
the +7 oxidation state and the +4 state:
3K2MnO4 + 2H2O → 2KMnO4 + MnO2↓ + 4KOH;
3K2MnO4 + 2H2SO4 → 2KMnO4 + MnO2↓ + 2K2SO4 + 2H2O.
35
The +7 oxidation state
Manganese(VII) oxide (Mn2O7) is an acidic oxide and has all properties
of acidic oxides. It reacts with water, bases and basic oxides forming
permanganates (salts of permanganic acid, HMnO4):
Mn2O7 + H2O → 2HMnO4;
Mn2O7 + 2KOH → 2KMnO4 + H2O;
Mn2O7 + K2O → 2KMnO4 .
–
The permanganate ion (MnO4 ) has a deep purple color.
Permanganates decompose at heating up:
∆ K MnO + MnO + O ↑.
2KMnO →
4
2
4
2
2
2+
All permanganates are strong oxidizing agents and reduce to Mn in an
4+
6+
acidic solution, to Mn (MnO2) in a neutral solution or to Mn in a basic
solution:
2KMnO4 + 5K2SO3 + 3H2SO4 → 2MnSO4 + 6K2SO4 + 3H2O
pH<7
2KMnO4 + 3K2SO3 + H2O → 2MnO2↓ + 3K2SO4 + 2KOH
pH=7
pH>7.
2KMnO4 + K2SO3 + 2KOH → 2K2MnO4 + K2SO4 + H2O
Uses
Manganese is used mainly in the form of alloys with iron. The most
important of these alloys, which are used in steelmaking, are
ferromanganese, containing about 78 percent manganese, and
spiegeleisen, containing from 12 to 33 percent of manganese. Small
amounts of manganese are added to steel as a deoxidizer; large amounts
are used to produce a very tough alloy, resistant to wear. Safes, for
example, are made of manganese steel containing about 12 percent of
manganese. Nonferrous manganese alloys include manganese bronze
(composed of manganese, copper, tin and zinc), which resists corrosion
from seawater and is used for propeller blades on boats and torpedoes,
and manganin (containing manganese, copper, and nickel), used in the
form of wire for accurate electrical measurements because its electrical
conductivity does not vary appreciably with temperature.
Manganese dioxide is used in dry-cell batteries as a depolarizer, in paint
and varnish oils, for coloring glass and ceramics, and in preparing of
chlorine and iodine. Potassium permanganate forms dark purple crystals,
which are used as oxidizers in analytical chemistry and disinfectants in
medicine.
Manganese compounds are essential for life. They are necessary for
the action of those enzymes that mediate phosphate group transfer (ie.
phosphate transferases and phosphate dehydrogenases), particularly
those concerned with the citric acid cycle including arginase, alkaline
phosphatase and hexokinase. It is an essential component of the enzyme
pyruvate carboxylase. As a cofactor or component of several key enzyme
systems, manganese
is essential for bone formation (i.e.
36
mucopolysaccharide synthesis), the regeneration of red blood cells,
carbohydrate metabolism, and the reproductive cycle.
Soil deficiencies lead to infertility in mammals and to bone deformation
in growing chicks.
Compounds and alloys containing technetium can prevent the corrosion
of iron caused by water. Technetium-99 is used for imaging in medicine.
Rhenium is used in electrical filaments, welding rods, thermocouples,
cryogenic magnets, and photographic flashbulb filaments; it is also used as
a catalyst.
6. Group VIIIB. Iron, Cobalt, Nickel and the Platinum
Metals
Group VIIIB of the periodic table contains nine chemical elements – iron
Fe, cobalt Co, nickel Ni and the platinum metals: platinum Pt, ruthenium
Ru, rhodium Ph, palladium Pd, osmium Os and iridium Ir.
Atoms of these elements have the following configurations:
26Fe
2
2
6
2
6
6
2
44Ru
2
2
6
2
6
10
2
6
7
1
2
2
6
2
6
10
2
6
10
14
– 1s 2s 2p 3s 3p 3d 4s
2
2
6
2
6
7
2
27Co – 1s 2s 2p 3s 3p 3d 4s
2
2
6
2
6
8
2
– 1s 2s 2p 3s 3p 3d 4s
28Ni
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s
2
2
6
2
6
10
2
6
8
1
45Rh – 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s
2
2
6
2
6
10
2
6
10
0
46Pd – 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s
76Os
77Ir
78Pt
6
2
7
1
[Ar]3d 2s
7
2
[Ar]3d 2s
8
2
[Ar]3d 2s
[Kr]4d 5s
8
1
[Kr]4d 5s
10
0
[Kr]4d 5s
2
6
6
2
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s
2
2
6
2
6
10
2
6
10 14
2
6
7
2
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s
2
2
6
2
6
10
2
6
10 14
2
6
9
1
– 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s
6
2
[Xe]5d 6s
7
2
[Xe]5d 6s
9
1
[Xe]5d 6s
Table 6.1 lists some physical properties of group VIIIB chemical
elements.
Table 6.1. Properties of Group VIIIB Elements
Elements
Iron
Cobalt
Nickel
Ruthenium
Rhodium
Palladium
Osmium
Iridium
Platinum
Melting
point, °C
Boiling
point, °C
1535
1495
1455
2310
1966
1554
2700
2410
1772
2750
2870
2730
3900
3727
2970
5000
4130
3827
Property
Density,
Covalent
3
o
g/cm
radius, A
1.17
7.86
1.16
8.9
1.15
8.9
1.34
12.3
1.34
12.4
1.37
12.02
1.35
22.6
1.35
22.6
1.38
21.45
37
Ionic radius
2+
o
(Me ), A
0.75
0.79
0.83
0.80
0.80
0.88
0.89
0.89
0.90
Physical Properties and Occurrence of Elements
Iron, (Latin ferrum, “iron”) symbol Fe, is a magnetic, malleable, silvery
white metallic element. The atomic number of iron is 26. Pure iron has a
hardness that ranges from 4 to 5. It is soft, malleable, and ductile. Iron is
easily magnetized at ordinary temperatures; it is difficult to magnetize it
when heated, and at about 790 °C the magnetic property disappears. Pure
iron melts at about 1535 °C, boils at 2750 °C, and has a specific gravity of
7.86. The atomic weight of iron is 55.847.
The metal exists in three different forms: ordinary, or α-iron (alpha-iron);
γ-iron (gamma-iron); and δ-iron (delta-iron). The internal arrangement of the
atoms in the crystal lattice changes in the transition from one form to
another. The transition from α-iron to γ-iron occurs at about 910 °C, and the
transition from γ-iron to δ-iron occurs at about 1400 °C.
Metallic iron was known and used for ornamental purposes and
weapons in prehistoric ages; the earliest specimen still extant, a group of
oxidized iron beads found in Egypt, dates from about 4000 BC.
Metallic iron occurs in the free state in a few localities only, notably
western Greenland. It is found in meteorites, usually alloyed with nickel. In
chemical compounds the metal is widely distributed and ranks fourth in
abundance among all the elements in the earth's crust; next to aluminium it
is the most abundant of all metals. The principal ore of iron is hematite
(Fe2O3). Other important ores are magnetite (Fe3O4), siderite (FeCO3), bog
iron (limonite) (Fe2O3·nH2O) and pyrite (FeS2). Small amounts of iron occur
in combination in natural waters, in plants, and as a constituent of blood.
Cobalt, symbol Co, is a silvery-white, magnetic, metallic element. The
atomic number of cobalt is 27. Cobalt was discovered in 1735 by the
Swedish chemist George Brandt. It has a relatively low strength and little
ductility at normal temperatures, but is ductile at high temperatures. Cobalt
melts at about 1495 °C, boils at about 2870 °C, and has a specific gravity
of 8.9; the atomic weight of cobalt is 58.933.
th
Cobalt is about the 30 most abundant element in crustal rocks. Cobalt
occurs as the arsenide (CoAs2), known as smaltite or speiss cobalt; as
cobalt sulfarsenide (CoAsS), known as cobalt glance or cobaltite; and as a
hydrated arsenate of cobalt (Co(AsO4)2 · 8H2O), known as cobalt bloom or
erythrite.
Nickel, symbol Ni, is a silvery white, magnetic metallic element. The
atomic number of nickel is 28. Nickel is a hard, malleable, ductile metal,
capable of taking a high polish. It is magnetic below 345 °C. Nickel melts at
about 1455 °C, boils at about 2730 °C, and has a specific gravity of 8.9.
The atomic weight of nickel is 58.69.
38
Nickel was used as coinage in nickel-copper alloys for several thousand
years, but was not recognized as an elemental substance until 1751 when
the Swedish chemist Baron Axel Frederic Cronstedt isolated the metal from
niccolite ore.
Nickel occurs as a metal in meteors. Combined with other elements, it
nd
occurs in a mineral garnierite ((Ni,Mg)SiO3·nH2O). Nickel ranks about 22
in natural abundance among elements in crustal rock.
The platinum metals
Ruthenium, symbol Ru, is a grayish-white metallic element. The atomic
number of ruthenium is 44. Ruthenium was discovered in 1844 by the
Russian chemist Karl Karlovich Klaus. The name of the element is derived
from the region of Ruthenia, which now is a part of Ukraine. Ruthenium
melts at about 2310 °C, boils at about 3900 °C, and has a specific gravity
of 12.3. The atomic weight of ruthenium is 101.07.
The metal occurs in the metallic state in platinum ores. Ruthenium ranks
th
80 in natural abundance among elements in crustal rocks.
Rhodium (Greek rhodon, “rose”), symbol Rh, is a brilliant silvery white
metallic element. The atomic number of rhodium is 45. Rhodium was
discovered in 1803 by the British chemist William Hyde Wollaston.
Rhodium melts at about 1966 °C, boils at about 3727 °C, and has a specific
gravity of 12.4. The atomic weight of rhodium is 102.905.
The metal occurs as an alloy in platinum ores, in osmiridium, and in
st
gold-rhodium ores called rhodite. It ranks 81 in order of abundance of the
elements in the crust of the earth.
Palladium, symbol Pd, is a relatively rare, silvery white, soft metallic
element. The atomic number of palladium is 46. Palladium was discovered
in 1804 by the British chemist William Hyde Wollaston. Palladium has a
hardness of 4.8. Finely divided palladium is an excellent adsorbent for
some gases; it adsorbs hydrogen or acetylene gas from 1000 to 3000 times
of its volume when heated to 100 °C. Palladium melts at about 1554 °C,
boils at about 2970 °C, and has a specific gravity of 12.02. The atomic
weight of Palladium is 106.4.
st
Palladium ranks about 71 in natural abundance among the elements in
crustal rocks. The metal occurs in the pure state in platinum ores and in the
combined state in Canadian nickel ore.
Osmium, symbol Os, is a bluish-white, brittle metallic element. Osmium
is generally considered to be the most dense elements. The atomic number
of osmium is 76. Osmium was discovered in 1803 by the British chemist
Smithson Tennant. Osmium has a hardness of 7; it melts at 2700 °C and
has a specific gravity of 22.6. The atomic weight of osmium is 190.2.
The metal occurs naturally in platinum ores and as an alloy, osmiridium,
th
with iridium. Osmium ranks about 74 in natural abundance among the
elements in crustal rocks.
39
Iridium, symbol Ir, is a white, brittle, extremely hard, metallic element.
The atomic number of iridium is 77. Iridium was discovered by the British
chemist Smithson Tennant in 1804 and was named for the iridescent
nature of some of its compounds. The atomic weight of iridium is 192.22. It
melts at about 2410 °C and boils at about 4130 °C, and has a specific
gravity of 22.6.
th
It is an extremely rare metal, ranking 77 in order of abundance of the
elements in the crust of the earth. Iridium is found in alluvial deposits
alloyed with platinum as platiniridium and with osmium as osmiridium.
Platinum, symbol Pt, is a grayish-white metal with a hardness of 4.3. It
has a high fusing point, is malleable and ductile, expands slightly upon
heating, and has high electrical resistance. The atomic number of platinum
is 78. Platinum melts at about 1772 °C, boils at about 3827 °C, and has a
specific gravity of about 21.45. The atomic weight of platinum is 195.09.
Platinum were probably used in alloyed forms in ancient Greece and
Rome and were first mentioned in European literature in the early 16th
nd
century. Platinum ranks about 72 in natural abundance among the
elements in crustal rocks. Except the mineral sperrylite, which is platinum
arsenide and is found only sparingly in a few localities, platinum occurs in
the metallic state, often alloyed with other platinum metals.
Preparation of the Elements
The basic materials used for the manufacture of iron are iron ore, coke,
and limestone. The coke is burned as a fuel to heat the furnace; as it burns,
the coke gives off carbon monoxide:
2C + O2 → CO↑;
which combines with the iron oxides in the ore, reducing them to metallic
iron. This is the basic chemical reaction in the blast furnace; its equation is:
∆ 3CO ↑ + 2Fe.
Fe O + 3CO →
2
3
2
The limestone in the furnace charge is used as an additional source of
carbon monoxide and as a “flux” combining with the infusible silica present
in the ore to form fusible calcium silicate:
∆ CaSiO + CO ↑.
CaCO + SiO →
3
2
3
2
Without the limestone, iron silicate will form, with a resulting loss of
metallic iron. Calcium silicate plus other impurities form a slag that floats on
the top of the molten metal at the bottom of the furnace.
Small amounts of pure iron can be made through the purification of
crude iron with carbon monoxide. The intermediate in this process is iron
pentacarbonyl, Fe(CO)5. The carbonyl decomposes at heating to about
250°C forminf pure iron powder:
Fe + CO → Fe(CO)5;
T = 250 °C
Fe(CO)5 
 → Fe + 5CO↑.
40
Fe(CO)5 is a volatile oily complex which is easily flushed from the
reaction vessel leaving the impurities behind. Other way to obtaining small
samples of pure iron includes the reduction of iron oxide, Fe2O3, with
hyrogen, H2:
∆ 2Fe + 3H O.
Fe O + 3H →
2
3
2
2
Pure cobalt and nickel prepared by reduction of their oxides:
− by aluminum: 3CoO + 2Al → 3Co + Al2O3;
− by carbon: NiO + C → Ni + CO;
− or by electrolyzing of sulfates:
2+
CoSO4 ⇄ Co
On cathode:
2+
–
Co + 2e → Co
2–
+ SO4
On anode:
–
+
2H2O – 4e → 4H + O2
is
2CoSO4 + 2H2O electrolys


→ 2Co + O2↑ + 2H2SO4.
Chemical Properties of Iron
Oxidation states. The principal oxidation states of iron are +2, +3 and
+6. Generally iron(II) compounds are reducing agents and readily oxidizes
to the +3 state, which is the most stable oxidation state of iron.
Iron metal reacts directly with chlorine, nitrogen, oxygen and water at
heating:
∆ 2FeCl ;
∆ Fe N ;
2Fe + 3Cl →
3Fe + N →
2
3
2
3
2
∆ 2Fe O ;
∆ Fe(OH) + H ↑.
4Fe + 3O2 →
Fe + 2H2O →
2 3
2
2
With dilute acids (except nitric acid) iron metal reacts releasing
2+
hydrogen gas and Fe ions:
Fe + 2HCl → FeCl2 + H2↑;
Fe + H2SO4 → FeSO4 + H2↑.
3+
Concentrated sulfuric and dilute nitric acids oxidize iron to Fe at
heating:
∆ Fe(NO ) + NO↑ + 2H O;
Fe + 4HNO (dil) →
3
3 3
2
∆ Fe (SO ) + 3SO ↑ + 6H O .
2Fe + 6H2SO4(conc.) →
2
4 3
2
2
When iron is dipped into concentrated nitric acid, it forms a layer of
oxide that renders it passive – that is, it does not react chemically with
acids or other substances.
The +2 oxidation state
Iron(II) oxide (FeO), an amorphous black powder, is a basic oxide and
has all properties of basic oxides:
FeO + H2SO4 → FeSO4 + H2O;
FeO + CO2 → FeCO3.
41
Iron(II) hydroxide (Fe(OH) 2) has basic properties only and in reactions
with acids formes salts of iron(II):
Fe(OH)2 + 2HCl → FeCl2 + 2H2O;
Fe(OH)2 + SO3 → FeSO4 + H2O.
Iron(II) hydroxide is easily oxidized by O2 to iron(III) hydroxide:
4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3.
2+
3+
In solutions Fe can be oxidized with oxidizing agents to Fe :
10FeSO4 + 2KMnO4 + 8H2SO4 → 5Fe2(SO4)3 + 2MnSO4 + K2SO4 + 8H2O.
2+
The Fe
ions combine with cyanides forming complex cyanide
compounds:
FeSO4 + 6KCN → K4[Fe(CN)6] + K2SO4.
Potassium hexacyanoferrate(II) (K4[Fe(CN)6]), is called yellow prussiate
of potash.
2+
The reaction of identification of the Fe ion
2+
Potassium hexacyanoferrate(III) (K3[Fe(CN)6]) reacts with the Fe ion
forming a dark-blue, amorphous precipitate:
2+
3+
FeSO4 + K3[Fe(CN)6] → K Fe [ Fe (CN)6]↓ + K2SO4.
Iron(II)-potassium hexacyanoferrate(III) (KFe[Fe(CN)6]) is
Prussian blue.
called
The +3 oxidation state
Iron(III) oxide (Fe2O3), an amorphous red powder, has amphoteric
properties. It reacts with acids and acidic oxides:
Fe2O3 + 6HCl → 2FeCl3 + 3H2O;
Fe2O3 + 3SO3 → Fe2(SO4)3 .
Strong bases, basic oxides and some salts react with Fe2O3 and form
ferrites (salts of ferrous acid, HFeO2):
∆ 2NaFeO + H O;
Fe O + 2NaOH →
2
3
2
2
∆ Ca(FeO ) ;
Fe2O3 + CaO →
2 2
∆
Fe2O3 + K2CO3 → 2KFeO2 + CO2↑.
The ferrite ion is unstable in an aqueous solution and hydrolyzes:
2NaFeO2 + H2O → 2NaOH + Fe2O3↓ .
Iron(III) hydroxide (Fe(OH)3) has amphoteric properties in reactions with
acids and concentrated strong bases:
Fe(OH)3 + 3HCl → FeCl3 + 3H2O;
Fe(OH)3 + 3NaOH → Na3[Fe(OH)6].
3+
2+
All salts contain Fe are strong oxidizing agents and reduce to Fe :
2FeCl3 + 2KI → 2FeCl2 + I2 + 2KCl.
3+
The Fe
ions reacted with cyanides forming complex cyanide
compounds:
FeСl3 + 6KCN → K3[Fe(CN)6] + 3KCl .
42
Potassium hexacyanoferrate(III) (K3[Fe(CN)6]) is called red prussiate of
potash.
3+
The reaction of identification of the Fe ion
3+
1. Potassium hexacyanoferrate(II) (K4[Fe(CN)6]) reacts with the Fe ion
forming a dark-blue, amorphous precipitate:
3+
2+
FeCl3 + K4[Fe(CN)6] → K Fe [ Fe (CN)6]↓ + 3KCl .
Iron(III)-potassium hexacyanoferrate(II) (KFe[Fe(CN)6]) is called
Turnbull's blue.
3+
2. Ammonia thiocyanide (NH4SCN) reacts with the Fe ion forming a
dark-red solution:
FeCl3 + 3NH4SCN → Fe(SCN)3 + 3NH4Cl .
dark-red
The +6 oxidation state
2–
The ferrate ion (FeO4 ) form when Fe2O3 heat with oxidizing agents
and strong bases
∆ 2K FeO + 3KNO + 2H O.
Fe2O3 + 3KNO3 + 4KOH →
2
4
2
2
Chemical Properties of Cobalt and Nickel
Oxidation states. The principal oxidation states of cobalt and nickel are
+2 and +3. Generally cobalt(III) and nickel(III) compounds are strong
oxidizing agents and readily reduces to the +2 state, which is the most
stable oxidation state of cobalt and nickel.
Cobalt and nickel metals react directly with chlorine and oxygen at
heating:
∆ CoCl ;
∆ NiCl ;
Co + Cl →
Ni + Cl →
2
2
2
2
∆ 2CoO;
∆ 2NiO.
2Co + O2 →
2Ni + O2 →
With dilute strong acids (except nitric acid) cobalt and nickel metals
react releasing hydrogen gas:
Co + 2HCl → CoCl2 + H2↑;
Ni + 2HCl → NiCl2 + H2↑;
Ni + H2SO4 → NiSO4 + H2↑.
Co + H2SO4 → CoSO4 + H2↑;
Dilute nitric acid oxidizes cobalt and nickel at heating:
∆ 3Co(NO ) + 2NO↑ + 4H O;
3Co + 8HNO (dil.) →
3
3 2
2
∆ 3Ni(NO ) + 2NO↑ + 4H O.
3Ni + 8HNO3(dil.) →
3 2
2
Concentrated nitric acid turns cobalt and nickel to become passive.
The +2 oxidation state
Cobalt(II) oxide (CoO) and cobalt(II) hydroxide (Co(OH)2) have
amphoteric properties.
Cobalt(II) hydroxide is oxidized by strong oxidizing agents only to
cobalt(III) hydroxide:
2Co(OH)2 + NaClO + H2O → 2Co(OH)3 + NaCl.
43
2+
The Co ions combine with ammonia, cyanides, thiocyanides and
hydroxide-ions forming complex compounds:
CoCl2 + 8NH4OH → [Co(NH3)6](OH)2 + 2NH4Cl;
CoCl2 + 4KCN → K2[Co(CN)4] + 2KCl;
CoCl2 + 4KSCN → K2[Co(SCN)4] + 2KCl;
CoCl2 + 4KOH → K2[Co(OH)4] + 2KCl.
Nickel(II) oxide (NiO) and nickel(II) hydroxide (Ni(OH)2) have basic
properties only.
2+
The Ni ions combine with ammonia, cyanides and thiocyanides
forming complex compounds:
NiCl2 + 8NH4OH → [Ni(NH3)6](OH)2 + 2NH4Cl;
NiCl2 + 4KCN → K2[Ni(CN)4] + 2KCl;
NiCl2 + 4KSCN → K2[Ni(SCN)4] + 2KCl.
Nickel(II) compounds are often identified by adding an organic reagent,
dimethylgloxime, which reacts with nickel forming a red, flocculent
precipitate:
N
H3C C N OH
H3C C N OH
+ Ni
2+
+ 2NH3
O
H3C C
N
C CH3 + 2NH4+
Ni
C
H3C
O
N
N
OH OH
C
CH3
The +3 oxidation state
Cobalt(III) and nickel(III) oxides (Co2O3, Ni2O3) are strong oxidizing
agents:
Co2O3 + 6HCl → 2CoCl2 + Cl2↑ + 3H2O;
Ni2O3 + 6HCl → 2NiCl2 + Cl2↑ + 3H2O;
2Co2O3 + 4H2SO4 → 4CoSO4 + O2↑ + 4H2O;
2Ni2O3 + 4H2SO4 → 4NiSO4 + O2↑ + 4H2O.
Complex compounds of cobalt(III) can be prepared when solution of
2+
Co reacts with oxidizing agents and ligands:
2+
+
3+
4Co + 4NH4 + 20NH3 + O2 → 4[Co(NH3)6] + 2H2O.
Chemical Properties of the Platinum Metals
Chemically the platinum metal is relatively inert and resists attack by air,
1
water, single acids and ordinary reagents. It dissolves slowly in aqua regia
only, forming chloroplatinic acid (H2[PtCl6]):
3Pt + 18HCl + 4HNO3 →3H2[PtCl6] + 4NO↑ + 8H2O.
The ruthenium and rhodium metals are superior to platinum in
resistance to attack by acids, including aqua regia.
1
Aqua Regia (Latin, “royal water”), mixture of concentrated hydrochloric and nitric acids,
containing one part by volume of nitric acid (HNO3) to three parts of hydrochloric acid (HCl).
44
Palladium dissolves readily in aqua regia:
3Pd + 18HCl + 4HNO3 → 3H2[PdCl6] + 4NO↑ + 8H2O;
and slowly in concentrated nitric and sulfuric acids:
Pd + 4HNO3 → Pd(NO3)2 + 2NO2↑ + 2H2O.
∆ PdSO + SO ↑ + 2H O.
Pd + 2H SO →
2
4
4
2
2
Osmium is not attacked by ordinary acids, but dissolves in aqua regia:
Os + HCl + HNO3 → H2[OsCl6] + NO + 8H2O;
or fuming nitric acid:
Os + HNO3 → OsO4 + 8NO2↑ + 4H2O.
Iridium is extremely inert chemically, resisting even the action of aqua
regia.
Uses
Commercial iron invariably contains small amounts of carbon and other
impurities that alter its physical properties, which are considerably
improved by the further addition of carbon and other alloying elements.
By far the greatest amount of iron is used in processed forms, such as
wrought iron, cast iron, and steel. Commercially pure iron is used for the
production of galvanized sheet metal and of electromagnets. Iron
compounds are used for medicinal purposes in the treatment of anemia,
when the amount of hemoglobin or the number of red blood corpuscles in
the blood is lowered. Iron is used in tonics also.
Iron is an essential component of the respiratory pigments haemoglobin
and myoglobin. It is an essential component of various enzyme systems
including the cytochromes, catalases, peroxidases, and the enzymes
xanthine and aldehyde oxidase, and succinic dehydrogenase. As a
component of the respiratory pigments and enzymes concerned in tissue
oxidation, iron is essential for oxygen and electron transport within the
body.
Thermally resistant alloys, called superalloys, containing cobalt are used
in industry and aircraft gas turbine engines. An alloy with steel known as
cobalt steel is used for making permanent magnets. With tungsten carbide,
cobalt forms Carboloy, a hard material used for cutting and machining
steel; alloyed with chromium, cobalt produces Stellite, used for the same
purpose. Cobalt is used also in ceramics and paint driers, and as a
catalyst. The radioactive cobalt-60 is the most important. It has a half-life of
5.7 years and produces intensive gamma radiation. Cobalt-60 is used
extensively in industry and in radioisotope therapy.
Cobalt salts in small amounts are essential for many species, including
humans. It is at the core of vitamin-B12.
Nickel is used as a protective and ornamental coating for metals,
particularly iron and steel, that are susceptible to corrosion. The nickel plate
is deposited by electrolysis in a nickel solution. Finely divided nickel
45
absorbs 17 times of its own volume of hydrogen and is used as a catalyst
in many processes, including the hydrogenation of oils.
Nickel is used mainly in the form of alloys. It imparts great strength and
corrosion resistance to steel. Nickel steel, containing about 2 to 4 percent
of nickel, is used in automobile parts such as axles, crankshafts, gears,
valves, and rods; in machine parts; and in armor plate. Some of the most
important nickel-containing alloys are German silver, Invar, Monel metal,
Nichrome, and Permalloy. The nickel coins used for currency are an alloy
of 25 percent of nickel and 75 of percent copper. Nickel is also a key
component of nickel-cadmium batteries.
Nickel is an essential trace element for many species. Chicks and rats
raised on nickel-deficient diets show liver damage. Nickel is a key metal for
several hydrogenases and plant ureases.
Due to its chemical inertness and high fusing point, platinum is valuable
for laboratory apparatus, such as crucibles, tongs, funnels, combustion
boats, and evaporating dishes. Small amounts of iridium are usually added
to increase its hardness and durability. Platinum is used also for contact
points in electrical apparatus and in instruments used for measuring high
temperatures. Finely divided platinum in the form of platinum sponge or
platinum black is used extensively as a catalyst in the chemical industry.
The addition of ruthenium to platinum and palladium alloys makes them
very hard. Such alloys have a high resistance to wear and are used in the
manufacture of jewelry, in porcelain-metal restorations in dentistry, as tips
for fountain-pen nibs, and for nonmagnetic instrument pivots. The
ruthenium-molybdenum alloy is a superconductor at temperatures below
-263 °C.
Rhodium is used mostly as an alloy with platinum; the resulting alloy has
the desirable properties of platinum and is hard and durable also. Rhodiumplatinum alloys are used in thermocouples, measuring high temperatures.
Pure rhodium is used as a mirror surface in searchlights and as a plating
finish for jewelry and silverware. Rhodium black is a finely divided metal
that contains some oxide and hydride. It is used both as a catalyst and as a
black pigment for porcelain ware.
The main use of palladium is in the field of communications, where it is
used to face electrical contacts in automatic switchgear. It is used also in
dentistry; for nonmagnetic springs in clocks and watches; for coating
special mirrors; and in jewelry, alloyed with gold, in what is called white
gold.
The main use of osmium is in the alloy osmiridium. Alloyed with
platinum, it is used for standard weights and measures.
Iridium is used mainly as an alloying material for platinum; the alloy,
which contains about 10 percent of iridium, is much harder than pure
platinum. Platinum-iridium alloys containing larger percentages of iridium
46
are used in making precision instruments, surgical tools, pen points, and
standard weights and lengths.
7. Group IB. Copper, Silver and Gold
Group IB of the periodic table contain three chemical elements – copper
Cu, silver Ag and gold Au.
Atoms of these elements have the following configurations:
29Cu
2
2
6
2
6
10
1
– 1s 2s 2p 3s 3p 3d 4s
2
2
6
2
6
10
2
6
10
1
47Ag – 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s
2
2
6
2
6
10
2
6
10 14
2
6
10
1
79Au – 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s
10
1
[Ar]3d 2s
10
1
[Kr]4d 5s
10
1
[Xe]5d 6s
Table 7.1 lists some physical properties of group IB chemical elements.
Table 7.1. Properties of Group IB Elements
Copper
Silver
Property
Electron configuration
Melting point, °C
Boiling point, °C
3
Density, g/cm
o
Covalent radius, A
+
o
Ionic radius (Me ), A
10
[Ar] 3d 2s
1245
1962
1
10
1
Gold
10
1
[Xe] 5d 6s
3390
5600
20.53
8.9
[Kr] 4d 5s
2200
4567
11.5
1.28
1.44
1.44
0.96
1.66
1.37
Physical Properties and Occurrence of the Elements
Copper, symbol Cu, is a brownish-red metallic element. The atomic number
of copper is 29. Copper melts at about 1083 °C, boils at about 2567 °C,
and has a specific gravity of 8.9. The atomic weight of copper is 63.546.
Copper was known to prehistoric people and was probably the first
th
metal from which useful articles were made. Сopper is about the 25 most
abundant element in crustal rocks. It is usually found admixed with other
metals, such as gold, silver, bismuth, and lead, and exists in small specks
in rock. The principal source of copper is chalcopyrite (CuFeS2). Other
important ore minerals are chalcocite (Cu2S), azurite or malachite (basic
carbonate of copper, (CuOH)2CO3).
Silver, symbol Ag, is a white, lustrous metallic element. The atomic
number of silver is 47. Silver melts at about 962 °C, boils at about 2212 °C,
and has a specific gravity of 10.5. The atomic weight of silver is 107.868.
Silver has been known and valued as an ornamental and coinage metal
th
since ancient times. Silver ranks about 66 among elements in natural
abundance in crustal rocks. It occurs in the pure state to a small extent.
Pure silver is found associated with pure gold in the form of an alloy known
47
as electrum, and considerable amounts are recovered in the processing of
gold. Silver is usually found combined with sulfur in argentite (Ag2S).
Gold, symbol Au (from Latin aurum, “gold”), is a soft, dense, bright
yellow metallic element. Its atomic number is 79. Gold melts at about
1064 °C, boils at about 2808 °C, and has a specific gravity of 19.3; its
atomic weight is 196.97.
Gold is found in nature in quartz veins and secondary alluvial deposits
as a free metal or in a combined state. It is widely distributed although it is
th
rare, being 75 in order of abundance of the elements in the crust of the
earth. It is almost always associated with varying amounts of silver.
Preparation of the Elements
The copper ore chalcopyrite (CuFeS2) is first roasted in a limited supply
of air. This roasting converts the iron to FeO and removes some of the
sulfur compounds like SO2 but leaves the copper as Cu2S:
2CuFeS2 + 4O2 → Cu2S + 2FeO + 3SO2↑.
The FeO forms a slag of FeSiO3 with any silica that is present or is
added, and this slag is removed, leaving a relatively pure molten copper(I)
sulfide, Cu2S. Then air is blown through the molten Cu2S, converting it to
the metal:
Cu2S + O2 → 2Cu + SO2↑.
This copper is further purified by electrolysis.
Silver is usually recovered from silver ores (argentite, Ag2S) by roasting
the ore in a furnace to convert the sulfides to sulfates and then chemically
precipitating metallic silver:
Ag2S + 2O2 → Ag2SO4.
Silver and gold are usually recovered by lixiviation methods. Metals are
dissolved in a solution of a salt, usually sodium cyanide:
4Ag + 8NaCN + O2 + 2H2O → 4Na[Ag(CN)2] + 4NaOH;
4Au + 8NaCN + O2 + 2H2O → 4Na[Au(CN)2] + 4NaOH,
after that metals are precipitated by bringing the solution in contact with
metallic zinc or aluminium:
2Na[Ag(CN)2] + Zn → Na2[Zn(CN)4] + 2Ag;
2Na[Au(CN)2] + Zn → Na2[Au(CN)4] + 2Ag.
Several metallurgical processes are used to extract silver and gold from
ores of other metals. In the amalgamation process, liquid mercury, which
forms an amalgam with the silver or gold, is added to the crushed ore. After
the amalgam is washed out of the ore, the mercury is removed by
distillation, leaving metallic silver or gold.
Chemical Properties of Copper
Copper has two common oxidation states, +1 and +2. The compounds
+
2+
of copper are predominately ionic and contain Cu or Cu . The electron
48
10
1
+
configuration of copper is [Ar]3d 4s . It losses the 4s electron giving Cu ,
2+
and it can lose also one of the 3d electrons to give Cu .
Copper metal reacts directly with chlorine and oxygen only at heating:
∆ CuCl ;
∆ 2CuO.
Сu + Cl →
2Cu + O →
2
2
2
With dilute acids (except nitric acid) copper metal does not react.
Concentrated sulfuric and nitric acids react with copper at heating:
∆ CuSO + SO ↑ + 2H O;
Сu + 2H SO (conc.) →
2
4
4
2
2
∆ 3Cu(NO ) + 2NO↑ + 4H O;
3Cu + 8HNO3(dil) →
3 2
2
∆
Cu + 4HNO (conc.) → Cu(NO ) + 2NO ↑ + 2H O.
3
3 2
2
2
The +1 oxidation state
When CuO is heated strongly with copper filings or copper powder, red
copper(l) oxide, Cu2O, formes:
∆ Cu O.
CuO + Cu →
2
Black copper(l) sulfide, Cu2S, can be obtained simply by heating
copper(ll) sulfide, CuS:
∆ Cu S + S.
2CuS →
2
It can be made also by heating the requisite amounts of copper and sulfur:
∆ Cu S.
2Cu + S →
2
White copper(I) chloride, CuCI, is formed when HCI gas is passed over
heated copper:
∆ 2CuCI + H ↑.
2Cu + 2HCI(g) →
2
2+
–
Copper(l) iodide can be obtained by the reduction of Cu with I in a
reaction:
2CuSO4 + 4KI → 2Cul↓ + l2 + 2K2SO4.
Copper(l) iodide is insoluble in water. The only copper(l) compounds that
can be made in aqueous solution are insoluble compounds like Cul,
+
2+
because Cu is not stable in aqueous solution; it reacts giving Cu and Cu :
2CuCl → CuCl2 + Cu↓.
The +2 oxidation state
When copper is heated strongly in air or oxygen, black copper(ll) oxide,
CuO, formes:
∆ 2CuO.
2Cu + O →
2
It can be prepared also by heating copper(ll) nitrate:
∆ 2CuO + 4NO ↑ + O ↑.
2Cu(NO3)2 →
2
2
Copper(ll) nitrate can be obtained as blue hydrated crystals,
Cu(NO3)2×6H2O, by evaporating a solution made by dissolving copper in
nitric acid. The most common salt of copper(ll) is copper(II) sulfate, CuSO4.
It can be prepared by the reaction of CuO with dilute sulfuric acid:
CuO + H2SO4 → CuSO4 + H2O.
49
It crystallizes from solution as large blue crystals of the hydrate
CuSO4×5H2O. Four of the water molecules of copper(ll) sulfate
2+
pentahydrate are associated with Cu , and the fifth is hydrogen-bonded to
the sulfate ion as well as to the water molecules on the copper ion. We can
write the formula as [Cu(H2O)4]SO4×H2O to represent its structure better.
White CuSO4 formed when the hydrate is heated strongly:
∆ CuSO + 5H O,
CuSO ×5H O →
4
2
4
2
It is called anhydrous copper(ll) sulfate. It can be used for detection small
amounts of water in solvents such as alcohol and ether; when it is added to
these liquids, it becomes blue if water is present because the hydrate
CuSO4×5H2O is forming.
Copper(ll) sulfide is a black solid which formes by heating copper with
2+
the excess of sulfur or by passing H2S into a solution of a Cu salt:
∆ CuS;
Cu + S →
CuSO4 + H2S → CuS↓ + 2H2SO4.
–
Copper(ll) hydroxide formes as a pale blue precipitate when OH is
2+
added to a solution of Cu salt:
CuSO4 + 2KOH → Cu(OH)2 + K2SO4.
This precipitate dissolves in an aqueous NH3 solution giving a deep blue
2+
solution of the tetraaminecopper(ll) ion, [Cu(NH3)4] :
Cu(OH)2 + 4NH4OH → [Cu(NH3)4](OH)2 + 4H2O,
and in a concentrated solution of sodium hydroxide giving a deep blue
solution of the tetrahydroxocuprate(ll) ion:
Cu(OH)2 + 2NaOH → Na2[Cu(OH)4].
lf Cu(OH)2 is heated, black CuO formes:
∆ CuO + H O .
Cu(OH) →
2+
2
2
ion formes stable chelates with aminoacids and ethylene
H
O
The Cu
glycol:
H2C
Cu
2+
O
+ CH2(OH)CH 2(OH)
O
CH2
O
CH2
Cu
O
H
O
H2C
O
CH2
Cu
2+
C
H2N
+ 2H2NCH2COOH
O
Cu
H2N
O
CH2
C
O
50
Chemical Properties of Silver
Chemically silver is not very active. It is insoluble in dilute acids but
dissolves in concentrated nitric or sulfuric acid:
∆ Ag SO + SO ↑ + 2H O;
Ag + 2H SO (conc.) →
2
4
2
4
2
2
∆ AgNO + NO ↑ + H O;
Ag + 2HNO3(conc.) →
3
2
2
and it does not react with oxygen or water at ordinary temperatures.
–
Silver oxide formes as a dark brown precipitate when OH is added to a
+
solution of Ag salt:
2AgNO3 + 2KOH → Ag2O↓ + 2KNO3 + H2O.
It dissolves in a solution of ammonia hydroxide:
Ag2O + 4NH4OH + H2O → 2[Ag(NH3)2]OH + 4H2O.
Silver formes insoluble in water sulfide (black color), chloride (white
color), bromide (whity-yellow color) and iodide (yellow color):
AgNO3 + H2S → Ag2S↓ + 2HNO3;
AgNO3 + HCl → AgCl↓ + HNO3;
AgNO3 + HBr → AgBr↓ + HNO3;
AgNO3 + HI → AgI↓ + HNO3.
Chloride, bromide and iodide of silver dissolve in a solution of sodium
thiosulfate:
AgCl(s) + 2Na2S2O3 → Na3[Ag(S2O3)2](aq) + NaCl;
AgBr(s) + 2Na2S2O3 → Na3[Ag(S2O3)2](aq) + NaBr;
AgI(s) + 2Na2S2O3 → Na3[Ag(S2O3)2](aq) + NaI.
+
Ag -ion is oxidizing agent and react with formaldehyde (HCHO) or
glucose (C6H12O6) (reaction of silver mirror):
HCHO + 2[Ag(NH3)2]OH → HCOOH + 2Ag↓ + 4NH3↑ + H2O;
C6H12O6 + 2[Ag(NH3)2]OH → C6H12O7 + 2Ag↓ + 4NH3↑ + H2O.
Chemical Properties of Gold
Gold is extremely inactive and in compounds it has two common
oxidation states, +1 and +3.
It is unaffected by air, heat, moisture, and most solvents. However it
dissolves in aqueous solution containing chlorine and hydrochloric acid:
Au + 3Cl + HCl → H[AuCl4].
It dissolves in some oxidizing mixtures, such as cyanide ion with
oxygen:
4Au + O2 + 8NaCN + 2H2O → 4Na[Au(CN)2] + 4NaOH,
and in aqua regia (a mixture of hydrochloric and nitric acids):
Au + HNO3 + 4HCl → H[AuCl4] + NO↑ + 2H2O.
Gold (III) oxide and hydroxide are amphoteric compounds:
Au(OH)3 + NaOH → Na[Au(OH)4];
Au(OH)3 + 4HNO3 → H[Au(NO3)4] + 3H2O.
51
3+
Au -ion is strong oxidizing agent, for example:
3Na2[Sn(OH)4] + 2AuCl3 + 6H2O → 3Na[Sn(OH)6] + 2Au↓ +6HCl .
Uses
The main usage of copper are electrical, because of copper's extremely
high conductivity, which is surrendering to that of silver only. Copper can be
electroplated easily, alone or as a base for other metals. Large amounts
are used for this purpose, particularly in making electrotypes, reproductions
of type for printing.
Pure copper is soft but can be hardened somewhat by being worked.
Alloys of copper, which are far harder and stronger than the pure metal,
have higher resistance and therefore cannot be used for electrical
purposes. However they have corrosion resistance almost as good as that
of pure copper and are very easily worked in machine shops. The two most
important alloys are brass, a zinc alloy, and bronze, a tin alloy. Both tin and
zinc are sometimes added to the same alloy, and no sharp dividing line can
be drawn between brass and bronze. Both are used in enormous
quantities. Copper is alloyed also with gold, silver, and nickel, and is an
important constituent of such alloys as Monel metal, gunmetal, and
German silver.
Copper is an essential component of numerous oxidation-reduction
enzyme systems. For example, copper is a component of the enzymes
cytochrome oxidase, uricase, tyrosinase, superoxide dismutase, amine
oxidase, lysyl oxidase, and caeruloplasmin. As a component of the enzyme
caeruloplasmin (ferroxidase), copper is intimately involved with iron
metabolism, and therefore haemoglobin synthesis and red blood cell
production and maintenance. Copper is also believed to be necessary for
the formation of the pigment melanin and consequently skin pigmentation,
for the formation of bone and connective tissue, and for maintaining the
integrity of the myelin sheath of nerve fibres.
Silver is used to coat smooth glass surfaces for mirrors by vaporization
of the metal or by precipitation from a solution; however, aluminum has
replaced silver in this application largely. Silver is widely used also in the
circuitry of electrical and electronic components. Colloidal silver, dilute
solutions of silver nitrate (AgNO3), and some insoluble compounds, such as
potassium, are used in medicine as antiseptics and bactericides. Argyrol, a
silver-protein compound, is a local antiseptic for the eyes, ears, nose, and
throat.
The silver-halide salts – silver bromide, silver chloride, and silver iodide
– which darken when exposured to light, are used in emulsions for
photographic plates, film, and paper.
Gold is used in the form of gold sheet in the arts of gilding and lettering.
Chlorauric acid (H[AuCl4]) is used in photography for toning silver images.
52
Potassium dicyanourate(I) (K[Au(CN)2]) is used in electrogilding. Gold is
used also in dentistry. Radioisotopes of gold are used in biological research
and in the treatment of cancer.
8. Group IIB. Zinc, Cadmium and Mercury
Group IIB of the periodic table contains three chemical elements – zinc
Zn, cadmium Cd and mercury Hg.
Atoms of these elements have the following configurations:
2
30Zn
2
6
2
6
10
2
– 1s 2s 2p 3s 3p 3d 4s
2
2
6
2
6
10
2
6
10
2
48Cd – 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s
2
2
6
2
6
10
2
6
10 14
2
6
10
2
80Hg – 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s
10
2
[Ar]3d 2s
10
2
[Kr]4d 5s
10
2
[Xe]5d 6s
Table 8.1 lists some physical properties of group IIB chemical elements.
Table 8.1. Properties of Group IIB Elements
Cadmium
Zinc
Property
Electron configuration
10
[Ar] 3d 2s
2
10
[Kr] 4d 5s
2
Mercury
10
Melting point, °C
420
321
–38
Boiling point, °C
907
765
357
7.14
8.64
13.5
1.39
1.56
1.60
0.83
0.99
1.12
Density, g/cm
3
o
Covalent radius, A
+2
o
Ionic radius (Me ), A
2
[Xe] 5d 6s
Physical Properties and Occurrence of the Elements
Zinc, symbol Zn, is a bluish-white metallic element. The atomic number
of zinc is 30. Zinc melts at about 420 °C, boils at about 907 °C, and has a
specific gravity of 7.14. The atomic weight of zinc is 65.38.
The zinc ores have been known for a long time, but zinc was not
recognized as a separate element until 1746, when the German chemist
Andreas Sigismund Marggraf isolated the pure metal by heating calamine
and charcoal.
th
Zinc ranks 24 in abundance among the elements in the crust of the
earth. It never occurs free in nature, but is found as zinc oxide, ZnO, in the
mineral zincite; as zinc silicate, 2ZnO·SiO2·H2O, in the mineral
hemimorphite; as zinc carbonate, ZnCO3, in the mineral smithsonite; as a
mixed oxide of zinc and iron, Zn(FeO2)O2, in the mineral franklinite; and as
zinc sulfide, ZnS, in the mineral sphalerite, or zinc blende. The ores most
commonly used as a source of zinc are smithsonite and sphalerite.
53
Cadmium, symbol Cd, is a silvery-white metallic element. The atomic
number of cadmium is 48. Cadmium melts at 321 °C, boils at 765 °C, and
has a specific gravity of 8.64; the atomic weight of cadmium is 112.41.
Cadmium was discovered in 1817 by the German chemist Friedrich
Stromeyer, who found it in incrustations in zinc furnaces. The element
th
ranks about 65 in natural abundance among the elements in the earth's
crust. Cadmium occurs as the principal constituent of a mineral in the rare
greenockite only.
Mercury, (Latin, hydrargyrum, “liquid silver”), symbol Hg, metallic
element that is a free-flowing liquid at room temperature. The atomic
number of mercury is 80. Mercury melts at about –39 °C, boils at about 357
°C, and has a specific gravity of 13.5. The atomic weight of mercury is
200.59.
Mercury, once known as liquid silver and as quicksilver, was studied by
the alchemists. First it was distinguished as an element by the French
chemist Antoine Laurent Lavoisier in his experiment on the air composition.
th
Mercury ranks about 67 in natural abundance among the elements in
crustal rocks. It occurs in its pure form or combined with silver in small
amounts but is found most often in the form of the sulfide, the ore cinnabar
(HgS).
Preparation of Zinc
The first step in the metallurgy process of zinc metal preparation is
transforming the ores into oxides by subjecting them to high temperatures:
∆ 2ZnO + 2SO .
2ZnS + 3O →
2
2
Then the oxides are reduced by carbon in an electric furnace:
∆ Zn + CO↑,
ZnO + C →
where zinc is boiling and distilling in the retort in which the reduction takes
place. The zinc obtained by distillation contains small amounts of iron,
arsenic, cadmium, and lead and is known in metallurgy as spelter. In
another method of zinc refining, the roasted ores are leached with sulfuric
acid:
ZnO + H2SO4 → ZnSO4 + H2O.
After removing the impurities, the solution is electrolyzed:
2+
ZnSO4 ⇄ Zn
On cathode:
2+
–
Zn + 2e → Zn
2–
+ SO4
On anode:
–
+
2H2O – 4e → 4H + O2
is
2ZnSO4 + H2O electrolys


→ 2Zn + O2↑ + 2H2SO4.
Electrolytic zinc is pure and has superior qualities, such as high
resistance to corrosion.
Mercury is obtained easily from the ore by heating it in the air:
54
∆ Hg + SO ↑.
HgS + O2 →
2
Chemical Properties of Zinc
Zinc has one oxidation state only, the +2 state. Its valence-shell electron
10
2
2+
configuration is 3d 4s . It can lose the two 4s electrons to give the Zn ion,
10
which has a d arrangement of unused valence electrons.
Zinc metal reacts directly with chlorine, oxygen and sulfur at heating;
∆ ZnCl ;
∆ 2ZnO;
Zn + Cl →
2Zn + O →
2
2
2
∆ ZnS.
Zn + S →
It dissolves dilute acids (except nitric acid) forming hydrogen gas:
Zn + HCl → ZnCl + H2↑;
Zn + H2SO4 → ZnSO4 + H2↑.
Nitric and concentrated sulfuric acids reacts with zinc also:
∆ ZnSO + SO ↑ + 2H O;
Zn + 2H SO (conc.) →
2
4
4
2
2
∆ 4Zn(NO ) + NH NO + 3H O;
4Zn + 10HNO3(dil.) →
3 2
4
3
2
∆
3Zn + 8HNO (conc.) → 3Zn(NO ) + 2NO↑ + 4H O.
3
3 2
2
Zinc dissolves in concentrated solutions of strong bases:
Zn + 2KOH + 2H2O → K2[Zn(OH)4] + H2↑.
Addition of an alkali metal hydroxide or aqueous ammonia to a solution
of a zinc(ll) salt gives a white gelatinous precipitate of amphoteric zinc(ll)
hydroxide:
ZnSO4 + 2NaOH → Zn(OH)2↓ + Na2SO4.
The precipitate dissolves in excess hydroxide because of the formation of
2–
the [Zn(OH)4] ion:
Zn(OH)2 + 2KOH → K2[Zn(OH)4];
2+
and in the excess of ammonia because of the formation of the [Zn(NH3)4]
ion:
Zn(OH)2 + 4NH4OH → [Zn(NH3)4](OH)2 + 4H2O.
Zinc(ll) hydroxide dissolves in acids also:
Zn(OH)2 + 2HCl → ZnCl2 + 2H2O.
When hydrogen sulfide is passed into a basic solution of a zinc(ll) salt, a
white precipitate of zinc sulfide, ZnS, is obtained:
ZnCl2 + H2S → ZnS↓ + 2HCl.
Chemical Properties of Mercury
10
2
Mercury has the 5d 6s valence-shell electronic configuration. Only the
2+
s-electrons are used in compounds formation. Mercury forms a usual Hg
2+
ion as well as the unusual [Hg–Hg] ion in which two mercury atoms are
joined by a covalent bond.
55
Mercury forms an alloy with almost every metal except iron. These
mercury alloys are known as amalgams. The metal in these amalgams is
often less reactive than it is in the free state.
Mercury does not react with dilute acids, but it does react with
concentrated nitric acid and, when heated, with concentrated sulfuric acid
giving mercury(II) nitrate and mercury(II) sulfate, respectively:
∆ Hg(NO ) + 2NO ↑ + 2H O;
Hg + 4HNO (conc.) →
3
3 2
2
2
∆ HgSO + SO ↑ + 2H O.
Hg + 2H2SO4(conc.) →
4
2
2
Mercury(ll) chloride, HgCl2, can be made by heating mercury(ll) sulfate
with sodium chloride:
∆ Na SO + HgCl .
HgSO + 2NaCl →
4
2
4
2
The volatile HgCl2 distills off.
Mercury(ll) chloride is a covalent compound that even in aqueous
solution is present mainly as unionized covalent molecules. Since there are
only two bonding electron pairs in the valence shell of the mercury atom,
HgCl2 is a linear AX2 molecule.
Mercury combines with oxygen when heated in the air giving mercury(ll)
oxide:
∆ 2HgO.
2Hg + O →
2
However, when HgO is heated more strongly, it decomposes to mercury
and oxygen:
∆ 2Hg + O ↑ .
2HgO →
2
Mercury(ll) oxide can be obtained also by adding hydroxide ion to a
2+
solution of an Hg salt:
HgCl2 + 2KOH → HgO↓ + 2KCl + H2O.
When H2S is passed into a solution of a mercury(ll) salt, insoluble black
HgS is precipitated:
HgCl2 + H2S → HgS↓ + 2HCl.
2+
A very unusual property of mercury is that it forms the Hg2 ion in which
mercury atoms are joined by a covalent bond. Cadmium is the only other
metal known to form an analogous ion. This ion can be obtained, for
example, by the reaction between mercury and a solution of mercury(ll)
nitrate:
Hg + Hg(NO3)2 → 3Hg2(NO3)2.
2+
The Hg2 ion contains mercury in the +1 oxidation state and is called
the mercury(l) ion. When chloride ion is added to a solution of mercury(l)
nitrate, a white insoluble precipitate of mercury(l) chloride, Hg2Cl2, is
obtained:
Hg2(NO3)2 + 2HCl → Hg2Cl2↓ + 2HNO3.
56
Like mercury(II) chloride, Hg2Cl2 is a covalent molecule with a linear
structure in which both mercury atoms have the expected linear AX2
geometry.
Dimethyl mercury (CH3–Hg–CH3) is a volatile covalent compound. It is
an example of an organometallic compound, in which alkyl groups, or
substituted alkyl groups, are attached to a metal atom. Dimethyl mercury
can be prepared by the reaction of mercury-sodium amalgam with
chloromethane:
Hg + 2Na + 2CH3Cl → Hg(CH3)2 + 2NaCl .
It is a strong-smelling, very toxic, volatile substance (Tb= 96 °C). It has a
linear covalent structure. Reaction of (CH3)2Hg with HgCl2 gives
methylmercury chloride, CH3HgCl, another linear covalent molecule:
(CH3)2Hg + HgCl2 → 2CH3HgCl .
When the Cl is replaced by sulfate or nitrate, an ionic salt that contains the
+
covalent methylmercury cation, CH3–Hg , is obtained, for example,
+
–
CH3Hg NO3 .
Uses
Zinc is used mainly as a protective coating, or galvanizer, for iron and
steel; as an ingredient of various alloys, especially brass; as plates for dry
electric cells; and for die castings. Zinc oxide, known as zinc white or
Chinese white, is used as a paint pigment. It is used also as a filler in
rubber tires and in medicine as an antiseptic ointment. Zinc chloride is used
as a wood preservative and as a soldering fluid. Zinc sulfide is useful in
applications involving electroluminescence, photoconductivity, and
semiconductivity and has other electronic uses. It is used as a phosphor for
the screens of television tubes and in fluorescent coatings.
Zinc is an essential component of more than 80 metalloenzymes,
including carbonic anhydrase (required for the transport of carbon dioxide
by the blood and for the secretion of HCI in the stomach), glutamic
dehydrogenase, alkaline phosphatase, pyridine nucleotide dehydrogenase,
alcohol
dehydrogenase,
superoxide
dismutase,
pancreatic
carboxypeptidase, and tryptophan desmolase. It serves as a cofactor in
many enzyme systems, including arginase, enolase, several peptidases,
and oxalacetic decarboxylase. As an active component or cofactor for
many important enzyme systems zinc plays a vital role in lipid, protein, and
carbohydrate metabolism; being particularly active in the synthesis and
metabolism of nucleic acids (RNA) and proteins. Although not proven, it
has been suggested that zinc plays a role in the action of hormones such
as insulin, glucagon, corticotrophin, FSH and LH. Zinc is believed to play a
positive role in wound healing.
Cadmium can be deposited electrolytically as a coating on metals,
mainly iron or steel, on which it forms a chemically resistant coating.
57
Cadmium lowers the melting point of metals with which it is alloyed; it is
used with lead, tin, and bismuth in the manufacture of fusible metals for
automatic sprinkler systems, fire alarms, and electric fuses. An alloy of
cadmium with lead and zinc is used as a solder for iron. Cadmium salts are
used in photography and in the manufacture of fireworks, rubber,
fluorescent paints, glass, and porcelain. Cadmium has been used as a
control or shielding material in atomic energy plants because of its high
absorption of low-energy neutrons. Cadmium sulfide is used in a type of
photovoltaic cell, and nickel-cadmium batteries are in common use for
specialized purposes.
Cadmium and solutions of its compounds are highly toxic, with
cumulative effects similar to those of mercury poisoning.
Mercury is used in thermometers because its coefficient of expansion is
nearly constant; the change in volume for each degree of rise or fall in
temperature is the same. It is used also in other types of scientific
apparatus, such as vacuum pumps, barometers, and electric rectifiers and
switches. Mercury-vapor lamps are used as a source of ultraviolet rays in
homes and for sterilizing water. A mercury amalgam called dental amalgam
is used to fill teeth. It consists of about 50 % of mercury and 50 % of "dental
alloy" which is mainly silver and tin. Mercury(I) chloride (Hg2Cl2), or
calomel, is used for electrodes, and formerly used as a cathartic;
mercury(II) chloride (HgCl2), or corrosive sublimate; and medicinals such as
Mercurochrome.
Mercury Poisoning
Mercury is acutely hazardous as a vapor and in the form of its watersoluble salts, which corrode membranes of the body. Chronic mercury
poisoning, which occurs when small amounts of the metal or its fat-soluble
salts, particularly methylmercury, are repeatedly ingested over long periods
of time, causes irreversible brain, liver, and kidney damage. Because of
increasing water pollution, significant quantities of mercury have been
found in some species of fish, which has aroused concern regarding
uncontrolled discharge of the metal into the environment.
58
Contents
Inorganic chemistry .............................................................................................. 4
1. Hydrogen .......................................................................................................... 4
Physical Properties and Occurrence ................................................................. 5
Preparation Methods ........................................................................................ 5
Reactions of Hydrogen ..................................................................................... 6
Uses .................................................................................................................. 7
Water ................................................................................................................ 7
Physical Properties and Occurrence of Water .................................................. 8
Chemical Properties of Water .......................................................................... 8
Water Purification .......................................................................................... 10
Hydrogen Peroxide ......................................................................................... 10
Chemical Properties of Hydrogen Peroxide ................................................... 11
Uses ................................................................................................................ 12
Metals ................................................................................................................. 12
Physical Properties of Metals ......................................................................... 13
Chemical Properties of Metals ....................................................................... 13
Electron Structure of Metals........................................................................... 13
2. Group IA. Alkali Metals ................................................................................. 14
Physical Properties and Occurrence of the Alkali Metals .............................. 14
Preparation of the Elements ........................................................................... 16
Chemical Properties of the Alkali Metals....................................................... 16
Oxides and Hydroxides of the Alkali Metals ................................................. 17
Uses ................................................................................................................ 17
3. Group IIA. Beryllium, Magnesium and Alkaline Earth Metals ..................... 19
Physical Properties and Occurrence ............................................................... 19
Preparation of Elements ................................................................................. 20
Chemical Properties of Beryllium and its Compounds .................................. 21
Chemical Properties of Magnesium and its Compounds ................................ 23
Chemical Properties of the Alkaline Earth Metals ......................................... 23
Oxides and Hydroxides of the Alkaline Earth Metals .................................... 24
Uses and biological role ................................................................................. 24
4. Group VIB. Chromium, Molybdenum, Tungsten .......................................... 26
Physical Properties and Occurrence of the Elements ..................................... 27
Preparation of the Elements ........................................................................... 28
Chemical Properties of Chromium ................................................................. 28
Uses ................................................................................................................ 31
5. Group VIIB. Manganese, Technetium, Rhenium ........................................... 32
Physical Properties and Occurrence of the Elements ..................................... 33
Preparation of Manganese .............................................................................. 34
Chemical Properties of Manganese ................................................................ 34
Uses ................................................................................................................ 36
59
6. Group VIIIB. Iron, Cobalt, Nickel and the Platinum Metals .......................... 37
Physical Properties and Occurrence of Elements ........................................... 38
Preparation of the Elements ........................................................................... 40
Chemical Properties of Iron ........................................................................... 41
Chemical Properties of Cobalt and Nickel ..................................................... 43
Chemical Properties of the Platinum Metals .................................................. 44
Uses ................................................................................................................ 45
7. Group IB. Copper, Silver and Gold ................................................................ 47
Physical Properties and Occurrence of the Elements ..................................... 47
Preparation of the Elements ........................................................................... 48
Chemical Properties of Copper ...................................................................... 48
Chemical Properties of Silver......................................................................... 51
Chemical Properties of Gold .......................................................................... 51
Uses ................................................................................................................ 52
8. Group IIB. Zinc, Cadmium and Mercury ....................................................... 53
Physical Properties and Occurrence of the Elements ..................................... 53
Preparation of Zinc ......................................................................................... 54
Chemical Properties of Zinc ........................................................................... 55
Chemical Properties of Mercury .................................................................... 55
Uses ................................................................................................................ 57
60