B.Sc. Chemistry

Semester III

18UCE504 : Inorganic and Physical Chemistry

Unit I: Acids and Bases (12 Hours)

Definitions - different approaches to protonic acid - base systems - strengths of Lewis Acids

and Bases - Hard and Soft Acids and Bases. Applications of HSAB concept. Basis of

hardness and softness - limitations of HSAB concept.

Unit II: Metallurgy (12 Hours)

Metallurgy : Occurrence of metals – concentration of ores – froth floatation, magnetic

separation, calcination, roasting, smelting, flux, aluminothermic process, purification of

metals – electrolysis, zone refining, van Arkel de-Boer process.

Ozone and hydrogen peroxide – preparation, properties, structure, uses and comparison

between the two. Selenium and Tellurium – Extraction, properties and Uses. Oxides and

oxyacids of Se and Te. A comparative study of Sulphur, Selenium, Tellurium and their

compounds (hydrides, oxides, halides).

UNIT III: Equilibrium and Spontaneity (12 Hours)

General conditions of equilibrium and spontaneity- conditions of equilibrium and spontaneity

under constants – definition of A and G – physical significance of dA and dG.Temperature

and pressure dependence of G, Gibbs – Helmholtz equation and chemical equilibrium –

concept of chemical potential – chemical potential in a mixture of ideal gases – van’t Hoff

isotherm and isochore – Third law of thermodynamics – statement and applications -

exception to third law.

UNIT IV: Ionic Equilibria (12 Hours)

Solubility and solubility product - determination of solubility product - applications of

solubility product principle. Dissociation of weak acids and bases - dissociation constants -

pH scale - common ion effect - buffer solutions - determination of pH values of buffer

mixtures - Henderson’s equation - Hydrolysis of salts - degree of hydrolysis.

UNIT V: Chromatography (12 Hours)

Chromatographic methods – partition adsorption – basic principles – differential migration,

adsorption phenomenon, nature of adsorbents, choice of solvents and Rf value – techniques

and applications of paper, column, TLC and gas chromatography. Basic principles of HPLC.

Unit II: Metallurgy

Metallurgy : Occurrence of metals

The crust of the earth provides metals and it is a good source to procure metals. Mostly,

the metals occur in nature in a combined state but sometimes they can also occur in the free

state. A native metal is a metal found in its metallic state naturally, either in pure form or in the

form of an alloy. Most metals can’t resist natural processes like oxidation, corrosion, etc. Hence,

only non-reactive metals like gold, silver, platinum, etc are found in the native or free state.

Most metals are obtained in the form of compounds and they need to be filtered from their

impurities to be further used for various applications. The process of procuring metals from ores

is called metallurgy and these naturally occurring compounds of metals are known as minerals.

Metals are usually extracted from the earth by mining.

Ores

All ores are minerals, but not all minerals are ores. Ores are those minerals which

contain metals. The picture shown above is a gold-bearing ore. Different methods are used to

extract different metals. For instance, common metals like iron are smelted using carbon as a

reducing agent. Few metals, such as aluminum and sodium, don’t have a commercially practical

reducing agent and are extracted using electrolysis instead. Sulfide ores are not reduced directly

to the metal but are roasted in air to convert them to oxides.

Gangue or Matrix

The gangue or matrix is an unwanted, commercially useless, rocky, earthy or sandy

materials that are found with the ores. These are the impurities are filtered out at a later stage.

Extraction of Metals

The occurrence of metals in nature is only a part of the process, these metals also have to

be extracted. Metallurgy is the process of extracting metals. The basic step of extracting metals

is called concentration of ores.

Concentration of Ores

The ores contain a lot of impurities that need to be filtered. The filtration of these

impurities from ores is known as concentration. Concentration can be done by gravity separation

or hydraulic washings, or froth flotation process.

Metals of high reactivity level

Metals in the top of the reactivity level i.e. metals like sodium, magnesium, calcium, etc

cannot be obtained from their compound by heating with carbon as they are very reactive.

Electrolytic reduction is used to obtain such metals.

Metals of medium reactivity level

Metals in the middle of the reactivity series i.e metals like lead, zinc, copper, etc are

moderately reactive and are generally present as carbonates or sulphides. Sulphides ores are

converted into oxides by heating strongly in the presence of excess air. Sulphur impurities can

escape as gas, this process is known as roasting. Carbonates ores are calcinated to eliminate

carbonate and moisture impurities. Calcination is a process where the ore is converted into

oxides by heating strongly in the presence of excess air.

Metals of low reactivity level

Metals that aren’t very reactive are placed low in the reactivity series. These metals can

be reduced to metals by heating alone. For example, Mercury is procured from its ore, cinnabar

(HgS), by the process of heating. Copper can also be obtained from its sulphide ore (Cu2S) by

heating.

Froth floatation;

Principle: The principle of froth floatation is that sulphide ores are preferentially wetted by

pine oil, whereas the gangue particles are wetted by water. Collectors are added to enhance

the non-wettability of the mineral particles. Examples of collectors are pine oil, fatty acids

and xanthates. Froth stabilisers are added to stabilise the froth. Examples of froth stabilisers

are cresols, aniline. If two sulphide ores are present, then it is possible to separate the two

sulphide ores by adjusting the proportion of oil to water or by adding depressants. Example:

For an ore containing ZnS and PbS, the depressant used is NaCN. It selectively prevents ZnS

from coming to froth but allows PbS to come with the froth.

Method:

This method employs a mixture of water and pine oil which is made to froth in a tank

to separate sulphide ores. The differences in the wetting properties of the ore and gangue

particles separate them.

A mixture of water, pine oil, detergent and powdered ore is first taken in a tank. A

blast of compressed air is blown through the pipe of a rotating agitator to produce froth. The

sulphide ore particles are wetted and coated by pine oil and rise up along with the froth (froth

being lighter). The gangue particles wetted by water sink to the bottom of the tank (water

being heavier). Sulphide being more electronegative attracts the covalent oil molecules. The

gangue being less electronegative is attracted by the water. The froth containing the sulphide

ore is transferred to another container, washed, and dried.

Magnetic separation:

This method is used in those cases where either ore or the impurities are of magnetic

in nature. In this method, the powdered impure ore in the form of thin layer is allowed to fall

on a rubber belt which moves horizontally over two rollers, one of which has electromagnetic

attached it. As the ore particles roll over the belt, the magnetic component in the ore gets

attracted towards the magnet. It gets collected in a heap while the non – magnetic component

forms a separate heap.

Chromite Fe(CrO2)2 an ore of chromium is magnetic in nature and contains non –

magnetic impurities casseterite (SnO2) is non – magnetic in nature while the impurities of

iron and chromium are magnetic in nature. These can be separated by magnetic separation

Calcination:

What is Calcination?

As per the popular definition, calcination is defined as the process of converting an

ore into an oxide by heating it strongly. The ore is heated below its melting point either in

absence of air or in a limited supply. This method is commonly used for converting

carbonates and hydroxides to their respective oxides. During calcination, moisture and

volatile impurities are also removed. Calcination can also be described as a thermal process

that is used to convert ores and other solid materials by bringing about thermal

decomposition. In calcination, reaction most of the time occurs at or above the thermal

decomposition temperature.

Calcination is derived from a Latin word calcinare which translates as “to burn lime”.

So, calcination is mostly used in the decomposition of limestone (calcium carbonate) to lime

(calcium oxide) and carbon dioxide.

CaCo3 —> CaO + CO2

Meanwhile, the products that are derived from calcination are know as calcine, and it is

regardless of actual compounds undergoing thermal treatment.

What is Roasting?

When we talk about roasting, it is basically a process of metallurgy where an ore is

converted into its oxide by heating it above its melting point in the presence of excess air.

While calcination is mostly used in the oxidation of carbonates, roasting is a method that is

used for converting sulphide ores. During roasting, moisture and non metallic impurities in

the form of volatile gases are released. Process of roasting consist of solid-gas thermal

reaction which includes oxidation, reduction, sulfation, chlorination and pyro hydrolysis.

However, roasting which involves sulphides are a major source of air pollution and the main

drawback of this process is that it releases large amount of metallic as well as toxic and acidic

compounds which causes harm to the environment.

An example of roasting is when Zinc sulphide is converted into zinc oxide.

2ZnS+3O2 —> 2ZnO + CO2

Important Differences Between Calcination and Roasting

Here are some of the major differences between calcination and roasting.

Calcination Roasting

Calcination is a process in which an ore is heated

in the absence of air or air might be supplied in

limited quantity

Roasting involves heating of ore higher

than its melting point in the presence of

air or oxygen.

Calcination involves thermal decomposition of

carbonate ores.

Roasting is carried out mostly for sulfide

minerals.

During calcination moisture is driven out from an

ore.

Roasting does not involve dehydrating an

ore.

Carbon dioxide is given out during calcination

During roasting large amount of toxic,

metallic and acidic compounds are

released.

Smelting:

Smelting, process by which a metal is obtained, either as the element or as a

simple compound, from its ore by heating beyond the melting point, ordinarily in the

presence of oxidizing agents, such as air, or reducing agents, such as coke. The first metal to

be smelted in the ancient Middle East was probably copper (by 5000 BCE), followed

by tin, lead, and silver. To achieve the high temperatures required for smelting, furnaces with

forced-air draft were developed; for iron, temperatures even higher were required. Smelting

thus represented a major technological achievement. Charcoal was the universal fuel

until coke was introduced in 18th-century England. Meanwhile, the blast furnace had

achieved a high state of development.

In modern ore treatment, various preliminary steps are usually carried out before

smelting in order to concentrate the metal ore as much as possible. In the smelting process a

metal that is combined with oxygen—for example, iron oxide—is heated to a high

temperature, and the oxide is caused to combine with the carbon in the fuel, escaping

as carbon monoxide or carbon dioxide. Other impurities, collectively called gangue, are

removed by adding a flux with which they combine to form a slag.

flux:

Flux, in metallurgy, any substance introduced in the smelting of ores to promote

fluidity and to remove objectionable impurities in the form of slag. Limestone is commonly

used for this purpose in smelting iron ores. Other materials used as fluxes are silica, dolomite,

lime, borax, and fluorite. In soldering, a flux is used to remove oxide films, promote wetting,

and prevent reoxidation of the surfaces during heating. Rosin is widely used as a

noncorrosive flux in soldering electronic equipment; for other purposes, a water solution

of zinc chloride and ammonium chloride may be used.

Aluminothermic process;

Aluminothermic reactions are exothermic chemical reactions using aluminium as

the reducing agent at high temperature. The process is industrially useful for production

of alloys of iron. The most prominent example is the thermite reaction between iron

oxides and aluminium to produce iron itself:

Fe2O3 + 2 Al → 2 Fe + Al2O3

This specific reaction is however not relevant to the most important application of

aluminothermic reactions, the production of ferroalloys. For the production of iron, a cheaper

reducing agent, coke, is used instead via the carbothermic reaction.

Aluminothermic reactions are exothermic chemical reactions using aluminium as

the reducing agent at high temperature. The process is industrially useful for production

of alloys of iron. The most prominent example is the thermite reaction between iron

oxides and aluminium to produce iron itself:

FeO + 2 Al → 2 Fe + AlO

This specific reaction is however not relevant to the most important application of

aluminothermic reactions, the production of ferroalloys. For the production of iron, a cheaper

reducing agent, coke, is used instead via the carbothermic reaction.

History

Aluminothermy started from the experiments of Russian scientist Nikolay Beketov at

the University of Kharkiv in Ukraine, who proved that aluminium restored metals from

their oxides under high temperatures. The reaction was first used for the carbon-free

reduction of metal oxides. The reaction is highly exothermic, but it has a high activation

energy since strong interatomic bonds in the solids must be broken first. The oxide was

heated with aluminium in a crucible in a furnace. The runaway reaction made it possible to

produce only small quantities of material. Hans Goldschmidt improved the aluminothermic

process between 1893 and 1898, by igniting the mixture of fine metal oxide and aluminium

powder by a starter reaction without heating the mixture externally. The process was patented

in 1898 and used extensively in the later years for rail trackwelding.

Applications

The aluminothermic reaction is used for the production of several ferroalloys, for

example ferroniobium from niobium pentoxide and ferrovanadium from iron, vanadium(V)

oxide, and aluminium. The process begins with the reduction of the oxide by the aluminium:

3 VO + 10 Al → 5 AlO + 6 V

Other metals can be produced from their oxides in the same way.

Aluminothermic reactions have been used to welding rail tracks on-site, useful for

complex installations or local repairs that cannot be done using continuously welded rail.

Another common use is the welding of copper cables (wire) for use in direct burial

(grounding/earthing) applications. It is still the only type of electrical connection recognized

by the IEEE (IEEE, Std 80-2001) as continuous un-spliced cable.

purification of metals – electrolysis

Electrolysis uses currents to give the energy needed to help a chemical reaction occur

to break apart chemical bonds between metals and impurities. This method was first used in

the late 18th century to separate tin and zinc from their salts. This method is also used for

purifying aluminum. Aluminum oxide is melted and a current is passed through it. The

electric current separates the aluminum from the oxygen. A cathode adds electrons on the

aluminum, making pure aluminum, and an anode collects the extra electrons on the oxygen,

combining it with carbon to form CO.

It is the process of refining of metal in which impure metal is made the anode and a

thin strip of pure metal is made the cathode. A solution of metal salt is used as an electrolyte.

The apparatus is shown in the figure given below. When electric current is passed through the

electrolyte, the impure metal from the anode is dissolved in the electrolyte and an equal

amount of pure metal from the electrolyte is deposited on the cathode. The soluble impurities

go into the solution whereas the insoluble impurities settle down at the bottom of the anode

which is known as anode mud. Following diagram shows the electrolytic refining of copper

in which anode is impure copper and cathode is thin strip of pure copper while the electrolyte

used is the copper sulphate solution.

Zone refining:

Zone refining is the most important of the zone-melting techniques. In zone refining,

a solid is refined by passing a number of molten zones through it in one direction. Each zone

carries a fraction of the impurities to the end of the solid charge, thereby purifying the

remainder. Zone refining was first described by the U.S. scientist W.G. Pfann and was first

used in the early 1950s to purify germanium for transistors. The purity achieved was hitherto

unheard of—less than one part of detectable impurity in 10,000,000,000 parts of germanium.

The method was adopted in transistor manufacture around the world.

The principles of zone refining are quite general, and so the method has been applied to many

substances. More than one-third of the elements and hundreds of inorganic and

organic compounds have been raised to their highest purity by zone refining. Many of these

were, for the first time, made pure enough for their intrinsic properties to be determined

zone remelting, in which two solutes are distributed through a pure metal. This is

important in the manufacture of semiconductors, where two solutes of opposite conductivity

type are used. For example, in germanium, pentavalent elements of group Vsuch

as antimony and arsenic produce negative (n-type) conduction and the trivalent elements

of group III such as aluminum and boronproduce positive (p-type) conduction. By melting a

portion of such an ingot and slowly refreezing it, solutes in the molten region become

distributed to form the desired n-p and p-n junctions.

Van Arkel de-Boer process:

Ultra pure metals are being prepared by the Van Arkel Method. Crude metal is heated

with a suitable substance so that the pure metal present in it may be converted into stable

volatile compound leaving behind impurities. The compound so formed is then decomposed

by heating to get the pure metal. Van Arkels method is used to purify crude titanium metal. It

is heated with iodine to about 500K to form volatile compound. Til4 leaving behind the

impurities .Til4 is further heated to 1700K when it decomposes to give pure titanium.

Ozone – preparation, properties, structure, uses;

Ozone is an allotrope of oxygen. It is highly unstable in nature. We can find its traces

for about 20 kilometres above the sea level. So, how is ozone present at such heights? It is

formed by the reaction of oxygen with the ultraviolet rays of the sun. The major role of this

gas is protecting the earth’s surface from the harmful ultraviolet radiations from the sun.

You have heard of the depletion of the ozone layer. You know why is that? We use

chlorofluorocarbons in refrigerators and other aerosols also that release harmful things into

the air leading to gaps in the layer. Due to this, we get the UV light and it causes a lot of skin

problems and malignancy in people.

Apart from this, the various oxides of nitrogen, particularly nitric oxide, react very

rapidly with ozone to produce oxygen and nitrogen dioxide. Hence, the various nitrogen

oxides of nitrogen that appear from the fumes frameworks of supersonic fly planes lead to

depleting the layer.

Preparation

We can prepare ozone by passing a silent electric discharge through dry,

unadulterated, and cold oxygen in an extraordinary device. This device is what we know as

the ozoniser. In this process, we obtain the gas of up to 10% concentration. This is an

endothermic process and we must take care to complete it at high temperatures. That is why

we use a silent electric discharge.

If we want to produce higher concentrations of ozone, we can do so by the process of

fractional liquefaction of an oxygen and ozone mixture.

In principle, this ozoniser is like the Siemen's ozoniser but dilute sulphuric acid

replaces the tin foil. Two carbon electrodes are dipped in the acid and connected to an

induction coil. A current of dry oxygen is passed through the space between the tubes.

Ozonised oxygen containing about 5% O3 comes out at the other end. If the apparatus is kept

cool, the proportion of ozone may go up 20-25%.

Structure of ozone:

The occurrence of ozone is in small amounts, in the upper layer of the atmosphere,

where it is formed due to the action of ultraviolet rays on the oxygen of the air. It is also

present in seawater where it is formed due to the reaction of fluorine with water.

In the structure of ozone, the bond length of 127.8 pm is intermediate between a single bond

(bond length 148 pm) and a double bond (bond length 110 pm). Ozone is, therefore,

considered to be a resonance hybrid of the following canonical forms:

Physical Properties

Ozone is a gas with a light blueish colour.

It has a fishy smell.

It condenses at – 120°C to give a dull blue fluid. On further cooling, it hardens to

give dark violet crystals.

Thermodynamically, it is very unstable and disintegrates to oxygen. This is an

exothermic process and is catalyzed by numerous materials. However, we must know

that high concentration of the gas can be very dangerous.

Oxidizing Action

Ozone is taken as a very strong oxidizing agent. This is mainly because of the ease

with which it gives out atoms of nascent oxygen. It helps in oxidising:

o Lead sulphide to lead sulphate:

4O3 + PbS → 4O2 + PbSO4

o Iodide ions to iodine:

2KI + H2O + O3 → 2KOH + I2 + O2

We use this particular reaction in the process of quantitative estimation of the gas.

When we bring ozone in contact with potassium iodide solution with a borate buffer

(pH 9.2), it liberates iodine. We can titrate this liberated iodine against a standard

solution of sodium thiosulphate using starch as an indicator. The reactions that

involve in this process are:

It also helps in oxidising nitrogen dioxide to dinitrogen pentoxide.

2NO2 + O3 → N2O5 + O2

Uses of ozone

For air purification at the crowded places like cinema halls and tunnel railways. Due

to its strong oxidizing power it also destroys the foul smell in slaughter houses.

In sterilizing drinking water by oxidizing all germs and bacteria.

For preservation of meat in cold storage.

For bleaching delicate fabrics such as silk, ivory, oils, starch and wax.

It helps to locate a double bond in any unsaturated organic compound by ozonolysis.

Hydrogen peroxide-preparation, properties, structure, uses and comparison between

the two

Hydrogen peroxide is the simplest kind of peroxide available (oxygen-oxygen single

bond). It is a colorless liquidand is used in aqueous solution for safety reasons. It acts as

a bleaching agent and is also used as a disinfectant. Concentrated hydrogen peroxide is a very

reactive oxygen species and is used as a propellant in rocketry. The chemical formula for

hydrogen peroxide is H2O2.

Laboratory Methods of Preparation

When barium peroxide is acidified and the excess water is removed by the process of

evaporation under reduced pressure, we obtain hydrogen peroxide. The following reaction

will clarify this:

BaO2.8H2O(s) + H2SO4(aq) → BaSO4(s) + H2O2(aq) + 8H2O(l)

Industrial Method of Preparation

Hydrogen peroxide is prepared by the electrolysis of 30% ice-cold H2SO4. When

acidified sulfate solution is electrolyzed at high current density, peroxodisulphate is obtained.

Peroxodisulphate is then hydrolyzed to get hydrogen peroxide.

2HSO–

4(aq)[Electrolysis]→HO3SOOSO3H(aq)[Hydrolysis]→2HSO–4(aq)+2H+(aq) +H2O2(aq)

Properties of Hydrogen Peroxide

Physical Properties:

In the pure state, hydrogen peroxide is almost colorless (very pale blue) liquid.

It melts at 272.4 K and has a boiling point of 423 K (extrapolated).

It is miscible in water in all proportions and forms hydrates.

Chemical Properties:

Hydrogen peroxide in both acidic and basic medium acts as an oxidizing as well as

the reducing agent. The following reactions will give a clear picture:

Oxidation Reactions of H2O2

Oxidizes black Pbs to white PbSO4. Reaction: Pbs + 4H2O2 → PbSO4 + 4H2O

Oxidises KI to Iodine. Reaction: 2KI + H2O2 → 2KOH + I2

Oxidizes nitrites to nitrates. Reaction: NaNO2 + H2O2 → NaNO3 + H2O

Oxidizes acidified Potassium ferrocyanide. Reaction: 2K4Fe(CN)6 + H2SO4 +

H2O2 → 2K3Fe(CN)6 + K2SO4 + 2H2O

Oxidizes sulphites to sulphates. Reaction: Na2SO3 + H2O2 → Na2SO4 + H2O

Oxidizes CrO-2

4 in an acid medium to chromium Peroxide CrO5. Reaction: CrO-2

4 +

2H+ + 2H2O2 → CrO5 + 3H2O

Structure of Hydrogen Peroxide

The structure of hydrogen peroxide is non-planar. H2O2 has an open book structure

with O – O spins. The dihedral angle is 111°. The O-O bond length is 145.8 pm and

the O-H bond length is 98.8 pm(which is equal to 9.88 × 10-13

m). The following

diagram will clearly show what an open book structure means.

Hydrogen Peroxide Structure

There are two planes in this structure and each plane has one O-H bond pair, the angle

between both the planes is 90.2°.

Uses of Hydrogen Peroxide

Since it acts as both an oxidizing as well as a reducing agent, it has a vast area of

applications:

Hydrogen peroxide is used as an antiseptic.

It is used in the preservation of wine and milk.

Our white blood cells produce hydrogen peroxide to kill bacteria.

It is used to clean stains from sensitive fabrics and clothes.

Soaking feet in water mixed with H2O2 for 15 min every day for a week will help to

get rid of fungus

Plant growth can be sustained by sprinkling mixture of hydrogen peroxide and water

in equal amounts

The textile and paper industry use it as a bleaching agent.

It is used for the synthesis of tartaric acid, food products, and many pharmaceuticals.

It is used in the manufacture of chemicals which in turn are used in making of high-

quality detergents.

The most significant use of H2O2 is in environmental chemistry where it is used in

pollution control treatment of domestic waste and industrial effluents.

Highly concentrated Hydrogen Peroxide is used as rocket propellant.

Selenium-Extraction, properties and Uses

Occurrence in nature

Selenium is a very rare element. Scientists estimate its abundance at about 0.05 to

0.09 parts per million. It ranks among the 25 least common elements in the Earth's crust. It is

widely distributed throughout the crust. There is no ore from which it can be mined with

profit. Instead, it is obtained as a by-product of mining other metals. It is now produced

primarily from copper, iron , and lead ores. The major producers of selenium in the world

are Japan, Canada, Belgium, the United States, and Germany.

Extraction

Selenium is a very rare element, and so scientists predict its abundance at 0.05 to 0.09

parts per million. Selenium is one of the 25 least common elements in Earth's Crust and it is

widely distributed throughout the crust. There is no ore in which in can be mined, so instead

it is obtained as a by-product of mining other metals. Now, Selenium is primarily produced

from iron,lead and copper ores. Pure Selenium is obtained from the slimes and sludges

formed in producing sulfuric acid. Within the sulfuric acid, the impure red Selenium is

dissolved. Selenious acid ( H2SeO3) and selenic acid ( H2SeO4) are formed and can be

drained from residual and insoluble material.Oxidation by air is used in other method.

Heating with sodium carbonate gives soluble sodium selenite and sodium selenate. Out of all

processes used to prepare/extract selenium, and in each process selenium compounds are

converted by water to become selenious acid. Selenium is then finally recoverd by treating

the selenious acid with sulfur dioxide.

Selenium is obtained as a by-product from other industrial processes. For example,

when copper is refined, small amounts of selenium are produced as by-products. This

selenium can be removed from the copper-refining process and purified. Selenium is also

obtained as a secondary product during the manufacture of sulfuric acid

Physical properties

Selenium exists in a number of allotropic forms. Allotropes are forms of an element

with different physical and chemical properties. One allotrope of selenium is an amorphous

red powder. Amorphous means "without crystalline shape." A lump of clay is an example of

an amorphous material. A second allotrope of selenium has a bluish, metallic appearance. A

number of other allotropes have properties somewhere between these two forms.

The amorphous forms of selenium do not have specific melting points. Instead, they

gradually become softer as they are heated. They may also change from one color and texture

to another.

The crystalline (metallic) form of selenium has a melting point of217°C (423°F ) and a

boiling point of 685°C (1,260°F). Its density is 4.5 grams per cubic centimeter.

Selenium from the Greek word for moon, selene .

Some of the most important physical characteristics of selenium are its electrical

properties. For example, selenium is a semiconductor. A semiconductor is a substance that

conducts an electric current better than non-conductors, but not as well as conductors.

Semiconductors have many very important applications today in the electronics industry.

Selenium is often used in the manufacture of transistors for computers, cellular phones, and

hand-held electronic games.

Selenium is also a photoconductor, a material that changes light energy into electrical

energy. Furthermore, it becomes better at making this conversion as the light intensity or

brightness increases.

Chemical properties

Selenium is a fairly reactive element. It combines easily with hydrogen, fluorine,

chlorine, and bromine. It reacts with nitric and sulfuric acids. It also combines with a

number of metals to form compounds called selenides. An example is magnesium selenide

(MgSe). One of its interesting reactions is with oxygen. It burns in oxygen with a bright blue

flame to form selenium dioxide (SeO 2 ). Selenium dioxide has a characteristic odor of rotten

horseradish.

Selenium and tellurium are often associated with each other. They tend to occur together in

the Earth and have somewhat similar properties.

Uses

The two most important uses of selenium are in glass-making and in electronics. Each

accounts for about 30 to 35 percent of all the selenium produced each year. The addition of

selenium to glass can have one of two opposite effects. First, it will cancel out the green color

that iron compounds usually add to glass. If a colorless glass is desired, a little selenium is

added to neutralize the effects of iron. Second, selenium will add its own color—a beautiful

ruby red—if that is wanted in a glass product.

Selenium is also added to glass used in architecture. The selenium reduces the amount

of sunlight that gets through the glass.

A growing use of selenium is in electronic products. One of the most important uses is

in plain-paper photocopiers and laser printers. The element is also used to make photovoltaic

("solar") cells. When light strikes selenium, it is changed into electricity. A solar cell is a

device for capturing the energy of sunlight on tiny pieces of selenium. The sunlight is then

changed into electrical energy.

Currently, that process is not very efficient. Too much sunlight is lost without being

converted into electricity. More efficient solar cells will be able to make use of all the free

sunlight that strikes the planet every day.

About a third of all selenium produced is used as pigments (coloring agents) for

paints, plastics, ceramics, and glazes. Depending on the form of selenium used, the color

ranges from deep red to light orange.

Selenium is also used to make alloys. An alloy is made by melting and mixing two or more

metals. The mixture has properties different from those of the individual metals. The addition

of selenium to a metal makes it more machinable. Machinability means working with a

metal: bending, cutting, shaping, turning, and finishing the metal, for example.

About 5 percent of all selenium produced is used in agriculture. It is added to soil or

animal feed to provide the low levels of selenium needed by plants and animals.

Tellurium -Extraction, properties and Uses

Occurrence in nature

Tellurium is one of the rarest elements in the Earth's crust. Its abundance is estimated

to be about 1 part per billion. That places it about number 75 in abundance of the elements in

the earth. It is less common than gold, silver, or platinum.

The most common mineral of tellurium is sylvanite. Sylvanite is a complex

combination of gold, silver, and tellurium. Tellurium is obtained commercially today as a by-

product in copper and lead refining.

Extraction

A common method for obtaining tellurium is to pass an electric current through

dissolved tellurium dioxide (TeO 2 ). The current breaks the tellurium dioxide down into

oxygen and tellurium:

Tellurium has the unusual property of combining with gold. Gold normally combines with

very few elements.

Physical properties

Tellurium is a grayish-white solid with a shiny surface. It has a melting point of

449.8°C (841.6°F) and a boiling point of 989.9°C (1,814°F). Its density is 6.24 grams per

cubic centimeter. It is relatively soft. Although it has many metal-like properties, it breaks

apart rather easily and does not conduct an electric current very well.

Chemical properties

Tellurium does not dissolve in water. But it does dissolve in most acids and some

alkalis. An alkali is a chemical with properties opposite those of an acid. Sodium

hydroxide (common lye, such as Drano) and limewater are examples of alkalis.

Tellurium also has the unusual property of combining with gold. Gold normally

combines with very few elements. The compound formed between gold and tellurium is

called gold telluride (Au 2 Te 3 ). Much of the gold found in the earth occurs in the form of

gold telluride.

Uses and compounds

About 75 percent of all the tellurium produced today is used in alloys. Its most

important alloy is a tellurium-steel alloy. It has better machinability than does steel without

tellurium. Machinability means working with a metal: bending, cutting, shaping, turning, and

finishing the metal, for example. Adding 0.04 percent tellurium to steel makes it much easier

to work with.

Tellurium is also added to copper to improve machinability. Tellurium-copper alloys

are also easier to work with than pure copper. And the essential ability of copper to conduct

an electric current is not affected. Tellurium is also added to lead. Tellurium-lead alloys are

more resistant to vibration and fatigue than pure lead. Metal fatigue is the tendency of a metal

to wear out and eventually break down after long use.

About 15 percent of all tellurium produced is used in the rubber and textile industries.

It is important in the vulcanization of rubber, for example. Vulcanization is the process by

which soft rubber is converted to a harder, longer-lasting product. Tellurium is also used as a

catalyst in the manufacture of synthetic fibers. A catalyst is a substance used to speed up or

slow down a chemical reaction without undergoing any change itself.

A growing application of tellurium is in a variety of electrical devices. For example, it

is used to improve picture quality in photocopiers and printers. A compound of

tellurium, cadmium, and mercury is also used in infrared detection systems. Infrared

radiation is heat. It can be made visible with special glass. Some satellites orbiting the Earth

study forests, crops, and other plant life by measuring the infrared radiation they give off.

Finally, a very small amount of tellurium is used for minor applications, such as a coloring

agent in glass and ceramics and in blasting caps for construction projects.

Tellurium gives a garlicky-odor to the breath.

Oxides of Se

Selenium Oxide (or Selenium Dioxide) is available in many powder and crystal sizes.

Ultra high purity and high purity forms include powder, submicron powder and nanoscale.

Selenium Oxide is generally immediately available in most volumes. Oxide compounds are

not conductive to electricity. However, certain perovskite structured oxides are electronically

conductive finding application in the cathode of solid oxide fuel cells and oxygen generation

systems. They are compounds containing at least one oxygen anion and one metallic cation.

They are typically insoluble in aqueous solutions (water) and extremely stable making them

useful in ceramic structures as simple as producing clay bowls to advanced electronics and in

light weight structural components in aerospace and electrochemical applications such as fuel

cells in which they exhibit ionic conductivity. Metal oxide compounds are basic anhydrides

and can therefore react with acids and with strong reducing agents in redoxreactions.

Selenium Oxide is also available in pellets, pieces, powder, sputtering targets, tablets,

and nanopowder (from American Elements' nanoscale productionfacilities). Metallic oxide

compounds are also available from American Elements' nanoscale production facilities.

American Elements produces to many standard grades when applicable, including Mil Spec

(military grade); ACS, Reagent and Technical Grade; Food, Agricultural and Pharmaceutical

Grade; Optical Grade, USP and EP/BP (European Pharmacopoeia/British Pharmacopoeia)

and follows applicable ASTM testing standards. Typical and custom packaging is available.

Additional technical, research and safety (MSDS) information is available as is a Reference

Calculator for converting relevant units of measurement.

Effects of Exposure:

To the best of our knowledge the chemical, physical and toxicological properties of

selenium oxide have not been thoroughly investigated and recorded.

Selenium in small amounts is essential for normal growth of some animals.

Deficiency or excess is associated with serious disease in livestock. Elemental selenium has

low acute systemic toxicity, but dust or fumes can cause serious irritation of the respiratory

tract. It is a suspected carcinogen of the liver and thyroid. Inorganic selenium compounds

can cause dermatitis. Garlic odor of breath is a common symptom. Pallor, nervousness,

depression, digestive disturbances have been reported in cases of chronic exposure. Selenium

compounds are common air contaminants. (Sax, Dangerous Properties of Industrial

Materials)

Selenium may cause amyotrophic lateral sclerosis, bronchial irritation, gastrointestinal

distress, nasopharyngeal irritation, garlic odor on breath and sweat, metallic taste, pallor,

irritability, excessive fatigue, loss of fingernails and hair, pulmonary edema, anemia and

weight loss.

Acute and Chronic Effects:

Inhalation: Inhalation of selenium dust or fume can cause irritation, metallic taste and a

garlic-odor to the breath.

Ingestion: May cause pallor, weakness, G.I. tract disturbances and nervousness.

Skin: Irritant. Repeated skin contact can cause dermatitis.

Eye: Irritant. SeO2 in the eye can cause a pink allergic-type reaction.

Oxides of Tellurium

Tellurium dioxide (TeO2) is a solid oxide of tellurium. It is encountered in two

different forms, the yellow orthorhombic mineral tellurite, β-TeO2, and the synthetic,

colourless tetragonal (paratellurite), α-TeO2. Most of the information regarding reaction

chemistry has been obtained in studies involving paratellurite, α-TeO2.

Preparation

Paratellurite, α-TeO2, is produced by reacting tellurium with O2:

Te + O2 → TeO2

An alternative preparation is to dehydrate tellurous acid, H2TeO3, or to thermally

decompose basic tellurium nitrate,Te2O4.HNO3 above 400 °C.

Physical properties

The speed of sound in Tellurium dioxide is 4,250 metres per second (13,900 ft/s).

Chemical properties

TeO2 is barely soluble in water and soluble in strong acids and alkali metal hydroxides.[4]

It is

an amphoteric substance and therefore can act both as an acid or as a base depending on the

solution it is in.[5]

It reacts with acids to make tellurium salts and bases to make tellurites. It

can be oxidized to telluric acid or tellurates.

Structure

Paratellurite, α-TeO2, converts at high pressure into the β-, tellurite form Both the α-,

(paratellurite) and β- (tellurite forms) contain four coordinate Te with the oxygen atoms at

four of the corners of a trigonal bipyramid. In paratellurite all vertices are shared to give

a rutile-like structure, where the O-Te-O bond angle are 140°. α-TeO2 In tellurite pairs of

trigonal pyramidal, TeO4 units, sharing an edge, share vertices to then form a layer. The

shortest Te-Te distance in tellurite is 317 pm, compared to 374 pm in paratellurite. Similar

Te2O6 units are found in the mineral denningite.

TeO2 melts at 732.6 °C, forming a red liquid.

Uses

It is used as an acousto-optic material.

Tellurium dioxide is also a conditional glass former, which means it will form a glass with

small molar% additions of a second compound such as an oxide or halide. TeO2 glasses have

high refractive indices and transmit into the mid-infrared part of the electromagnetic

spectrum, therefore they are of technological interest for optical waveguides. Tellurite glasses

have also been shown to exhibit Raman gain up to 30 times that of silica, useful in optical

fibre amplification.

Oxyacids of Selenium and Tellurium

Selenous and tellurous acids can be obtained by the hydrolysis of the tetrahalides,

TeCl4 + 3 H2O → 4 HCl + H2TeO3

or by slow evaporation of solutions of the dioxides in water. Unlike H2SO3, solid H2SeO3 can

be obtained. A large number of selenitesand tellurites are known, and typical preparations

include the following reactions.

CaO + SeO2 → CaSeO3

2 NaOH + SeO2 → Na2SeO3 + H2O

In either concentrated solutions or the molten acid, selenite is partially converted

by dehydration to pyroselenite, Se2O52−.

2HSeO3−⇆Se2O52

−+H2O

Tellurites also form a series of polytellurites (Te4O92−,

Te6O132−,

etc.) that can be obtained as

solid salts.

Selenic and telluric acids contain the central atom in the +6 oxidation state, but they are quite

different. Selenic acid, H2SeO4, behaves very much like H2SO4 in most of its chemical

properties. Most selenates are isomorphous with the corresponding sulfates, and some

pyroselenates are known. On the other hand, telluric acid has the formula H6TeO6 or

Te(OH)6. Although this formula is equivalent to H2TeO4·2H2O, the molecule is octahedral

and is not a dihydrate. It is isoelectronic with the

hydroxocomplexes Sn(OH)62− and Sb(OH)6

−. As a result of its structure, telluric acid is a

weak acid (see Chapter 6). Several types of salts of telluric acid are known in which one or

more of the hydrogen atoms is replaced (e.g., KH5TeO6 or K[TeO(OH)5], Hg3TeO6, and

Li2H4TeO6). Neither selenic nor telluric acid approaches the commercial importance of

sulfuric acid. But then, no compound does.

Oxides, Oxyacids, and Oxyanions of Se and Te

Selenium and tellurium form acidic dioxides and trioxides. Selenium dioxide, SeO2, reacts

with water, producing selenious acid, H2SeO3, a weak diprotic acid. Tellurium dioxide, TeO2,

is insoluble in water, but salts such as NaHTeO3 and Na2TeO3 form when TeO2 is dissolved

in an aqueous base.

Oxidation of H2SeO3 with hydrogen peroxide produces selenic acid, H2SeO4, a strong

diprotic acid that resembles sulfuric acid:

Telluric acid (also called orthotelluric acid) forms when TeO2 is treated with various aqueous

oxidizing agents. We might expect the formula for this acid to be H2TeO4 by analogy with

the corresponding acids of sulfur and selenium. However, orthotelluric acid is H6TeO6, which

can be viewed as the result of adding two molecules of H2O to H2TeO4. [The formula can

also be written Te(OH)6.] As in H5IO6(Figure 22.17), the central atom in H6TeO6 is

surrounded by six O atoms. The large sizes of Te and I permit them to bond to six

surrounding O atoms. Orthotelluric acid is a weak acid, capable of dissociating two protons in

aqueous solution (at 25°C, Ka1 = 2.4 × 10-8

, Ka2 = 1.0 × 10-11

).

A comparative study of Sulphur, Selenium, Tellurium and their compounds (hydrides,

oxides, halides).

Electronic Configuration of Group 16 Elements

Group 16 elements have 6 electrons in their valence shell and their general electronic

configuration is ns2np

4.

Element Electronic Configuration

Oxygen [He] 2s2 2p

4

Sulphur [Ne] 3s2 3p

4

Selenium [Ar] 3d10

4s2 4p

4

Tellurium [Kr] 4d10

5s2 5p

4

Polonium [Xe] 4f14

5d10

6s2 6

4

Atomic and Physical Properties and Their Trends

Atomic and Ionic Radii: The atomic and ionic radius increases as we move from

oxygen to polonium.

Ionization Enthalpy: Ionization enthalpy decreases with increase in the size of the

central atom. Therefore, it decreases as we move from oxygen to polonium since the size

of the atom increases as we move down.

Electron Gain Enthalpy: The electron gain enthalpy decreases with increase in the

size of the central atom moving down the group. Oxygen molecule has a less negative

electron gain enthalpy than sulfur. This is on the grounds that, oxygen, because of its

compressed nature encounter more repulsion between the electrons effectively present

and the approaching electron.

Electronegativity: The electronegativity decreases as we move down the group.

Therefore, it decreases as we move from oxygen to polonium due to increase in nuclear

size. Learn about Electronegativity here in detail.

Nature of the Group 16 Elements: Oxygen and sulfur are non-metals, selenium and

tellurium are metalloids and polonium is a metal under typical conditions. Polonium is a

radioactive element.

Allotropy: Each one of the element of group 16 displays allotropy. Oxygen has two

allotropes: Oxygen and Ozone. Sulphur exists as many allotropic forms but only two of

them are stable, which are: Rhombic Sulphur and Monoclinic Sulphur. Selenium and

Tellurium are found in both amorphous and crystalline forms.

The Melting and Boiling Points: As the atomic size increases from oxygen to

tellurium, the melting and boiling points also increase. The huge distinction between the

melting and boiling points of oxygen and sulfur might be clarified on the premise that

oxygen exists as a diatomic atom while sulfur exists as a polyatomic particle.

Oxidation States: The group 16 elements have a configuration of ns2 np

4 in their outer

shell, they may accomplish noble gas configuration either by the gain of two electrons,

framing M-2

or by sharing two electrons, in this manner shaping two covalent bonds.

Thus, these elements indicate both negative and positive oxidation states. The regular oxidation

states showed by the elements of group 16 incorporate -2, +2, +4, and + 6.

Chemical Properties

The group sixteen elements react with hydrogen to form hydrides of the sort H2E, where E could

be any element- oxygen, sulfur, selenium, tellurium or polonium.

The Physical States of Hydrides of Group 16 Elements

Water is an odourless and colourless liquid but the hydrides of the various elements of this group

are poisonous gases which are colourless with disagreeable smells.

The boiling point of these hydrides extraordinarily diminishes from water to hydrogen sulfide,

and after that increases. Water has an anomalously high boiling point since its particles are

bonded with each other by the hydrogen bonds in both its liquid as well as solid states.