Chemistry, Elements

A. NOBLE GASES

The noble gases are colorless, odorless, tasteless, and nonflammable under standard

conditions. They were once labeled group 0 in the periodic table because it was believed they

had a valence of zero, because the noble gases have full valence electron shells, meaning

their atoms cannot combine with those of other elements to form compounds..

Z Element No. of electrons/shell

2 helium 2

10 neon 2, 8

18 argon 2, 8, 8

36 krypton 2, 8, 18, 8

54 xenon 2, 8, 18, 18, 8

86 radon 2, 8, 18, 32, 18, 8

Helium - He

http://en.wikipedia.org/wiki/Chemical_compound

http://en.wikipedia.org/wiki/Atom

http://en.wikipedia.org/wiki/Electron_shells

http://en.wikipedia.org/wiki/Valence_(chemistry)

PropertiesGaseous chemical element, symbol: He, atomic number: 2 and atomic weight 4,0026 g/mol.

Helium is one of the noble gases of group O in the periodic table. It’s the second lightest

element. The main helium source in the world is a series of fields of natural gas in the United

States.

Helium is a colourless, odourless, insipid and non-toxic gas. It’s less soluble in water than any

other gas. It’s the less reactive element and doesn’t essentially form chemical compounds.

The density and viscosity of helium vapour are very low. The termic conductivity and the

caloric content are exceptionally high. Helium can be liquefied, but its condensation

temperature is the lowest among all the known substances.

Applications:Helium has many unique properties: low boiling point, low density, low solubility, high

thermal conductivity and inertness, so it is use for any application which can explioit these

properties. Helium was the first gas used for filling balloons and dirigibles. This application

goes on in altitude research and for meteorological balloons. The main use of helium is as an

inert protection gas in autogenous welding. Its biggest potential is found in applications at

very low temperatures. Helium is the only cooler which is capable of reaching temperatures

lower than 15 K (-434ºF). The main application of ultralow temperature is in the development

of the superconductivity state, in which the resistance to the electricity flux is almost zero.

Other applications are its use as pressurizing gas in liquid propellants for rockets, in helium-

oxygen mixtures for divers, as working fluid in nuclear reactors cooled down by gas and as

gas carrier in chemical analysis by gas chromatography.

Helium in the environment:Helium is the second most abundant element in the known universe, after hydrogen. Helium

constitutes the 23% of all elemental matter measured by mass. Helium is formed in The Earth

Atomic number 2Atomic mass 4.00260 g.mol -1

Electronegativity according to Pauling unknown

Density 0.178*10 -3 g.cm -3 at 20 °CMelting point - 272.2 (26 atm) °CBoiling point - 268.9 °CVanderwaals radius 0.118 nmIonic radius unknownIsotopes 2Electronic shell 1s 2

Energy of first ionization 2372 kJ.mol -1

Discovered by Sir Ramsey in 1895

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by natural radioactive decay of heavier elements. Most of this helium migrates to the surface

and enters the atmosphere. It could be logical to think that the helium concentration in the

atmosphere was higher than it is (5,25 parts per million at sea level). Nevertheless, its low

molecular weight allows it to escape to space at the same rate of its formation. There is an

about 1000 km layer in the heterosphere at 600 miles where helium is the dominant gas

(although the total pressure is very low). Natural gases contain higher helium concentrations

than the atmosphere.

Helium is the 71st most abundant element in the Earth's crust where it is found in 8 parts per

billion (109).

Health effects of helium:Effects of exposure: The substance can be absorbed into the body by inhalation. Inhalation:

High voice. Dizziness. Dullness. Headache. Suffocation. Skin: on contact with liquid:

frostbite. Eyes: on contact with liquid: frostbite. Inhalation risk: On loss of containment this

gas can cause suffocation by lowering the oxygen content of the air in confined areas. Check

oxygen content before entering area.

Neutral helium at standard conditions is non-toxic, plays no biological role and is found in

trace amounts in human blood

Neon - Ne

Atomic number: 10

Atomic mass: 20.179 g.mol -1

Electronegativity according to Pauling: unknown

Density: 0.9*10 -3 g.cm-3 at 20°C

Melting point: -249 °C

Boiling point: -246 °C

Vanderwaals radius: 0.16 nm

Ionic radius: unknown

Isotopes: 3

Electronic shell: [ He ] 2s22p6

Energy of first ionization: 2080 kJ.mol -1

Energy o second ionization: 3952 kJ.mol -1

Standard potential: 6122 kJ.mol -1

Date of Discovery: 1898

Discoverer: Sir William Ramsay

Name Origin: Form the Greek word neos (new)

Uses: lighting

Obtained From: liquid air

Properties:Neon is the second-lightest noble gas, its colour is reddish-orange in a vacuum discharge tube

and in neon lamps. The the refrigerating capacity of helium is over 40 times the one of

liquid helium and three times that of liquid hydrogen (on a per unit volume basis). It is a less

expensive refrigerant than helium in most applications.

Even though neon is for most practical purposes an inert element, it can form an exotic

compound with fluorine in the laboratory. It is not known for certain if this or any neon

compound exists naturally but some evidence suggests that this may be true. The ions, Ne+,

(NeAr)+, (NeH)+, and (HeNe+) are have also been observed from optical and mass

spectrometric research. In addition, neon forms an unstable hydrate.

Applications:The reddish-orange color emitted in neon lights is widely used to make advertising signs.

Neon is also used generically for these types of lights when in reality many other gases are

used to produce different colors of light. Other uses of neon include high-voltage indicators,

lightning arrestors, wave meter tubes and television tubes. Neon and helium are used to make

a type of gas laser. Liquefied neon is commercially used as an economical cryogenic

refrigerant.

Neon in the environment:Although neon is the forth most abundant element in the universe, only 0.0018% in volume

of the earth's atmosphere is neon.

Neon is usually found in the form of a gas with molecules consisting of a single Neon atom.

Neon is a rare gas that is found in the Earth's atmosphere at 1 part in 65,000.

Health effects of neon:Routes of exposure: The substance can be absorbed into the body by inhalation.

Inhalation risk: On loss of containment this liquid evaporates very quickly causing

supersaturation of the air with serious risk of suffocation when in confined areas.

Effects of exposure: Inhalation: Simple asphyxiant. Skin: On contact with liquid: frostbite.

Eyes: On contact with liquid: frostbite.

Inhalation: This gas is inert and is classified as a simple asphyxiant. Inhalation in excessive

concentrations can result in dizziness, nausea, vomiting, loss of consciousness, and death.

Death may result from errors in judgment, confusion, or loss of consciousness which prevent

self-rescue. At low oxygen concentrations, unconsciousness and death may occur in seconds

without warning.

The effect of simple asphyxiant gases is proportional to the extent to which they diminish the

amount (partial pressure) of oxygen in the air that is breathed. The oxygen may be diminished

to 75% of it's normal percentage in air before appreciable symptoms develop. This in turn

requires the presence of a simple asphyxiant in a concentration of 33% in the mixture of air

and gas. When the simple asphyxiant reaches a concentration of 50%, marked symptoms can

be produced. A concentration of 75% is fatal in a matter of minutes.

Symptoms: The first symptoms produced by a simple asphyxiant are rapid respirations and

air hunger. Mental alertness is diminished and muscular coordination is impaired. Later

judgment becomes faulty and all sensations are depressed. Emotional instability often results

and fatigue occurs rapidly. As the asphyxia progresses, there may be nausea and vomiting,

prostration and loss of consciousness, and finally convulsions, deep coma and death.

Environmental effects of neon:

Neon is a rare atmospheric gas and as such is non-toxic and chemically inert. Neon poses no

threat to the environment, and can have no impact at all because it's chemically unreactive

and forms no compounds.

No known ecological damage caused by this element.

Argon - Ar

Name: Argon

Symbol: Ar

Atomic Number: 18

Atomic Mass: 39.948 amu


Melting Point: -189.3 °C (83.85 K, -308.74 °F)

Boiling Point: -186.0 °C (87.15 K, -302.8 °F)

Number of Protons/Electrons: 18

Number of Neutrons: 22

Classification: Noble Gas

Crystal Structure: Cubic

Density @ 293 K: 1.784 g/cm3

Color: Colorless Gas

Date of Discovery: 1894

Discoverer: Sir William Ramsay

Name Origin: From the Greek word argon (inactive)

Uses: Lighting

Obtained From: air

Description:Argon was suspected to be present in air by Henry Cavendish in 1785 but wasn't discovered

until 1894 by Lord Rayleigh and Sir William Ramsay.

Argon is the third noble gas, in period 8, and it makes up about 1% of the Earth's atmosphere.

Argon has approximately the same solubility as oxygen and it is 2.5 times as soluble in water

as nitrogen . This chemically inert element is colorless and odorless in both its liquid and

gaseous forms. It is not found in any compounds.

This gas is isolated through liquid air fractionation since the atmosphere contains only 0.94%

argon. The Martian atmosphere in contrast contains 1.6% of Ar-40 and 5 ppm Ar-36. World

production exceeds 750.000 tonnes per year, the supply is virtually inexhaustible.

Applications:Argon does not react with the filament in a lightbulb even under high temperatures, so is used

in lighting and in other cases where diatomic nitrogen is an unsuitable (semi-)inert gas.

Argon is perticularly important for the metal industry, being used as an inert gas shield in arc

welding and cutting. Other uses incude non-reactive blanket in the manufacture

of titanium and other reactive elements and as a protective atmosphere for

growing silicon and germanium crystals. Argon-39 has been used for a number of

applications, primarily ice coring. It has also been used for ground water dating. Argon is also

used in technical SCUBA diving to inflate the drysuit, due to its nonreactive, heat isolating

effect.

Argon as the gap between the panes of glass provides better insulation because it is a poorer

conductor of heat than ordinary air. The most exotic use of argon is in the tyre of luxury cars.

Argon in the environment:

http://www.lenntech.com/Periodic-chart-elements/Ge-en.htm

http://www.lenntech.com/Periodic-chart-elements/Si-en.htm

http://www.lenntech.com/Periodic-chart-elements/Ti-en.htm

http://www.lenntech.com/Periodic-chart-elements/N-en.htm


http://chemicalelements.com/groups/noblegases.html

In earth's atmosphere, Ar-39 is made by cosmic ray activity, primarily with Ar-40. In the

subsurface environment, it is also produced through neutron-capture by K-39 or alpha

emission by calcium. Argon-37 is produced from the decay of calcium-40, the result of

subsurface nuclear explosions. It has a half-life of 35 days.

Argon is present in some potassium minerals because of radiactive decay of the

isotope potassium-40

Health effects of argon:Routes of exposure: The substance can be absorbed into the body by inhalation.

Inhalation risk: On loss of containment this liquid evaporates very quickly causing

supersaturation of the air with serious risk of suffocation when in confined areas.

Effects of exposure: Inhalation: Dizziness. Dullness. Headache. Suffocation. Skin: On

contact with liquid: frostbite. Eyes: On contact with liquid: frostbite.

Inhalation: This gas is inert and is classified as a simple asphyxiant. Inhalation in excessive




without warning.







Symptoms: The first symptoms produced by a simple asphyxiant are rapid respirations and

air hunger. Mental alertness is diminished and muscular coordination is impaired. Later

judgment becomes faulty and all sensations are depressed. Emotional instability often results

and fatigue occurs rapidly. As the asphyxia progresses, there may be nausea and vomiting,


Environmental effects of argon:

No known ecological damage caused by argon.

No adverse environmental consequences are expected. Argon gas occurs naturally in the

environment. The gas will dissipate rapidly in well ventilated areas.

The effects of argon on plants or animals is not currently known. It is not expected to harm

aquatic life.

Argon does not contain any ozone depleting chemicals and is not listed as a marine pollutant

by DOT (Department of Transportation, USA).

http://en2.wikipedia.org/wiki/Potassium

http://www.lenntech.com/Periodic-chart-elements/Ca-en.htm

Krypton - Kr

Atomic number: 36


Density: 3.73 10-3 g.cm-3 at 20°C

Melting point: - 157 °C

Boiling point: - 153° C


Isotopes: 15

Electronic shell: [ Ar ] 3d10 4s2 4p6


Energy of second ionization: 2350.4 kJ.mol -1

Energy of third ionization: 3565 kJ.mol -1

Discovered by: Sir Ramsay in 1898

Description: Krypton is present in the air at about 1 ppm. The atmosphere of Mars contains a little (about

0.3 ppm) of krypton. It is characterised by its brilliant green and orange spectral lines. The

spectral lines of krypton are easily produced and some are very sharp. In 1960 it was

internationally agreed that the fundamental unit of length, the metre, should be defined as 1 m

= 1,650,763.73 wavelengths (in vacuo) of the orange-red line of Kr-33.

Under normal conditions krypton is colourless, odourless, fairly expensive gas. Solid krypton

is a white crystalline substance with a face-centered cubic structure which is common to all

the "rare gases". Krypton difluoride, KrF2, has been prepared in gram quantities and can be

made by several methods. Other compounds are unstable, unless isolated in a matrix at very

low temperatures.

Applications:Krypton is used to fill electric lamp bulbs which are filled with a mixture of krypton and

argon, and for various electronic devices. Krypton is also used in photographic projection

lamps, in very high-powered electric arc lights used at airports and in some strobo-lamps,

because it has an extremely fast respons to an electric current.

A mixture of stable and unstable isotopes of krypton is produced by slow neutron fission of

uranium in nuclear reactors as Kripron-85, its most stable isotope. It is used to detect leaks in

sealed containers, to excite phosphors in light sources with no external source of energy, and

in medicine to detect abnormal heart openings.

Krypton in the environment

Krypton might be one of the rarest gases in the atmosphere, but in total there are more than 15

billion tonnes of this metal circulating in the planet, of which only about 8 tonnes a year are

extracted, via liquid air.

Health effects of krypton:Inhalation: This gas is inert and is classified as a simple asphyxiant. Inhalation in excessive




without warning.







Symptoms: The first symptoms produced by a simple asphyxiant are rapid respirations and air

hunger. Mental alertness is diminished and muscular coordination is impaired. Later judgment

becomes faulty and all sensations are depressed. Emotional instability often results and

fatigue occurs rapidly. As the asphyxia progresses, there may be nausea and vomiting,


Environmental effects of krypton

Krypton is a rare atmospheric gas and as such is non-toxic and chemically inert. The extreme

cold temperature (-244oC) will freeze organisms on contact, but no long term ecological

effects are anticipated.

Disposal considerations: When disposal becomes necessary, vent gas slowly to a well-

ventilated out door location remote from personnel work areas and building air intakes. Do

not dispose of any residual gas in compressed gas cylinders. Return cylinders to the supplier

with residual pressure, the cylinder valve tightly closed. Please be advised that state and local

requirements for waste disposal may be more restrictive or otherwise different from federal

regulations. Consult state and local regulations regarding the proper disposal of this material.

Xenon - Xe

Atomic number: 54



Density: 5.9*10-3g.cm-3 at 20°C


Boiling point: - 107 °C


Isotopes: 21

Electronic shell: [ Kr ] 4d10 5s25p6


Discovered by: Sir Ramsay 1898

Description:Xenon is a rare, odorless, colourless, tasteless, chamically unreactive gas. It was regarded as

completely inert until, in 1962, Neil Barlett reported synthesis of xenon haxafluoroplatinate.

In a gas filled tube xenon emits blue light when excited by electrical discharge.

Applications:Xenon has relatively little commercial use. It is used in photographic flash lamps,

stroboscopic lamps, high-intensitive arc-lamps for motion picture projection and high-

pressure arc lamps to product ultraviolet light (solar simulators). Other uses are as general

anaesthetic, xenon 'blue' headlights and fog lights are used on some vehicles and are said to

be less tiring on the eyes. They illuminate road signs and markings better than conventional

lights.

Xenon in the environment:Xenon si a trace gas in the Earth's atmosphere, occurring in 1 part in 20 million. The only

commercial source of xenon is from industrial liquid-air plants. World production is less than

1 tonne per year, although reserved of xenon gas in the atmosphere amount to 2 billion

tonnes.

Health effects of xenon:Inhalation: This gas is inert and is classified as a simple asphyxiant. Inhalation in excessive




without warning.






be produced. A concentration of 75% is fatal in a matter of minutes. Symptoms: The first


symptoms produced by a simple asphyxiant are rapid respirations and air hunger. Mental

alertness is diminished and muscular coordination is impaired. Later judgment becomes faulty

and all sensations are depressed. Emotional instability often results and fatigue occurs rapidly.

As the asphyxia progresses, there may be nausea and vomiting, prostration and loss of

consciousness, and finally convulsions, deep coma and death.

This agent is not considered a carcinogen.

Effects \of xenon on the environment:Xenon is a rare atmospheric gas and as such is non-toxic and chemically inert. The extreme

cold temperature (-244oC) will freeze organisms on contact, but no long term ecological

effects are anticipated.

Disposal considerations: When disposal becomes necessary, vent gas slowly to a well-

ventilated out door location remote from personnel work areas and building air intakes. Do

not dispose of any residual gas in compressed gas cylinders. Return cylinders to the supplier

with residual pressure, the cylinder valve tightly closed. Please be advised that state and local

requirements for waste disposal may be more restrictive or otherwise different from federal

regulations. Consult state and local regulations regarding the proper disposal of this material.

Radon - Rn

Atomic number: 86

Atomic mass: (222) g.mol -1


Density: 9.96*10-3g.cm3 at 20°C


Boiling point: - 62 °C

Vanderwaals radius: unknown

Ionic radius: unknown

Isotopes: 7Electronic shell: [ Xe ] 4f14 5d10 6s2 6p6


Discovered: Fredrich Ernst Dorn in 1898

Description:

Radon is colorless at standard temperature and pressure and it is the most dense gas known.

At temperature below it's freezing point is has a brilliant yellow phosphorescence. It is

chemically unreactive, it is highly radioactive and has a short half life.

Applications:Radon was sometimes used in hospitals to treat cancer and was produced as needed and

delivered in sealed gold needles. Radon is used in hydrologic research, because of it's rapid

loss to air. It is also used in geologic research and to track air masses.

Radon in the environment:Radon can be found in some spring water and hot springs. There is anyway a detectable

amount of radon in the atmosphere. Radon collects over samples of radium 226 at the rate of

around 0.001 cm3/day per g of radium.

Health effects of radon:Radon occurs in the environment mainly in the gaseous phase. Consequently, people are

mainly exposed to radon through breathing air. Background levels of radon in outside air are

generally quite low, but in indoor locations radon levels in air may be higher. In homes,

schools and buildings radon levels are increased because radon enters the buildings through

cracks in the foundations and basements. Some of the deep wells that supply us with drinking

water may also contain radon. As a result a number of people may be exposed to radon

through drinking water, as well as through breathing air. Radon levels in groundwater are

fairly high, but usually radon is quickly released into air as soon as the groundwater enters

surface waters. Exposure to high levels of radon through breathing air is known to cause lung

diseases. When long-term exposure occurs radon increases the chances of developing lung

cancer. Radon can only cause cancer after several years of exposure. Radon may be

radioactive, but it gives of little actual gamma radiation. As a result, harmful effects from

exposure to radon radiation without actual contact with radon compounds are not likely to

occur. It is not known whether radon can cause health effects in other organs besides the

lungs. The effects of radon, which is found in food or drinking water, are unknown.

Environmental effects of radon:

Radon is a radioactive compound, which rarely occurs naturally in the environment. Most of

the radon compounds found in the environment derive from human activities. Radon enters

the environment through the soil, through uranium and phosphate mines, and through coal

combustion. Some of the radon that is located in the soil will move to the surface and enter

the air through vaporization. In the air, radon compounds will attach to dust and other

particles. Radon can also move downwards in the soil and enter the groundwater. However,

most of the radon will remain in the soil. Radon has a radioactive half-life of about four days;

this means that one-half of a given amount of radon will decay to other compounds, usually

less harmful compounds, every four days.

B. HALOGEN GASES

Owing to their high reactivity, the halogens are found in the environment only

in compounds or as ions. Halide ions and oxoanions such as iodate(IO3−) can be found in

many minerals and in seawater. Halogenated organic compounds can also be found as natural

products in living organisms. In their elemental forms, the halogens exist as diatomic

molecules, but these only have a fleeting existence in nature and are much more common in

the laboratory and in industry. At room temperature and pressure, fluorine and chlorine are

gases, bromine is a liquid and iodine and astatine are solids; Group 17 is therefore the only

periodic table group exhibiting all three states of matter at room temperature.

http://en.wikipedia.org/wiki/Temperature

http://en.wikipedia.org/wiki/State_of_matter

http://en.wikipedia.org/wiki/Diatomic_molecule

http://en.wikipedia.org/wiki/Diatomic_molecule

http://en.wikipedia.org/wiki/Organic_halide

http://en.wikipedia.org/wiki/Iodate

http://en.wikipedia.org/wiki/Oxyanion

http://en.wikipedia.org/wiki/Halide

http://en.wikipedia.org/wiki/Ion

http://en.wikipedia.org/wiki/Chemical_compound

http://en.wikipedia.org/wiki/Reactivity_(chemistry)

Z Element No. of electrons/shell

9 fluorine 2, 7

17 chlorine 2, 8, 7

35 bromine 2, 8, 18, 7

53 iodine 2, 8, 18, 18, 7

85 astatine 2, 8, 18, 32, 18, 7

Fluorine - F

Atomic number: 9Atomic mass: 18.998403 g.mol-1

Electronegativity according to Pauling: 4Density: 1.8*10-3 g.cm-3at 20°C

Melting point: -219.6 °C

Boiling point: -188 °C


Ionic radius: 0.136 nm (-1) ; 0.007 (+7)

Isotopes: 2Electronic shell: [ He ] 2s22p5

Energy of first ionization: 1680.6 kJ.mol -1

Energy of second ionization: 3134 kJ.mol -1

Energy of third ionization: 6050 kJ mol-1

Standard potential: 2.87 V

Discovered by: Moissan in 1886

Desciption: Fluorine is an univalent poisonous gaseous halogen, it is pale yellow-green and it is the most

chemically reactive and electronegative of all the elements. Fluorine readily forms

compounds with most other elements, even with the noble gases krypton, xenon and radon. It

is so reactive that glass, metals, and even water, as well as other substances, burn with a

bright flame in a jet of fluorine gas.

In aqueous solution, fluorine commonly occurs as the fluoride ion F-. Fluorides are

compounds that combine fluoride with some positively charged counterpart.

Applications:Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor

manufacturing, flat panel display production and MEMs fabrication.

Fluorine is indirectly used in the production of low friction plastics such as teflon and in

halons such as freon, in the production of uranium. Fluorochlorohydrocarbons are used

extensively in air conditioning and in refrigeration.

Fluorides are often added to toothpaste and, somewhat controversially, to municipal water

supplies to prevent dental cavities. Fore more information visit our page on mineral water.

Fluorine in the environment:Annual world production of the mineral fluorite in around 4 million tonnes, and there are

around 120 million tonnes of mineral reserves. The main mining areas for fluorite are China,

Mexico and Western Europe.

Fluorine occurs naturally in the earth's crust where it can be found in rocks, coal and clay.

Fluorides are released into the air in wind-blown soil. Fluorine is the 13th most aboundant

element in the Earth's crust: 950 ppm are contanined in it. Soils contain approximatively 330

ppm of fluorine, ranging from 150 to 400 ppm. Some solis can have as much as 1000 ppm

and contaminated solis have been found with 3500 ppm. Hydrogen fluorides can be released

into air through combustion processes in the industry. Fluorides that are found in air will

eventually drop onto land or into water. When fluorine is attached to very small particles it

can remain in the air for a long period of time.

In the atmosphere 0.6 ppb of fluorine are present as salt spray and organicochloride

compounds. Up to 50 ppb has been recorded in city environments.

Health effects of fluorine:Small amounts of fluorine are naturally present in water, air, plants and animals. As a result

humans are exposed to fluorine through food and drinking water and by breathing air. Fluorine

can be found in any kind of food in relatively small quantities. Large quantities of fluorine can

be found in tea and shellfish. luorine is essential for the maintenance of solidity of our bones.

Fluorine can also protect us from dental decay, if it is applied through toothpaste twice a day.

If fluorine is absorbed too frequently, it can cause teeth decay, osteoporosis and harm to

kidneys, bones, nerves and muscles. luorine gas is released in the industries. This gas is very

dangerous, as it can cause death at very high concentrations. At low concentrations it causes

eye and nose irritations.

Environmental effects of fluorine:When fluorine from the air ends up in water it will settle into the sediment. When it ends up

in soils, fluorine will become strongly attached to soil particles. In the environment fluorine

cannot be destroyed; it can only change form. luorine that is located in soils may accumulate

in plants. The amount of uptake by plants depends upon the type of plant and the type of soil

and the amount and type of fluorine found in the soil. With plants that are sensitive for

fluorine exposure even low concentrations of fluorine can cause leave damage and a decline

in growth. Too much fluoride, wheater taken in form the soil by roots, or asdorbed from the

atmosphere by the leaves, retards the growth of plants and reduces crop yields. Those more

affected are corns and apricots. Animals that eat fluorine-containing plants may accumulate

large amounts of fluorine in their bodies. Fluorine primarily accumulates in bones.

Consequently, animals that are exposed to high concentrations of fluorine suffer from dental

decay and bone degradation. Too much fluorine can also cause the uptake of food from the

paunch to decline and it can disturb the development of claws. Finally, it can cause low birth-

weights.

Chlorine - Cl

Atomic number: 17


Electronegativity according to Pauling: 3.0

Density: 3.21*10 -3 g.cm -3 at 20 °C

Melting point: -101 °C

Boiling point: -34.6 °C


Ionic radius: 0.184 (-2) nm ; 0.029 nm (+6)

Isotopes: 4Electronic shell: [Ne] 3s23p5




Standard potential: - 1.36 V

Discovered by: Carl Wilhelm Scheele in 1774

Description:The pure chemical element has the physical form of a diatomic green gas. The name chlorine

is derived from chloros, meaning green, referring to the color of the gas. Chlorine gas is two

and one half times as heavy as air, has an intensely disagreeable suffocating odor, and is

exceedingly poisonous. In its liquid and solid form it is a powerful oxidizing, bleaching, and

disinfecting agent.

This element is a part of the halogen series forming salts. It is extracted from chlorides

through oxidation and electrolysis. Chlorine gas is greenish-yellow and combines readily with

nearly all other elements.

Applications:Chlorine is an important chemical in water purification, in disinfectants, in bleach and in

mustard gas. Chlorine is also used widely in the manufacture of many products and items

directly or indirectly, i.e. in paper product production, antiseptic, dyestuffs, food, insecticides,

paints, petroleum products, plastics, medicines, textiles, solvents, and many other consumer

products. It is used to kill bacteria and other microbes from drinking water supplies.

Chlorine is involved in beaching wood pulp for paper making, bleach is also used industrially

to remove ink from recycle paper.

Chlorine often imparts many desired properties in an organic compound when it is substituted

for hydrogen (synthetic rubber), so it is widely use in organic chemistry, in the production of

chlorates, chloroform, carbon tetrachloride, and in the bromine extraction.

Chlorine in the environment:In nature it is only found combined with other elements chiefly sodium in the form of

common salt (NaCl), but also in carnallite, and sylvite. Chlorides make up much of the salt

dissolved in the earth's oceans: about 1.9 % of the mass of seawater is chloride ions.

The amount of chloride in soils varies according to the distance from the sea. The average in

top soils is about 10 ppm. Plants contain various amount of chlorine; it is an essential

microutrient for higher plants where is concentrates in the chloroplasts. Growth suffers if the

amount of chloride in the soil fall below 2 ppm, but it rarely happens. The upper limit of

tolerance varies according to the crop.

Health effects of chlorine:Chlorine is a highly reactive gas. It is a naturally occurring element. The largest users of

chlorine are companies that make ethylene dichloride and other chlorinated solvents,

polyvinyl chloride (PVC) resins, chlorofluorocarbons, and propylene oxide. Paper companies

use chlorine to bleach paper. Water and wastewater treatment plants use chlorine to reduce

water levels of microrganisms that can spread disease to humans (disinfection).Exposure to

chlorine can occur in the workplace or in the environment following releases to air, water, or

land. People who use laundry bleach and swimming pool chemicals containing chlorine

products are usually not exposed to chlorine itself. Chlorine is generally found only in

industrial settings. Chlorine enters the body breathed in with contaminated air or when

consumed with contaminated food or water. It does not remain in the body, due to its

reactivity. Effects of chlorine on human health depend on how the amount of chlorine that is

present, and the length and frequency of exposure. Effects also depend on the health of a

person or condition of the environment when exposure occurs. Breathing small amounts of

chlorine for short periods of time adversely affects the human respiratory system. Effects

differ from coughing and chest pain, to water retention in the lungs. Chlorine irritates the

skin, the eyes, and the respiratory system. These effects are not likely to occur at levels of

chlorine that are normally found in the environment.Human health effects associated with

breathing or otherwise consuming small amounts of chlorine over long periods of time are not

known. Some studies show that workers develop adverse effects from repeat inhalation

exposure to chlorine, but others will not.

Environmental effects of chlorine:Chlorine dissolves when mixed with water. It can also escape from water and enter air under

certain conditions. Most direct releases of chlorine to the environment are to air and to surface

water. Once in air or in water, chlorine reacts with other chemicals. It combines with inorganic

material in water to form chloride salts, and with organic material in water to form chlorinated

organic chemicals.Because of its reactivity chlorine is not likely to move through the ground

and enter groundwater. Plants and animals are not likely to store chlorine. However, laboratory

studies show that repeat exposure to chlorine in air can affect the immune system, the blood,

the heart, and the respiratory system of animals. Chlorine causes environmental harm at low

levels. Chlorine is especially harmful to organisms living in water and in soil.

Bromine - Br

Atomic number35

Atomic mass 79.904 g.mol -1

Electronegativity according to Pauling 2.8

Density 3.1 g.cm-3 at 20°C

Melting point - 7.2 °C

Boiling point 58.8 °C

Vanderwaals radius 0.165 nm

Ionic radius 0.195 nm (-1)

Isotopes 10

Electronic shell [ Ar ] 3d10 4s2 4p5

Energy of first ionization 1142.7 kJ.mol -1

Standard potential 1.08 V

Discovered by Anthoine Balard in 1826

Description:At ambient temperature bromine is a brownish-red liquid. It has a similarly colored vapor

with an offensive and suffocating odor. It is the only nonmetallic element that is liquid under

ordinary conditions, it evaporates easily at standard temperature and pressures in a red vapor

that has a strong disagreeable odor resembling that of chlorine. Bromine is less active

chemically than chlorine and fluorine but is more active than iodine; its compounds are

similar to those of the other halogens. Bromine is soluble in organic solvents and in water.

Applications:Bromine is used in industry to make organobromo compounds. A major one was

dibromoethane an agent for leaded gasoline, before they were largely phased out due to

environmental considerations. Other organobromines are used as insecticides, in fire

extinguishers and to make pharmaceuticals. Bromine is used in making fumigants, dyes,

flameproofing agents, water purification compounds, sanitizes, medicinals, agents for

photography and in brominates vegetable oil, used as emulsifier in many citrus-flavoured solft

drinks.

Bromine in the environmentBromine is a naturally occurring element that can be found in many inorganic substances.

Humans however, have many years ago started the introduction of organic bromines in the

environment. These are all compounds that are not natural and can cause serious harm to

human health and the environment.In diffuse crustal rock bromine naturally occurs as

bromide salts. Bromine salts have accumulated in sea water (85 ppm), from which bromine is

extracted. World production of bromine is more than 300.000 tonnes per year; the three main

producing countries are US, Istrael and the UK. In this last case it is extracted from sea water

at a plant on the coast of Anglesey, Wales. In diffuse crustal rock bromine naturally occurs as

bromide salts. Bromine salts have accumulated in sea water (85 ppm), from which bromine is

extracted. World production of bromine is more than 300.000 tonnes per year; the three main

producing countries are US, Istrael and the UK. In this last case it is extracted from sea water

at a plant on the coast of Anglesey, Wales.

Health effects of bromine:Bromine is corrosive to human tissue in a liquid state and its vapors irritate eyes and throat.

Bromine vapors are very toxic with inhalation. Humans can absorb organic bromines through

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the skin, with food and during breathing. Organic bromines are widely used as sprays to kill

insects and other unwanted pests. But they are not only poisonous to the animals that they are

used against, but also to larger animals. In many cases they are poisonous to humans, too. The

most important health effects that can be caused by bromine-containing organic contaminants

are malfunctioning of the nervous system and disturbances in genetic materials.But organic

bromines can also cause damage to organs such as liver, kidneys, lungs and milt and they can

cause stomach and gastrointestinal malfunctioning. Some forms of organic bromines, such as

ethylene bromine, can even cause cancer. Inorganic bromines are found in nature, but

whereas they occur naturally humans have added too much through the years. Through food

and drinking water humans absorb high doses of inorganic bromines. These bromines can

damage the nervous system and the thyroid gland.

Environmental effects of bromine:Organic bromines are often applied as disinfecting and protecting agents, due to their

damaging effects on microorganisms. When they are applied in greenhouses and on farmland

they can easily rinse off to surface water, which has very negative health effects on daphnia,

fishes, lobsters and algae. Organic bromines are also damaging to mammals, especially when

they accumulate in the bodies of their preys. The most important effects on animals are nerve

damage and next to that DNA damage, which can also enhance the chances of development

of cancer. The uptake of organic bromine takes place through food, through breathing and

through the skin. Organic bromines are not very biodegradable; when they are decomposed

inorganic bromines will consist. These can damage the nerve system when high doses are

absorbed. It has occurred in the past that organic bromines ended up in the food of cattle.

Thousands of cows and pigs had to be killed in order to prevent contagion of humans. The

cattle suffered from symptoms such as liver damage, loss of sight and depletion of growth,

decrease of immunity, decreasing milk production and sterility and malformed children.

http://www.lenntech.com/water-disinfection/disinfectants-bromine.htm

Iodine - I

Atomic number 53




Melting point 114 °C

Boiling point 184 °C


Ionic radius 0.216 nm (-1) ; 0.05 nm (+7)

Isotopes 15

Electronic shell [ Kr ] 4d10 5s25p5

Energy of first ionisation 1008.7 kJ.mol -1

Standard potential + 0.58 V ( I2/ I- )

Discovered Bernard Courtois in 1811

Description:Iodine is a non-metallic, dark-gray/purple-black, lustrous, solid element. Iodine is the most

electropositive halogen and the least reactive of the halogens even if it can still form

compounds with many elements. Iodine sublime easily on heating to give a purple vapour.

Iodine dissolves in some solvents, such as carbon tetrachloride and it is only slightly soluble

in water.

Applications:Iodine is used in medical treatment as tincture and iodioform, it is employed in the

preparation of certain drugs and in the manufacture of some printing inks and dyes. Silver

iodine is used in photography. Iodine is added to almost all the table salt and is used as a

supplement to animal feed. It is also an ingredient of water purification tablets that are used

for drinking water preparation.

For many of these uses iodine is turned into iodides.

Iodine in the environment:Iodine is added to nearly any kind of salt that is applied. It is an ingredient of bread, sea fish

and oceanic plants. Iodine is naturally present in the ocean and some sea fish and water plants

will store it in their tissues.

Iodine can be found naturally in air, water and soil. The most important sources of natural

iodine are the oceans. About 400.000 tonnes of iodine escape from the oceans every year as

iodide in sea spray or as iodide, hydrichloric acid and methyl iodide, produced by marine

organisms. Much of it is deposited on land where it may become part of the biocycle.

There are some iodine-containing minerals, such as alutarite, found in Chile and iodargyte,

found in Colorado, Nevada and New Mexico. World-wide industrial production of iodine is

about 13.000 tonnes per year, mainly in Chile and Japan, plus small amounts in Russia and

USA. Iodine is extracted from natural brines and oil brines, which have up to 100 ppm of the

element or form chilean nitrate deposits. Known reserves of easily accessible iodine amount

is around 2 million tonnes.

Health effects of iodine:Many medicines and cleansers for skin wounds contain iodine. Iodine is a building material of

thyroid hormones that are essential for growth, the nervous system and the metabolism.

Humans that eat little to no bread can experience iodine shortages. The function of the thyroid

gland will than slow down and the thyroid gland will start swelling up. This phenomenon is

called struma. This condition is rare now as table salt is dosed with a little iodide. Large

quantities of iodine can be dangerous because the thyroid gland will labour too hastily. This

affects the entire body; it causes disturbed heartbeats and loss of weight. Elemental iodine, I2,

is toxic, and its vapour irritates the eyes and lungs. The maximum allowable concentration in

air when working with iodine is just 1 mg m-3. All iodides are toxic if taken in excess. Iodine

131 is one of the radionuclides involved in atmospheric testing of nuclear weapons, which

began in 1945, with a US test, and ended in 1980 with a Chinese test. It is among the long-

lived radionuclides that have produced and will continue to produce increased cancers risk for

decades and centuries to come. Iodine 131 increases the risk of cancer and possibly other

diseases of the thyroid and those caused by thyroid hormonal deficiency.

Environmental effects of iodine:Iodine in air can combine with water particles and precipitate into water or soils. Iodine in

soils will combine with organic matter and remain in the same place for a long time. Plants

that grow on these soils may absorb iodine. Cattle and other animals will absorb iodine when

they eat these plants. Iodine in surface water will vaporize and re-enter the air as a result.

Humans also add iodine gas to the air, by burning coal or fuel oil for energy. But the amount

of iodine that enters the air through human activity is fairly small compared to the amount

that vaporizes from the oceans. Iodine may be radioactive. The radioactive isotopes are

formed naturally during chemical reactions in the atmosphere. Most radioactive isotopes of

iodine have very short half-lives and will reshape into stable iodine compounds quickly.

However, there is one radioactive form of iodine that has a half-live of millions of years and

that is seriously harmful to the environment. This isotope enters the air from nuclear power

plants, where it is formed during uranium and plutonium processing. Accidents in nuclear

power plants have caused the release of large amounts of radioactive iodine into air.

Astatine - At

Atomic number85

Atomic mass (210) g.mol -1


Density unknown


Boiling point 337 °C (estimation)

Vanderwaals radius unknown

Ionic radius unknown

Isotopes 7

Electronic shell [ Xe ] 4f14 5d10 6s2 6p5

Energy of first ionisation (926) kJ.mol -1

Discovered by D.R. Corson 1940

Description:

Astatine is a highly radioactive element and it is the heaviest known halogen. Its chemical

properties are believed to be similar to those of iodine. Is has been little researched because

all its isotopes have short half lives. All that is known about the element has been estimated

from knowing its position in the periodic table below iodine and by studying its chemistry in

extreme diluted solutions.

Applications:Astatine is never encountered outside nuclear facilities or research laboratories.

Astatine in the environment:Total world production of astatine to date is estimated to be less than a millionth of a gram,

and virtually all of this has now decayed away.

Health effects of astatine:The total amount of astatine in the earth's crust at any particular time is less than 30 grams

and only a few micrograms have ever been artificially produced. This, together with its short

lifetime, leaves no reason for considering the effects of astatine on human health.

Astatine is studied in a few nuclear research laboratories where its high radioactivity requires

special handling techniques and precautions.

Astatine is a halogen and possibly accumulates in the thyroid like iodine. From a chemical

point of view, one can speculate that its toxicity would mimic that of iodine.

Environmental effects of astatine:Astatine does not occur to any significant extent in the biosphere and so normally never

presents a risk.

http://www.lenntech.com/Periodic-chart-elements/I-en.htm

C. TRANSITION METALS GROUP IVA

21 Sc Scandium Transition metal [Ar] 3d1 4s2

22 Ti Titanium Transition metal [Ar] 3d2 4s2

23 V Vanadium Transition metal [Ar] 3d3 4s2

24 Cr Chromium Transition metal [Ar] 3d5 4s1 (*)

25 Mn Manganese Transition metal [Ar] 3d5 4s2

26 Fe Iron Transition metal [Ar] 3d6 4s2

27 Co Cobalt Transition metal [Ar] 3d7 4s2

28 Ni Nickel Transition metal [Ar] 3d9 4s1 (*)

29 Cu Copper Transition metal [Ar] 3d10 4s1 (*)

30 Zn Zinc Transition metal [Ar] 3d10 4s2

Scandium - Sc

Atomic number: 21



Density: 3.0 g.cm-3 at 20°C

Melting point: 1541 °C

Boiling point: 2836 °C


Ionic radius: 0.083 nm (+3)

Isotopes: 7Electronic shell: [ Ar ] 3d1 4s2



Energy of third ionistion: 2389 kJ.mol -1

Energy of fourth ionization: 7089 kJ.mol -1

Discovered by: Lars Nilson in 1879

Description:Scandium is a soft, silvery transition element which occurs in rare minerals from Scandinavia.

It develops a slightly yellowish or pinkish cast when exposed to air. Scandium tarnished in air

and burn easily, once it has been ignited. It reacts with water to form hydrogen gas and will

dissolve in many acids. Pure scandium is produced by heating scandium fluoride (ScF3) with

calcium metal.

Applications:Scandium is one of the rare chemicals, that can be found in houses in equipment such as

colour televisions, fluorescent lamps, energy-saving lamps and glasses. The use of scandium

is still growing, due to the fact that it is suited to produce catalysers and to polish glass.

The main application by volume is in aluminium-scandium alloys for the aerospace industry

and for sports equipment (bikes, baseball bats, etc.) which rely on high performance

materials. It has been shown to reduce solidification cracking during the welding of high

strength aluminium alloys.

Scandium in the environment:Scandium can rarely be found in nature, as it occurs in very small amounts. Scandium is

usually found only in two different kinds of ores. Thortveitite is the primary source of

scandium with uranium mill tailings by-products also being an important source. World

production amount to only 50 kg per year. There is no estimate of how much is potentially

available. Scandium is only the 50th most abundant element on hearth, it is distributed

widely, occurring in trace quantities in over 800 minerals. The blue color of the aquamarine

variety of beryl is thought to be caused by scandium.

Only about 3% of plants that were analysed for scandium shows its presence, and even those

amounts were tiny, with vegetable having only 5 ppb although grass has 70 ppb.

Health effects of scandium:Scandium has no biological role. Only trace amounts reach the food chain, so the average

person's daily intake is less than 0.1 microgram.

Scadium is not toxic, although there have been suggestions that some of its compounds might

be cancerogenic. Scandium is mostly dangerous in the working environment, due to the fact

that damps and gasses can be inhaled with air. This can cause lung embolisms, especially

during long-term exposure. Scandium can be a threat to the liver when it accumulates in the

human body.

Effects of scandium on the environment:Scandium is dumped in the environment in many different places, mainly by petrol-producing

industries. It can also enter the environment when household equipment is thrown away.

Scandium will gradually accumulate in soils and water soils and this will eventually lead to

increasing concentrations in humans, animals and soil particles. With water animals scandium

causes damage to cell membranes, which has several negative influences on reproduction and

on the functions of the nervous system.

Titanium - Ti

Atomic number: 22







Ionic radius: 0.09 nm (+2) ; 0.068 nm (+4)






Discovered by: William Gregor in 1791

Description:Chemical element, Ti, atomic number 22 and atomic weight 47.90. Its chemical behavior

shows many similarities with that or silica and zirconium, as an element belonging to the first

transition group. Its chemistry in aqueous solution, especially in the lower oxidation states,

has some similarities with that of chrome and vanadium. Titanium is a transition metal light

with a white-silvery-metallic colour. It is stong, lustrous, corrosion-resistant. Pure titanium is

not soluble in water but is soluble in concentrated acids. This metal forms a passive but

protective oxide coating (leading to corrosion-resistance) when exposed to elevated

temperatures in air but at room temperatures it resists tarnishing.

The main oxidation state is 4+, although the states 3+ and 2+ are also known, but are less

stable. This element burns in the air when it’s heated up to obtain the dioxide, TiO2, and when

it is combined with halogens. It reduces the water vapor to form the dioxide and hydrogen,

and it reacts in a similar way with hot concentrated acids, although it forms trichloride with

chlorhydric acid. The metal absorbs hydrogen to give TiH2, and forms the nitride, TiN, and

the carbide, TiC. Other known compounds are the sulfur TiS2, as well as the lowest oxides,

Ti2O3 and TiO, and the sulfurs Ti2S3 and TiS. Salts are known in the three oxidation states.

Applications:The titanium dioxide is extensively used as a white pigment in outside paintings for being

chemically inert, for its great coating power, its opacity to UV light damage and its

autocleaning capacity. The dioxide was also used once as a bleaching and opicifying agent in

porcelain enamels, giving them a final touch of great brightness, hardness and acid resistance.

A typical lipstick contais 10% titanium.

Titaium alloys are characterized by very high tensile strength even at high temperatures, light

weight, high corrosion resistance, and ability to withstand extreme temperatures. ue to these

properties they are principally used in aircraft, pipes for power plants, armour plating, naval

ships, spacecraft and missiles. Titanium is as strong as steel but 45% lighter.

In medicine titanium is used to make hip and knee replacements, pace-makers, bone-plates

and screws and cranial plates for skull fractures. It has also been used to attach false theet.

The alkaline earth titanates have some remarkable properties. The level of dielectric constants

varies from 13 for the MgTiO3, to various milliards for solid solutions of SrTiO3 in BaTiO3.

The barium titanate also has a dielectric constant of 10.000 close to 120ºC, which is its Curie

point; it has low dielectric histeresis. The ceramic transductors that contain barium titanate are

favorably compared with Rochelle salt in terms of thermal stability and with quartz in terms

of the strength of the effect and the capacity to form the ceramics in various forms. The

compound has bee used as ultrasonic vibrations generator and as a sound detector.

Titanium in the environment:Althoug it is not found unbound to other elements in nature, titamuim is the ninth most

abundant element in the Earth's crust (0.63% by mass) and is present in most igneous rocks

and in sediments derived from them. Important titanium minerals are rutile, brookite, anatase,

illmenite, and titanite. The chief mined ore, ilmenite, occurs as vast deposits of sand in

Western Australia, Norway, Canada and Ukraine. Large deposits of rutile in North America

and South Africa also contribute significantly to the world supply of titanium. World

production of the metal is about 90.000 tonnes per year, and that of titanium dioxide is 4.3

million tonnes per year.

The titanium dioxide, TiO2, is commonly found in a black or brownish form known as rutile.

The natural forms that are less frequently found in nature are the anatasite and the brooquite.

Both the pure rutile and the pure anatasite are white. The black basic oxide, FeTiO3, is found

in the natural form as the natural mineral called ilmenite; this is the main commercial source

of titanium.

Health effects of titanium:There is no known biological role for titanium. There is a detectable amount of titanium in the

human body and it has been hestimated that we take in about 0.8 mg/day, but most passes

through us without being adsorbed. It is not a poisoun metal and the human body can tolerate

titanium in large dosis. Elemental titanium and titanium dioxide is of a low order of toxicity.

Laboratory animals (rats) exposed to titanium dioxide via inhalation have developed small-

localized areas of dark-colored dust deposits in the lungs. Excessive exposure in humans may

result in slight changes in the lungs.

Effects of overexposure to titanium powder: Dust inhalation may cause tightness and pain in

chest, coughing, and difficulty in breathing. Contact with skin or eyes may cause irritation.

Routes of entry: Inhalation, skin contact, eye contact.

Carcinogenicity: The International Agency for Research on Cancer (IARC) has listed

titanium dioxide within Group 3 (The agent is not classifiable as to its carcinogenicity to

humans.)

Environmental effects of titanium:Low toxicity. When in a metallic powdered form, titanium metal poses a significant fire

hazard and, when heated in air, an explosion hazard.

No environmental effects have been reported

Vanadium - V

Atomic number: 23







Ionic radius: 0.074 nm (+3) ; 0.059 (+5)






Discovered by: Nils Sefstrom in 1830

Description:Vanadium is a rare, soft, ductile gray-white element found combined in certain minerals and

used mainly to produce certain alloys. Vanadium resists corrosion due to a protective film of

oxide on the surface. Common oxidation states of vanadium include +2, +3, +4 and +5.

Applications:Most of the vanadium (about 80%) produced is used as ferrovanadium or as a steel additive.

Mixed with aluminium in titanium alloys is used in jet engines and high speed air-frames, and

steel alloys are used in axles, crankshafts, gears and other critical components. Vanadium

alloys are also used in nuclear reactors because vanadium has low neutron-adsorption abilities

and it doesn not deform in creeping under high temperatures.

Vanadium oxide (V2O5) is used as a catalyst in manufacturing sulfuric acid and maleic

anhydride and in making ceramics. It is added to glass to produce green or blue tint. Glass

coated with vanadium dioxide (VO2) can block infrared radiation at some specific

temperature.

Vanadium in the environment:Vanadium is never found unbound in nature. Vanadium occurs in about 65 different minerals

among which are patronite, vanadinite, carnotite and bauxite. Vanadium occurs in carbon

containing deposits such as crude oil, coal, oil shale and tar sands.

Various vanadium ores are known but none is mined as such for the metal, which is generally

obtained as a byproducts of other ores. The largest resources of vanadium are to be found in

South Africa and in Russia. World production of vanadium ore is around 45.000 tonnes a

year. Production of the metal itself comes to about 7000 tonnes per year.

Watering is an important way in which vanadium is redistributed around the environment

because venedates are generally very soluble.

Vanadium is abundant in most soils, in variable amounts, and it is taken up by plants at levels

that reflect its availability.

In biology, a vanadium atom is an essential component of some enzymes, particularly the

vanadium nitrogenase used by some nitrogen-fixing microorganisms.

Health effects of vanadium:Vanadium compounds are not regarded as serious hazard, however, workers exposed to

vanadium peroxide dust were found to suffer severe eye, nose and throat irritation.

The uptake of vanadium by humans mainly takes place through foodstuffs, such as

buckwheat, soya beans, olive oil, sunflower oil, apples and eggs.Vanadium can have a

number of effects on human health, when the uptake is too high. When vanadium uptake

takes places through air it can cause bronchitis and pneumonia. The acute effects of vanadium

are irritation of lungs, throat, eyes and nasal cavities. Other health effects of vanadium uptake

are:

- Cardiac and vascular disease

- Inflammation of stomach and intestines

- Damage to the nervous system

- Bleeding of livers and kidneys

- Skin rashes

- Severe trembling and paralyses

- Nose bleeds and throat pains

- Weakening

- Sickness and headaches

- Dizziness

- Behavioural changes

The health hazards associated with exposure to vanadium are dependent on its oxidation state.

This product contains elemental vanadium. Elemental vanadium could be oxidized to

vanadium pentoxide during welding. The pentoxide form is more toxic than the elemental

form. Chronic exposure to vanadium pentoxide dust and fumes may cause severe irritation of

the eyes, skin, upper respiratory tract, persistent inflammations of the trachea and bronchi,

pulmonary edema, and systemic poisoning. Signs and symptoms of overexposure include;

conjunctivitis, nasopharyngitis, cough, labored breathing, rapid heart beat, lung changes,

chronic bronchitis, skin pallor, greenish-black tongue and an allergic skin rash.

Effects of vanadium on the environment:Vanadium can be found in the environment in algae, plants, invertebrates, fishes and many

other species. In mussels and crabs vanadium strongly bioaccumulates, which can lead to

concentrations of about 105 to 106 times greater than the concentrations that are found in

seawater. Vanadium causes the inhibition of certain enzymes with animals, which has several

neurological effects. Next to the neurological effects vanadium can cause breathing disorders,

paralyses and negative effects on the liver and kidneys. Laboratory tests with test animals

have shown, that vanadium can cause harm to the reproductive system of male animals, and

that it accumulates in the female placenta. Vanadium can cause DNA alteration in some

cases, but it cannot cause cancer with animals.

Chromium - Cr

Atomic number: 24


Electronegativity: 1.6










Standard potential: - 0.71 V (Cr3+ / Cr )

Discovered by: Vaughlin in 1797

Description:Chromium is a lustrous, brittle, hard metal. Its colour is silver-gray and it can be highly

polished. It does not tarnish in air, when heated it borns and forms the green chromic oxide.

Chromium is unstable in oxygen, it immediately produces a thin oxide layer that is

impermeable to oxygen and protects the metal below.

Applications:Chromium main uses are in alloys such as stainless steel, in chrome plating and in metal

ceramics. Chromium plating was once widely used to give steel a polished silvery mirror

coating. Chromium is used in metallurgy to impart corrosion resistance and a shiny finish; as

dyes and paints, its salts colour glass an emerald green and it is used to produce synthetic

rubies; as a catalyst in dyeing and in the tanning of leather; to make molds for the firing of

bricks. Chromium (IV) oxide (CrO2) is used to manufacture magnetic tape.

Chromium in the environment:Chromium is mined as chromite (FeCr2O4) ore. Chromium ores are mined today in South

Africa, Zimbabwe, Finland, India, Kazakihstan and the Philippines. A total of 14 million

tonnes of chromite ore is extracted. Reserves are hestimated to be of the order of 1 billion

tonnes with unexploited deposits in Greenland, Canada e USA.

Health effects of chromium:People can be exposed to chromium through breathing, eating or drinking and through skin

contact with chromium or chromium compounds. The level of chromium in air and water is

generally low. In drinking water the level of chromium is usually low as well, but

contaminated well water may contain the dangerous chromium(IV); hexavalent chromium. For

most people eating food that contains chromium(III) is the main route of chromium uptake, as

chromium(III) occurs naturally in many vegetables, fruits, meats, yeasts and grains. Various

ways of food preparation and storage may alter the chromium contents of food. When food in

stores in steel tanks or cans chromium concentrations may rise.

Chromium(III) is an essential nutrient for humans and shortages may cause heart conditions,

disruptions of metabolisms and diabetes. But the uptake of too much chromium(III) can cause

health effects as well, for instance skin rashes. Chromium(VI) is a danger to human health,

mainly for people who work in the steel and textile industry. People who smoke tobacco also

have a higher chance of exposure to chromium .Chromium(VI) is known to cause various

health effects. When it is a compound in leather products, it can cause allergic reactions, such

as skin rash. After breathing it in chromium(VI) can cause nose irritations and nosebleeds.

Other health problems that are caused by chromium(VI) are:

- Skin rashes

- Upset stomachs and ulcers

- Respiratory problems

- Weakened immune systems

- Kidney and liver damage

- Alteration of genetic material

- Lung cancer

- Death

The health hazards associated with exposure to chromium are dependent on its oxidation state.

The metal form (chromium as it exists in this product) is of low toxicity. The hexavalent form

is toxic. Adverse effects of the hexavalent form on the skin may include ulcerations,

dermatitis, and allergic skin reactions. Inhalation of hexavalent chromium compounds can

result in ulceration and perforation of the mucous membranes of the nasal septum, irritation of

the pharynx and larynx, asthmatic bronchitis, bronchospasms and edema. Respiratory

symptoms may include coughing and wheezing, shortness of breath, and nasal itch.

Carcinogenicity- Chromium and most trivalent chromium compounds have been listed by the

National Toxicology Program (NTP) as having inadequate evidence for carcinogenicity in

experimental animals. According to NTP, there is sufficient evidence for carcinogenicity in

experimental animals for the following hexavalent chromium compounds; calcium chromate,

chromium trioxide, lead chromate, strontium chromate,and zinc chromate. International

Agency for Research on Cancer (IARC) has listed chromium metal and its trivalent

compounds within Group 3 (The agent is not classifiable as to its carcinogenicity to humans.)

Chromium is not regulated as a carcinogen by OSHA (29 CFR 1910 Subpart Z). ACGIH has

classified chromium metal and trivalent chromium compounds as A4,not classifiable as a

human carcinogen.

Manganese - Mn

Atomic number: 25











Standard potential: - 1.05 V ( Mn2+/ Mn )

Discovered: Johann Gahn in 1774

Description:Manganese is a pinkinsh-gray, chemically active element. It is a hard metal and is very brittle.

It is hard to melt, but easily oxidized. Manganese is reactive when pure, and as a powder it

will burn in oxygen, it reacts with water (it rusts like iron) and dissolves in dilute acids.

Applications:Manganese is essential to iron and steel production. At present steel making accounts 85% to

90% of the total demand, most of the total demand. Manganese is a key component of low-

cost stainless steel formulations and certain widely used alumimum alloys. Manganese

dioxide is also used as a catalyst. Manganese is used to decolorize glass and make violet

coloured glass. Potassium permanganate is a potent oxidizer and used as a disinfectant. Other

compound that find application are Manganese oxide (MnO) and manganese carbonate

(MnCO3): the first goes into fertilizers and ceramics, the second is the starting material for

making other manganese compounds.

Manganese in the environment:Manganese is one of the most abundant metals in soils, where it occurs as oxides and

hydroxides, and it cycles through its various oxidation states. Manganese occurs principally

as pyrolusite (MnO2), and to a lesser extent as rhodochrosite (MnCO3). More than 25 million

tonnes are mined every year, representing 5 million tons of the metal, and reserves are

estimated to exeed 3 billion tonnes of the metal. The main mining areas for manganese ores

are South Africa, Russia, Ukraine, Georgia, Gabon and Australia.

Manganese is an essential element for all species. Some organisms, such as diatoms, molluscs

and sponges, accumulate manganese. Fish can have up to 5 ppm and mammals up to 3 ppm in

their tissue, although normally they have around 1 ppm.

Health effects of manganese:Manganese is a very common compound that can be found everywhere on earth. Manganese is

one out of three toxic essential trace elements, which means that it is not only necessary for

humans to survive, but it is also toxic when too high concentrations are present in a human

body. When people do not live up to the recommended daily allowances their health will

decrease. But when the uptake is too high health problems will also occur. The uptake of

manganese by humans mainly takes place through food, such as spinach, tea and herbs. The

foodstuffs that contain the highest concentrations are grains and rice, soya beans, eggs, nuts,

olive oil, green beans and oysters. After absorption in the human body manganese will be

transported through the blood to the liver, the kidneys, the pancreas and the endocrine glands.

Manganese effects occur mainly in the respiratory tract and in the brains. Symptoms of

manganese poisoning are hallucinations, forgetfulness and nerve damage. Manganese can also

cause Parkinson, lung embolism and bronchitis. When men are exposed to manganese for a

longer period of time they may become impotent. A syndrome that is caused by manganese

has symptoms such as schizophrenia, dullness, weak muscles, headaches and

insomnia.Because manganese is an essential element for human health shortages of

manganese can also cause health effects. These are the following effects:

- Fatness

- Glucose intolerance

- Blood clotting

- Skin problems

- Lowered cholesterol levels

- Skeleton disorders

- Birth defects

- Changes of hair colour

- Neurological symptoms

Chronic Manganese poisoning may result from prolonged inhalation of dust and fume. The

central nervous system is the chief site of damage from the disease, which may result in

permanent disability. Symptoms include languor, sleepiness, weakness, emotional

disturbances, spastic gait, recurring leg cramps, and paralysis. A high incidence of pneumonia

and other upper respiratory infections has been found in workers exposed to dust or fume of

Manganese compounds. Manganese compounds are experimental equivocal tumorigenic

agents.

Iron - Fe

Atomic number: 26






Vanderwaalsradius: 0.126 nm






Standard potential: - o.44 V (Fe2+/ Fe ) ; 0.77 V ( Fe3+/ Fe2+ )

Discovered by: The ancients

Description:Iron is a lustrous, ductile, malleable, silver-gray metal (group VIII of the periodic table). It is

known to exist in four distinct crystalline forms. Iron rusts in dump air, but not in dry air. It

dissolves readily in dilute acids. Iron is chemically active and forms two major series of

chemical compounds, the bivalent iron (II), or ferrous, compounds and the trivalent iron (III),

or ferric, compounds.

Applications:Iron is the most used of all the metals, including 95 % of all the metal tonnage produced

worldwide. Thanks to the combination of low cost and high strength it is indispensable. Its

applications go from food containers to family cars, from scredrivers to washing machines,

from cargo ships to paper staples.

Steel is the best known alloy of iron, and some of the forms that iron takes include: pig iron,

cast iron, carbon steel, wrought iron, alloy steels, iron oxides.

Iron in the environment:Iron is believed to be the tenth most abundant element in the universe. Iron is also the most

abundant (by mass, 34.6%) element making up the Earth; the concentration of iron in the

various layers of the Earth ranges from high at the inner core to about 5% in the outer crust.

Most of this iron is found in various iron oxides, such as the minerals hematite, magnetite,

and taconite. The earth's core is believed to consist largely of a metallic iron-nickel alloy.

Iron is essential to almost living things, from micro-organisms to humans.

World production of new iron is over 500 million tonnes a year, and recycled iron add other

300 million tonnes. Economically workable reserves of iron ores exceed 100 billion tonnes.

The main mining areas are China, Brazil, Australia, Russia and Ukraine, with sizeable

amounts mined in the USA, Canada, Venezuela, Sweeden and India.

Health effects of iron:Iron can be found in meat, whole meal products, potatoes and vegetables. The human body

absorbs iron in animal products faster than iron in plant products. Iron is an essential part of

hemoglobin; the red colouring agent of the blood that transports oxygen through our bodies.

Iron may cause conjunctivitis, choroiditis, and retinitis if it contacts and remains in the

tissues. Chronic inhalation of excessive concentrations of iron oxide fumes or dusts may

result in development of a benign pneumoconiosis, called siderosis, which is observable as an

x-ray change. No physical impairment of lung function has been associated with siderosis.

Inhalation of excessive concentrations of iron oxide may enhance the risk of lung cancer

development in workers exposed to pulmonary carcinogens. LD50 (oral, rat) =30 gm/kg.

(LD50: Lethal dose 50. Single dose of a substance that causes the death of 50% of an animal

population from exposure to the substance by any route other than inhalation. Usually

expressed as milligrams or grams of material per kilogram of animal weight (mg/kg or g/kg).)

A more common problem for humans is iron deficency, which leads to anaemia. A man needs

an average daily intake pf 7 mg of iron and a woman 11 mg; a normal diet will generally

provided all that is needed.

Environmental effects of iron:Iron (III)-O-arsenite, pentahydrate may be hazardous to the environment; special attention

should be given to plants, air and water. It is strongly advised not to let the chemical enter

into the environment because it persists in the environment.

Cobalt – Co

Atomic number: 27












Standard potential: - 0.28 V ( Co2+/ Co ) ; +1.84 V ( Co3+/ Co2+ )

Discovered by: George Brandt in 1737

Description:Cobalt is a hard ferromagnetic, silver-white, hard, lustrous, brittle element. It is a member of

group VIII of the periodic table. Like iron, it can be magnetized. It is similar

to iron and nickel in its physical properties. The element is active chemically, forming many

compounds. Cobalt is stable in air and unaffected by water, but is slowly attacked by dilute

acids.

ApplicationsCobalt is used in many alloys (superalloys for parts in gas turbine aircrafr engines, corrosion

resistant alloys, high-speed steels, cemented carbides), in magents and magnetic recording

media, as catalysts for the petroleum and chemical industries, as drying agents for paints and

inks. Cobalt blue is an important part of artists' palette and is used bu craft workers in

porcelain, pottery, stained glass, tiles and enamel jewellery. The radioactive isotopes, cobalt-

60, is used in medical treatment and also to irradiate food, in order to preserve the food and

protect the consumer.

Cobalt in the enviromentMost of the Earth's cobalt is in its core. Cobalt is of relatively low abundance in the Earth's

crust and in natural waters, from which it is precipitated as the highly insoluble cobalt sulfine

CoS.

Although the average level of cobalt in soils is 8 ppm, there are soils with as little as 0.1 ppm

and others with as much as 70 ppm. In the marine environment cobalt is needed by blue-green

algae (cyanobacteria) and other nitrogen fixing organisms. Cobalt is not found as a free metal

and is generally found in the form of ores. Cobalt is usually not mined alone, and tends to be

produced as a by-product of nickel and copper mining activities. The main ores of cobalt are

cobaltite, erythrite, glaucodot, and skutterudite. The world's major producers of cobalt are the

Democratic Republic of the Congo, mainland China, Zambia, Russia and Australia. It is also

found in Finland, Azerbaijan, and Kazakhstan.

World production is 17.000 tonnes per year.

Health effects of cobaltAs cobalt is widely dispersed in the environment humans may be exposed to it by breathing

air, drinking water and eating food that contains cobalt. Skin contact with soil or water that

contains cobalt may also enhance exposure. Cobalt is not often freely available in the

environment, but when cobalt particles are not bound to soil or sediment particles the uptake

by plants and animals is higher and accumulation in plants and animals may occur. Cobalt is

beneficial for humans because it is a part of vitamin B12, which is essential for human health.

Cobalt is used to treat anaemia with pregnant women, because it stimulates the production of

red blood cells. The total daily intake of cobalt is variable and may be as much as 1 mg, but

almost all will pass through the body unadsorbed, except that in vitamine B12. However, too

high concentrations of cobalt may damage human health. When we breathe in too high

concentrations of cobalt through air we experience lung effects, such as asthma and

pneumonia. This mainly occurs with people that work with cobalt. When plants grow on

contaminated soils they will accumulate very small particles of cobalt, especially in the parts

of the plant we eat, such as fruits and seeds. Soils near mining and melting facilities may

contain very high amounts of cobalt, so that the uptake by humans through eating plants can

cause health effects. Health effects that are a result of the uptake of high concentrations of

cobalt are:

- Vomiting and nausea

- Vision problems

- Heart problems

- Thyroid damage

Health effects may also be caused by radiation of radioactive cobalt isotopes. This can cause

sterility, hair loss, vomiting, bleeding, diarrhoea, coma and even death. This radiation is

sometimes used with cancer-patients to destroy tumors. These patients also suffer from hair

loss, diarrhea and vomiting.

Cobalt dust may cause an asthma-like disease with symptoms ranging from cough, shortness

of breath and dyspnea to decreased pulmonary function, nodular fibrosis, permanent disability,

and death. Exposure to cobalt may cause weight loss, dermatitis, and respiratory

hypersensitivity. LD 50 (oral, rat)- 6171 mg/kg. (LD50 = Lethal dose 50 = Single dose of a

substance that causes the death of 50% of an animal population from exposure to the substance

by any route other than inhalation. LD50 is usually expressed as milligrams or grams of

material per kilogram of animal weight (mg/kg or g/kg).)

Carcinogenicity- International Agency for Research on Cancer (IARC) haslisted cobalt and

cobalt compounds within group 2B (agents which are possibly carcinogenic to humans).

ACGIH has placed cobalt and inorganic compounds in category A3 (Experimental animal

carcinogen- the agent is carcinogenic in experimental animals at a relatively high dose, by

route(s), histologic type(s), or by mechanism(s) that are not considered relevant to worker

exposure.) Cobalt has been classified to be carcinogenic to experimental animals by the

Federal Republic of Germany.

Environmental effects of cobaltCobalt is an element that occurs naturally in the environment in air, water, soil, rocks, plants

and animals. It may also enter air and water and settle on land through wind-blown dust and

enter surface water through run-off when rainwater runs through soil and rock containing

cobalt. Humans add cobalt by releasing small amounts into the atmosphere from coal

combustion and mining, processing of cobalt-containing ores and the production and use of

cobalt chemicals. The radioactive isotopes of cobalt are not present in the environment

naturally, but they are released through nuclear power plant operations and nuclear accidents.

Because they have relatively short half-lives they are not particularly dangerous. Cobalt

cannot be destroyed once it has entered the environment. It may react with other particles or

adsorb on soil particles or water sediments. Cobalt will only mobilize under acidic conditions,

but ultimately most cobalt will end up in soils and sediments. Soils that contain very low

amounts of cobalt may grow plants that have a deficiency of cobalt. When animals graze on

these grounds they suffer from lack of cobalt, which is essential for them. On the other hand,

soils near mining and melting facilities may contain very high amounts of cobalt, so that the

uptake by animals through eating plants can cause health effects. Cobalt will accumulate in

plants and in the bodies of animals that eat these plants, but cobalt is not known to bio

magnify up the food chain. Because of this fruits, vegetables, fish and other animals we eat

will usually not contain very high amounts of cobalt.

Nickel - Ni

Atomic number: 28








Isotopes: 10

Electronic shell: [ Ar ] 3d8 4s2




Standard potential: - 0.25 V

Discovered by: Alex Constedt 1751

Description:Nickel is silvery-white. hard, malleable, and ductile metal. It is of the iron group and it takes

on a high polish. It is a fairly good conductor of heat and electricity. In its familiar

compounds nickel is bivalent, although it assumes other valences. It also forms a number of

complex compounds. Most nickel compounds are blue or green. Nickel dissolves slowly in

dilute acids but, like iron, becomes passive when treated with nitric acid. Finely divided

nickel adsorbs hydrogen.

ApplicationsThe major use of nickel is in the preparation of alloys. Nickel alloys are characterized by

strength, ductility, and resistance to corrosion and heat. About 65 % of the nickel consumed

in the Western World is used to make stainless steel, whose composition can vary but is

typically iron with around 18% chromium and 8% nickel. 12 % of all the nickel consumed

goes into super alloys. The remaining 23% of consumption is divided between alloy steels,

rechargeable batteries, catalysts and other chemicals, coinage, foundry products, and plating.

Nickel is easy to work and can be drawn into wire. It resist corrosion even at high

temperatures and for this reason it is used in gas turbines and rocket engines. Monel is an

alloy of nickel and copper (e.g. 70% nickel, 30% copper with traces of

iron, manganese and silicon), which is not only hard but can resist corrosion by sea water, so

that it is ideal for propeller shaft in boats and desalination plants.

Nickel in the environmentMost nickel on Earth is inaccessible because it is locked away in the planet's iron-nickel

molten core, which is 10 % nickel. The total amount of nickel dissolved in the sea has been

calculated to be around 8 billion tons. Organic matter has a strong ability to absorb the metal

which is why coal and oil contain considerable amounts. The nickel content in soil can be as

low as 0.2 ppm or as high as 450 ppm in some clay and loamy soils. The average is around 20

ppm. Nickel occurs in some beans where it is an essential component of some enzymes.

Another relatively rich source of nickel is tea which has 7.6 mg/kg of dried leaves.

Nickel occurs combined with sulphur in millerite, with arsenic in the mineral niccolite, and

with arsenic and sulphur in nickel glance. Most ores from which nickel is extracted are iron-

nickel sulphides, such as pentlandite. The metal is mined in Russia, Australia, New

Caledonia, Cuba, Canada and South Africa. Annual production exceeds 500.000 tons and

easily workable reserves will last at least 150 years.

Health effects of nickel

Nickel is a compound that occurs in the environment only at very low levels. Humans use

nickel for many different applications. The most common application of nickel is the use as

an ingredient of steal and other metal products. It can be found in common metal products

such as jewelry. Foodstuffs naturally contain small amounts of nickel. Chocolate and fats are

known to contain severely high quantities. Nickel uptake will boost when people eat large

quantities of vegetables from polluted soils. Plants are known to accumulate nickel and as a

result the nickel uptake from vegetables will be eminent. Smokers have a higher nickel uptake

through their lungs. Finally, nickel can be found in detergents. Humans may be exposed to

nickel by breathing air, drinking water, eating food or smoking cigarettes. Skin contact with

nickel-contaminated soil or water may also result in nickel exposure. In small quantities

nickel is essential, but when the uptake is too high it can be a danger to human health. An

uptake of too large quantities of nickel has the following consequences:

- Higher chances of development of lung cancer, nose cancer, larynx cancer and prostate

cancer

- Sickness and dizziness after exposure to nickel gas

- Lung embolism

- Respiratory failure

- Birth defects

- Asthma and chronic bronchitis

- Allergic reactions such as skin rashes, mainly from jewelry

- Heart disorders

Nickel fumes are respiratory irritants and may cause pneumonitis. Exposure to nickel and its

compounds may result in the development of a dermatitis known as “nickel itch” in sensitized

individuals. The first symptom is usually itching, which occurs up to 7 days before skin

eruption occurs. The primary skin eruption is erythematous, or follicular, which may be

followed by skin ulceration. Nickel sensitivity, once acquired, appears to persist indefinitely.

Carcinogenicity- Nickel and certain nickel compounds have been listed by the National

Toxicology Program (NTP) as being reasonably anticipated to be carcinogens. The

International Agency for Research on Cancer (IARC) has listed nickel compounds within

group 1 (there is sufficient evidence for carcinogenicity in humans) and nickel within group

2B (agents which are possibly carcinogenic to humans). OSHA does not regulate nickel as a

carcinogen. Nickel is on the ACGIH Notice of Intended Changes as a Category A1,

confirmed human carcinogen.

Effects of nickel on the environment

Nickel is released into the air by power plants and trash incinerators. It will than settle to the

ground or fall down after reactions with raindrops. It usually takes a long time for nickel to be

removed from air. Nickel can also end up in surface water when it is a part

of wastewater streams. The larger part of all nickel compounds that are released to the

environment will adsorb to sediment or soil particles and become immobile as a result. In

acidic ground however, nickel is bound to become more mobile and it will often rinse out to

the groundwater. There is not much information available on the effects of nickel upon

organisms other than humans. We do know that high nickel concentrations on sandy soils can

clearly damage plants and high nickel concentrations in surface waters can diminish the

growth rates of algae. Micro organisms can also suffer from growth decline due to the

presence of nickel, but they usually develop resistance to nickel after a while. For animals

nickel is an essential foodstuff in small amounts. But nickel is not only favorable as an

essential element; it can also be dangerous when the maximum tolerable amounts are

exceeded. This can cause various kinds of cancer on different sites within the bodies of

animals, mainly of those that live near refineries. Nickel is not known to accumulate in plants

or animals. As a result nickel will not bio magnify up the food chain.

Copper - Cu

Atomic number: 29








Isotopes: 6

Electronic shell: [ Ar ] 3d10 4s1



Standard potential: + 0.522 V ( Cu+/ Cu ) ; + 0.345 V (Cu2+/ Cu )

Discovered by: The ancients

Description:Copper is a reddish metal with a face-centered cubic crystalline structure. It reflects red and

orange light and absorbs other frequencies in the visible spectrum, due to its band structure,

so it as a nice reddish color. It is malleable, ductile, and an extremely good conductor of both

heat and electricity. It is softer than zinc and can be polished to a bright finish. It is found in

group Ib of the periodic table, together with silver and gold. Copper has low chemical

reactivity. In moist air it slowly forms a greenish surface film called patina; this coating

protects the metal from further attack.

ApplicationsMost copper is used for electrical equipment (60%); construction, such as roofing and

plumbing (20%); industrial machinery, such as heat exchangers (15%) and alloys (5%). The

main long established copper alloys are bronze, brass (a copper-zinc alloy), copper-tin-zinc,

which was strong enough to make guns and cannons, and was known as gun metal, copper

and nickel, known as cupronickel, which was the preferred metal for low-denomination coins.

Copper is ideal for electrical wiring because it is easily worked, can be drawn into fine wire

and has a high electrical conductivity.

Copper in the environmentCopper is a very common substance that occurs naturally in the environment and spreads

through the environment through natural phenomena. Humans widely use copper. For

instance it is applied in the industries and in agriculture. The production of copper has lifted

over the last decades. Due to this, copper quantities in the environment have increased.

The world's copper production is still rising. This basically means that more and more copper

ends up in the environment. Rivers are depositing sludge on their banks that is contaminated

with copper, due to the disposal of copper-containing wastewater. Copper enters the air,

mainly through release during the combustion of fossil fuels. Copper in air will remain there

for an eminent period of time, before it settles when it starts to rain. It will then end up mainly

in soils. As a result soils may also contain large quantities of copper after copper from the air

has settled. Copper can be released into the environment by both natural sources and human

activities. Examples of natural sources are wind-blown dust, decaying vegetation, forest fires

and sea spray. A few examples of human activities that contribute to copper release have

already been named. Other examples are mining, metal production, wood production and

phosphate fertilizer production. Because copper is released both naturally and through human

activity it is very widespread in the environment. Copper is often found near mines, industrial

settings, landfills and waste disposals. Most copper compounds will settle and be bound to

either water sediment or soil particles. Soluble copper compounds form the largest threat to

human health. Usually water-soluble copper compounds occur in the environment after

release through application in agriculture. World production of copper amounts to 12 million

tons a year and exploitable reserves are around 300 million tons, which are expected to last

for only another 25 years. About 2 million tons a year are reclaimed by recycling. Today

copper is mined as major deposits in Chile, Indonesia, USA, Australia and Canada, which

together account for around 80% of the world's copper. The main ore is a yellow copper-iron

sulfide called chalcopyrite (CuFeS2).

Zinc - Zn

Atomic number 30





Boiling point 907 °C


Ionic radius 0.074 nm (+2)

Isotopes 10

Electronic shell [ Ar ] 3d10 4s2

Energy of first ionisation 904.5 kJ.mol -1

Energy of second ionisation 1723 kJ.mol -1

Standard potential - 0.763 V

Discovered Andreas Marggraf in 1746

Description:Zinc is a lustrous bluish-white metal. It is found in group IIb of the periodic table. It is brittle

and crystalline at ordinary temperatures, but it becomes ductile and malleable when heated

between 110°C and 150°C. It is a fairly reactive metal that will combine with oxygen and

other non-metals, and will react with dilute acids to release hydrogen.

ApplicationsIt is used principally for galvanizing iron, more than 50% of metallic zinc goes into

galvanizing steel, but is also important in the preparation of certain alloys. It is used for the

negative plates in some electric batteries and for roofing and gutters in building construction.

Zinc is the primary metal used in making American pennies, is used in die casting in the

automobile industry. Zinc oxide is used as a white pigment in watercolours or paints, and as

an activator in the rubber industry. As a pigment zinc is used in plastics, cosmetics,

photocopier paper, wallpaper, printing inks etc, while in rubber production its role is to act as

a catalyst during manufacture and as a heat disperser in the final product. Zinc metal is

included in most single tablet, it is believed to possess anti-oxidant properties, which protect

against premature aging of the skin and muscles of the body.

Zinc in the environmentZinc is a very common substance that occurs naturally. Many foodstuffs contain certain

concentrations of zinc. Drinking water also contains certain amounts of zinc, which may be

higher when it is stored in metal tanks. Industrial sources or toxic waste sites may cause the

zinc amounts in drinking water to reach levels that can cause health problems.

Zinc occurs naturally in air, water and soil, but zinc concentrations are rising unnaturally, due

to addition of zinc through human activities. Most zinc is added during industrial activities,

such as mining, coal and waste combustion and steel processing. Some soils are heavily

contaminated with zinc, and these are to be found in areas where zinc has to be mined or

refined, or were sewage sludge from industrial areas has been used as fertilizer.

Zinc is the 23rd most abundant element in the Earth's crust. The dominant ore is zinc blende,

also known as sphalerite. Other important zinc ores are wurzite, smithsonite and

hemimorphite. The main zinc mining areas are Canada, Russia, Australia, USA and Peru'.

World production exceeds 7 million tonnes a year and commercially exploitable reserves

exceed 100 million tonnes. More than 30% of the world's need for zinc is met by recycling.

Effects of zinc on the EnvironmentThe world's zinc production is still rising. This basically means that more and more zinc ends

up in the environment.

Water is polluted with zinc, due to the presence of large quantities of zinc in the wastewater of

industrial plants. This wastewater is not purified satisfactory. One of the consequences is that

rivers are depositing zinc-polluted sludge on their banks. Zinc may also increase the acidity of

waters. Some fish can accumulate zinc in their bodies, when they live in zinc-contaminated

waterways. When zinc enters the bodies of these fish it is able to bio magnify up the food

chain. Large quantities of zinc can be found in soils. When the soils of farmland are polluted

with zinc, animals will absorb concentrations that are damaging to their health. Water-soluble

zinc that is located in soils can contaminate groundwater.

Zinc cannot only be a threat to cattle, but also to plant species. Plants often have a zinc uptake

that their systems cannot handle, due to the accumulation of zinc in soils. On zinc-rich soils

only a limited number of plants has a chance of survival. That is why there is not much plant

diversity near zinc-disposing factories. Due to the effects upon plants zinc is a serious threat to

the productions of farmlands. Despite of this zinc-containing manures are still applied. Finally,

zinc can interrupt the activity in soils, as it negatively influences the activity of microrganisms

and earthworms. The breakdown of organic matter may seriously slow down because of this.

Documents

Chemistry, Elements