SPI Liquid Nitrogen Dewars & Supplies

Our concern is with presenting good maintenance and use information to protect the life of your new liquid nitrogen dewar but at the same time, make sure that the user is also protecting their health and well-being and at all times, in ways that would prevent injuries.

How to take care of your liquid nitrogen dewar while taking care of your own health!

Warning

Use only liquid nitrogen (or liquid argon) in liquid nitrogen dewars supplied by SPI Supplies or equivalent. Do not ever use liquid air or liquid oxygen in these dewars because either of which could present a combustion hazard with some materials used in the construction of these dewars, or materials stored in them.

Introduction

The safe handling and use of liquid nitrogen in liquid nitrogen dewars or flasks is possible only by knowing the potential hazards and using common-sense procedures based on that knowledge. There are two important properties of liquid nitrogen that present potential hazards:

1.It is extremely cold. At atmospheric pressure, liquid nitrogen boils at -320° F/-196° C.

2.Very small amounts of liquid vaporize into large amounts of gas. One liter of liquid nitrogen becomes 24.6 ft³/0.7 m³ of gas.

The safety precautions as outlined must be followed to avoid potential injury or damage which could result from these two characteristics. Do not attempt to handle liquid nitrogen until you read and fully understand the potential hazards, their consequences, and the related safety precautions. Keep a print out of this webpage handy for ready reference and review.

Maintenance:
Keep the unit clean and dry at all times. Do not use strong alkaline or acid cleaners that could damage the finish and corrode the metal shell.

Note:
Because argon is an inert gas whose physical properties are very similar to those of nitrogen, the precautions and safe practices for the handling and use of liquid argon are the same as those for liquid nitrogen.

Handling Liquid Nitrogen

Contact of liquid nitrogen or any very cold gas with the skin or eyes may cause serious freezing (frostbite) injury.
Protect hands at all times when working with liquid nitrogen with SPI Cryo Gloves.

Handle liquid nitrogen carefully
The extremely low temperature can freeze human flesh very rapidly. When spilled on a surface the liquid tends to cover it completely and intimately, cooling a large area. The gas issuing from the liquid is also extremely cold. Delicate tissue, such as that of the eyes, can be damaged by an exposure to the cold gas which would be too brief to affect the skin of the hands or face.

Never allow any unprotected part of your body to touch objects cooled by liquid nitrogen.
Such objects may stick fast to the skin and tear the flesh when you attempt to free yourself. Use tongs, preferably with insulated handles, to withdraw objects immersed in the liquid, and handle the object carefully.

Wear protective clothing
Protect your eyes with a face shield or safety goggles (safety glasses without side shields do not give adequate protection). Always wear cryo gloves when handling anything that is, or may have been, in immediate contact with liquid nitrogen. The gloves should fit loosely, so that they can be thrown off quickly if liquid should splash into them. When handling liquid in open containers, it is advisable to wear high-top shoes. Trousers (which should be cuffless if possible) should be worn outside the shoes.

Any kind of canvas shoes should be avoided because a liquid nitrogen spill can be taken up by the canvas resulting in a far more severe burn, in fact that would occur if the feet were essentially open or bare! Now we don't advocate going bare foot when using liquid nitrogen, but we also don't think that the wearing of canvas shoes is a safe practice either.

Use only containers designed for low-temperature liquids
Cryogenic containers are specifically designed and made of materials that can withstand the rapid changes and extreme temperature differences encountered in working with liquid nitrogen. Even these special containers should be filled slowly to minimize the internal stresses that occur when any material is cooled. Excessive internal stresses can damage the container.

Do not ever cover or plug the entrance opening of any liquid nitrogen dewar. Do not use any stopper or other device that would interfere with venting of gas.

These cryogenic liquid containers are generally designed to operate with little or no internal pressure. Inadequate venting can result in excessive gas pressure which could damage or burst the container. Use only the loose-fitting necktube core supplied or one of the approved accessories for closing the necktube. Check the unit periodically to be sure that venting is not restricted by accumulated ice or frost.

Use proper transfer equipment
Use a phase separator or special filling funnel to prevent splashing and spilling when transferring liquid nitrogen into or from a dewar. The top of the funnel should be partly covered to reduce splashing. Use only small, easily handled dewars for pouring liquid. For the larger, heavier containers, use a cryogenic liquid withdrawal device to transfer liquid from one container to another. Be sure to follow instructions supplied with the withdrawal device. When liquid cylinders or other large storage containers are used for filling, follow the instructions supplied with those units and their accessories.

Do not overfill containers
Filling above the bottom of the necktube (or specified maximum level) can result in overflow and spillage of liquid when the necktube core or cover is placed in the opening.

Never use hollow rods or tubes as dipsticks
When a warm tube is inserted into liquid nitrogen, liquid will spout from the bottom of the tube due to gasification and rapid expansion of liquid inside the tube. Wooden or solid metal dipsticks are recommended; avoid using plastics that may become very brittle at cryogenic temperatures which then become prone to shatter like a fragile piece of glass.

Nitrogen gas can cause suffocation without warning. Store and use liquid nitrogen only in a well ventilated place.
As the liquid evaporates, the resulting gas tends to displace the normal air from the area. In closed areas, excessive amounts of nitrogen gas reduce the concentration of oxygen and can result in asphyxiation. Because nitrogen gas is colorless, odorless and tasteless, it cannot be detected by the human senses and will be breathed as if it were air. Breathing an atmosphere that contains less than 18 percent oxygen can cause dizziness and quickly result in unconsciousness and death.

Note:
The cloudy vapor that appears when liquid nitrogen is exposed to the air is condensed moisture, not the gas itself. The gas actually causing the condensation and freezing is completely invisible.

Never dispose of liquid nitrogen in confined areas or places where others may enter.

Disposal of liquid nitrogen should be done outdoors in a safe place. Pour the liquid slowly on gravel or bare earth where it can evaporate without causing damage. Do not pour the liquid on the pavement.

First Aid Notice
If a person seems to become dizzy or loses consciousness while working with liquid nitrogen, move to a well-ventilated area immediately. If breathing has stopped, apply artificial respiration. If breathing is difficult, give oxygen. Call a physician. Keep warm and at rest.

If exposed to liquid or cold gas, restore tissue to normal body temperature 98.6° F (37° C) as rapidly as possible, followed by protection of the injured tissue from further damage and infection. Remove or loosen clothing that may constrict blood circulation to the frozen area. Call a physician. Rapid warming of the affected part is best achieved by using water at 108° F/42° C). Under no circumstances should the water be over 112° F/44° C, nor should the frozen part be rubbed either before or after rewarming. The patient should neither smoke, nor drink alcohol.

Most liquid nitrogen burns are really bad cases of frostbite. We don't mean to belittle the harm that can come from frostbite, but at the same time, we wanted to keep the dangers associated with liquid nitrogen burns in perspective. Indeed, liquid nitrogen burns could be treated as frostbite.

Handling Liquid Nitrogen Dewars

Keep unit upright at all times except when pouring liquid from dewars specifically designed for that purpose.

Tipping the container or laying it on its side can cause spillage of liquid nitrogen. It may also damage the container and any materials stored in it. If tipping is anticipated, be sure to purchase a dewar that can be outfitted with a tipping stand.

Rough handling can cause serious damage to dewars and refrigerators.
Dropping the container, allowing it to fall over on its side, or subjecting it to sharp impact or severe vibration can result in partial or complete loss of vacuum. To protect the vacuum insulation system, handle containers carefully. Do not "walk", roll or drag these units across a floor. Use a dolly or handcart when moving containers, especially the larger portable refrigerators. Large units are heavy enough to cause personal injury or damage to equipment if proper lifting and handling techniques are not used.

When transporting a liquid nitrogen dewar, maintain adequate ventilation and protect the unit from damage.
Do not place these units in closed vehicles where the nitrogen gas that is continuously vented from unit can accumulate. Prevent spillage of liquids and damage to unit by securing it in the upright position so that it cannot be tipped over. Protect the unit from sever jolting and impact that could cause damage, especially to the vacuum seal.

Keep the unit clean and dry
Do not store it in wet, dirty areas. Moisture, animal waste, chemicals, strong cleaning agents and other substances which could promote corrosion should be removed promptly. Use water or mild detergent for cleaning and dry the surface thoroughly. Do not use strong alkaline or acid cleaners that could damage the finish and corrode the metal shell.

Protect Dewar Contents

Materials stored in a liquid nitrogen dewar with a wide mouth are protected by the extremely low temperature of the liquid nitrogen or the gas that issues from the evaporating liquid nitrogen. When all of the liquid nitrogen has evaporated, the temperature inside the unit will rise slowly to ambient. The rate at which the liquid nitrogen will evaporate depends upon the pattern of container use and the age and condition of the container. Evaporation increases as insulation efficiency deteriorates with age and rough handling. Opening and closing to insert and remove materials and moving the unit will also increase the evaporation rate.

To protect valuable material stored in a liquid nitrogen refrigerator:

Check the liquid level in unit frequently
Great damage could result to laboratory equipment that requires constant cooling to protect some critical part of the equipment such as the Si (Li) detector on a modern EDS system. Or important experiments could be delayed, or critical samples spoiled, if one unexpectedly ran out of liquid nitrogen and then could not perform their cryo ultramicrotomy. So it is of the greatest importance to check the liquid nitrogen level constantly in order to anticipate any such kinds of problems that could arise.

Condensed moisture or frost on the outer shell of a refrigerator and abnormally rapid evaporation of the liquid nitrogen are indications of vacuum loss.
If vacuum loss is evident or suspected, start thinking immediately about the procurement of a replacement dewar. It is just not cost effective to continue to use a dewar with a bad vacuum and waste valuable liquid nitrogen in the process. There is also the safety issue of excessive boil-off in an enclosed area that is not large enough to "absorb" the higher rate of nitrogen boil off.

Nitrogen Safety

Rapid release of nitrogen gas into an enclosed space can displace oxygen, and therefore represents an asphyxiation hazard. This may happen with few warning symptoms, since the human carotid body is a relatively slow and a poor low-oxygen (hypoxia) sensing system. An example occurred shortly before the launch of the first Space Shuttle mission in 1981, when two technicians lost consciousness (and one of them died) after they walked into a space located in the Shuttle's Mobile Launcher Platform that was pressurized with pure nitrogen as a precaution against fire. The technicians would have been able to exit the room if they had experienced early symptoms from nitrogen-breathing.

When inhaled at high partial pressures (more than about 4 bar, encountered at depths below about 30 m in scuba diving) nitrogen begins to act as an anesthetic agent. It can cause nitrogen narcosis, a temporary semi-anesthetized state of mental impairment similar to that caused by nitrous oxide.

Nitrogen also dissolves in the bloodstream and body fats. Rapid decompression (particularly in the case of divers ascending too quickly, or astronauts decompressing too quickly from cabin pressure to spacesuit pressure) can lead to a potentially fatal condition called decompression sickness (formerly known as caisson sickness or more commonly, the "bends"), when nitrogen bubbles form in the bloodstream, nerves, joints, and other sensitive or vital areas.[33][34] Other "inert" gases (those gases other than carbon dioxide and oxygen) cause the same effects from bubbles composed of them, so replacement of nitrogen in breathing gases may prevent nitrogen narcosis, but does not prevent decompression sickness.[35]

Direct skin contact with liquid nitrogen will eventually cause severe frostbite (cryogenic "burns"). This may happen almost instantly on contact, or after a second or more, depending on the form of liquid nitrogen. Bulk liquid nitrogen causes less rapid freezing than a spray of nitrogen mist (such as is used to freeze certain skin growths in the practice of dermatology). The extra surface area provided by nitrogen-soaked materials is also important, with soaked clothing or cotton causing far more rapid damage than a spill of direct liquid to skin. Full "contact" between naked skin and large collected-droplets or pools of liquid nitrogen may be prevented for second or two, by a layer of insulating gas from the Leidenfrost effect. This may give the skin a second of protection from nitrogen bulk liquid. However, liquid nitrogen applied to skin in mists, and on fabrics, bypasses this effect, and causes local frostbite immediately.

Nitrogen Biological Role

Nitrogen is an essential building block of amino and nucleic acids, essential to life on Earth.

Elemental nitrogen in the atmosphere cannot be used directly by either plants or animals, and must be converted to a reduced (or 'fixed') state in order to be useful for higher plants and animals. Precipitation often contains substantial quantities of ammonium and nitrate, thought to result from nitrogen fixation by lightning and other atmospheric electric phenomena.^[24] This was first proposed by Liebig in 1827 and later confirmed.^[24] However, because ammonium is preferentially retained by the forest canopy relative to atmospheric nitrate, most fixed nitrogen reaches the soil surface under trees as nitrate. Soil nitrate is preferentially assimilated by tree roots relative to soil ammonium^{[citation needed]}.

Specific bacteria (e.g. Rhizobium trifolium) possess nitrogenase enzymes which can fix atmospheric nitrogen (see nitrogen fixation) into a form (ammonium ion) that is chemically useful to higher organisms. This process requires a large amount of energy and anoxic conditions. Such bacteria may live freely in soil (e.g. Azotobacter) but normally exist in a symbiotic relationship in the root nodules of leguminous plants (e.g. clover, Trifolium, or soybean plant, Glycine max). Nitrogen-fixing bacteria are also symbiotic with a number of unrelated plant species such as alders (Alnus) spp., lichens, Casuarina, Myrica, liverworts, and Gunnera.^[25]

As part of the symbiotic relationship, the plant converts the 'fixed' ammonium ion to nitrogen oxides and amino acids to form proteins and other molecules, (e.g. alkaloids). In return for the 'fixed' nitrogen, the plant secretes sugars to the symbiotic bacteria.^[25] Legumes maintain an anaerobic (oxygen free) environment for their nitrogen-fixing bacteria.

Plants are able to assimilate nitrogen directly in the form of nitrates which may be present in soil from natural mineral deposits, artificial fertilizers, animal waste, or organic decay (as the product of bacteria, but not bacteria specifically associated with the plant). Nitrates absorbed in this fashion are converted to nitrites by the enzyme nitrate reductase, and then converted to ammonia by another enzyme called nitrite reductase.

Nitrogen compounds are basic building blocks in animal biology as well. Animals use nitrogen-containing amino acids from plant sources, as starting materials for all nitrogen-compound animal biochemistry, including the manufacture of proteins and nucleic acids. Plant-feeding insects are dependent on nitrogen in their diet, such that varying the amount of nitrogen fertilizer applied to a plant can affect the reproduction rate of insects feeding on fertilized plants.^[26]

Soluble nitrate is an important limiting factor in the growth of certain bacteria in ocean waters.^[27] In many places in the world, artificial fertilizers applied to crop-lands to increase yields result in run-off delivery of soluble nitrogen to oceans at river mouths. This process can result in eutrophication of the water, as nitrogen-driven bacterial growth depletes water oxygen to the point that all higher organisms die. Well-known "dead zone" areas in the U.S. Gulf Coast and the Black Sea are due to this important polluting process.

Many saltwater fish manufacture large amounts of trimethylamine oxide to protect them from the high osmotic effects of their environment (conversion of this compound to dimethylamine is responsible for the early odor in unfresh saltwater fish.^[28] In animals, free radical nitric oxide (NO) (derived from an amino acid), serves as an important regulatory molecule for circulation.

Animal metabolism of NO results in production of nitrite. Animal metabolism of nitrogen in proteins generally results in excretion of urea, while animal metabolism of nucleic acids results in excretion of urea and uric acid. The characteristic odor of animal flesh decay is caused by the creation of long-chain, nitrogen-containing amines, such as putrescine and cadaverine which are (respectively) breakdown products of the amino acids ornithine and lysine in decaying proteins.

Decay of organisms and their waste products may produce small amounts of nitrate, but most decay eventually returns nitrogen content to the atmosphere, as molecular nitrogen. The circulation of nitrogen from atmosphere, to organic compounds, then back to the atmosphere, is referred to as the nitrogen cycle.

Nitrogen Compound

The main neutral hydride of nitrogen is ammonia (NH₃), although hydrazine (N₂H₄) is also commonly used. Ammonia is more basic than water by 6 orders of magnitude. In solution ammonia forms the ammonium ion (NH+4). Liquid ammonia (boiling point 240 K) is amphiprotic (displaying either Brønsted-Lowry acidic or basic character) and forms ammonium and the less common amide ions (NH−2); both amides and nitride (N³⁻) salts are known, but decompose in water. Singly, doubly, triply and quadruply substituted alkyl compounds of ammonia are called amines (four substitutions, to form commercially and biologically important quaternary amines, results in a positively charged nitrogen, and thus a water-soluble, or at least amphiphilic, compound). Larger chains, rings and structures of nitrogen hydrides are also known, but are generally unstable.

Other classes of nitrogen anions (negatively charged ions) are the poisonous azides (N−3), which are linear and isoelectronic to carbon dioxide, but which bind to important iron-containing enzymes in the body in a manner more resembling cyanide. Another molecule of the same structure is the colorless and relatively inert anesthetic gas Nitrous oxide (dinitrogen monoxide, N₂O), also known as laughing gas. This is one of a variety of nitrogen oxides that form a family often abbreviated as NOx. Nitric oxide (nitrogen monoxide, NO), is a natural free radical used in signal transduction in both plants and animals, for example in vasodilation by causing the smooth muscle of blood vessels to relax. The reddish and poisonous nitrogen dioxide NO₂ contains an unpaired electron and is an important component of smog. Nitrogen molecules containing unpaired electrons show an understandable tendency to dimerize (thus pairing the electrons), and are generally highly reactive. The corresponding acids are nitrous HNO₂ and nitric acid HNO₃, with the corresponding salts called nitrites and nitrates.

The higher oxides dinitrogen trioxide N₂O₃, dinitrogen tetroxide N₂O₄ and dinitrogen pentoxide N₂O₅, are unstable and explosive, a consequence of the chemical stability of N₂. Nearly every hypergolic rocket engine uses N₂O₄ as the oxidizer; their fuels, various forms of hydrazine, are also nitrogen compounds. These engines are extensively used on spacecraft such as the space shuttle and those of the Apollo Program because their propellants are liquids at room temperature and ignition occurs on contact without an ignition system, allowing many precisely controlled burns. Some launch vehicles, such as the Titan II and Ariane 1 through 4 also use hypergolic fuels, although the trend is away from such engines for cost and safety reasons. N₂O₄ is an intermediate in the manufacture of nitric acid HNO₃, one of the few acids stronger than hydronium and a fairly strong oxidizing agent.

Nitrogen is notable for the range of explosively unstable compounds that it can produce. Nitrogen triiodide NI₃ is an extremely sensitive contact explosive. Nitrocellulose, produced by nitration of cellulose with nitric acid, is also known as guncotton. Nitroglycerin, made by nitration of glycerin, is the dangerously unstable explosive ingredient of dynamite. The comparatively stable, but more powerful explosive trinitrotoluene (TNT) is the standard explosive against which the power of nuclear explosions are measured.

Nitrogen can also be found in organic compounds. Common nitrogen functional groups include: amines, amides, nitro groups, imines, and enamines. The amount of nitrogen in a chemical substance can be determined by the Kjeldahl method.

Nitrogen Property

Nitrogen is a nonmetal, with an electronegativity of 3.04. It has five electrons in its outer shell and is therefore trivalent in most compounds. The triple bond in molecular nitrogen (N₂) is the strongest. The resulting difficulty of converting N₂ into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N₂, have dominated the role of nitrogen in both nature and human economic activities.

At atmospheric pressure molecular nitrogen condenses (liquefies) at 77 K (−195.8 °C) and freezes at 63 K (−210.0 °C) into the beta hexagonal close-packed crystal allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the cubic crystal allotropic form (called the alpha-phase). Liquid nitrogen, a fluid resembling water in appearance, but with 80.8% of the density (the density of liquid nitrogen at its boiling point is 0.808 g/mL), is a common cryogen.

Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like N₃ and N₄.^[4] Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced using a diamond anvil cell, nitrogen polymerizes into the single-bonded cubic gauche crystal structure. This structure is similar to that of diamond, and both have extremely strong covalent bonds. N₄ is nicknamed "nitrogen diamond."

Nitrogen History

Nitrogen (Latin nitrogenium, where nitrum (from Greek nitron νιτρον) means "saltpetre" (see nitre), and genes γενης means "forming") is formally considered to have been discovered by Daniel Rutherford in 1772, who called it noxious air or fixed air.^[1] The fact that there was an element of air which did not support combustion was clear to Rutherford. Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as "mephitic air" or azote, from the Greek word άζωτος (azotos) meaning "lifeless". In it animals died and flames were extinguished. Lavoisier's name for nitrogen is used in many languages (French, Polish, Russian, etc.) and still remains in English in the common names of many compounds, such as hydrazine and compounds of the azide ion.

Nitrogen compounds were well known during the Middle Ages. Alchemists knew nitric acid as aqua fortis (strong water). The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold (the king of metals). The earliest military, industrial and agricultural applications of nitrogen compounds involved uses of saltpeter (sodium nitrate or potassium nitrate), notably in gunpowder, and later as fertilizer. In 1910, Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", an allotrope considered to be monatomic. The "whirling cloud of brilliant yellow light" produced by his apparatus reacted with quicksilver to produce explosive mercury nitride

Nitrogen Chemicals

Nitrogen ( /ˈnaɪtrɵdʒɪn/ NYE-tro-jin) is a chemical element that has the symbol N, atomic number of 7 and atomic mass 14.00674 u. Elemental nitrogen is a colorless, odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78.08% by volume of Earth's atmosphere. The element nitrogen was discovered as a separable component of air, by Scottish physician Daniel Rutherford, in 1772.

Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in breaking the bond to convert the N₂ into useful compounds, but at the same time causing release of large amounts of often useful energy when the compounds burn, explode, or decay back into nitrogen gas.

Nitrogen occurs in all living organisms, and the nitrogen cycle describes movement of the element from air into the biosphere and organic compounds, then back into the atmosphere. Synthetically-produced nitrates are key ingredients of industrial fertilizers, and also key pollutants in causing the eutrophication of water systems. Nitrogen is a constituent element of amino acids and thus of proteins, and of nucleic acids (DNA and RNA). It resides in the chemical structure of almost all neurotransmitters, and is a defining component of alkaloids, biological molecules produced by many organisms.

Element Formulas

Obj. 5. From the name of an element, determine the formula (if different from the symbol).

The only elements for which we have to deal with the formula being different from the symbol are the diatomic elements. The diatomic elements are hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine and iodine. Remember, these nonmetals are found across the top of the periodic table, plus the halogens. (You could also remember S₈ and P₄ and other elements in their families, but the formulas of their molecules can change with conditions, and they are usually represented by their symbols rather than their formulas.) With that information, try your hand at exercise 5.

Exercises

Give the formula for each of the following elements if it is different than the symbol for the element.

a. hydrogen
b. sodium
c. silicon
d. nitrogen
e. oxygen
f. fluorine
g. neon

Answers to Exercises

Give the formula for each of the following elements if it is different than the symbol for the element.

a. hydrogen - The formula is H₂.
b. sodium - Sodium does not have a different formula, so we just use the symbol Na.
c. silicon - For silicon, we use the symbol Si.
d. nitrogen - Nitrogen has the formula N₂.
e. oxygen - Oxygen has the formula O₂.
f. fluorine - Fluorine has the formula F₂.
g. neon - For neon we use the symbol Ne.

Types of Chemicals

Types of Chemicals

The first seven objectives for this lesson deal with identifying various types or classifications of chemicals. These range from the simple awareness of whether you are dealing with an element or a compound to determining whether a compound exists as a network or as molecules.

The links to the left will take you to the pages for each of these objectives as noted.

CHEMICALS IN THE WORKPLACE

This Module provides trainees with background information on chemical hazards in the workplace. Topics discussed include: types of chemical hazards found in the workplace, how chemicals can harm you, how to obtain and understand information about chemicals used at work, and the role of the health and safety representative in ensuring the safe use of chemicals found in the workplace.

Objectives

At the end of this Module, trainees will be able to: (1) give examples of several types of chemicals commonly found in the workplace; (2) give several examples of how chemicals can affect your health; (3) describe at least two ways to obtain and use information about chemicals used in the workplace.

Adhesive Products

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Eco Green Technology

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Chemicals Companies

The largest corporate producers worldwide, with plants in numerous countries, are BASF, Dow, Degussa, Eastman Chemical Company, Shell, Bayer, INEOS, ExxonMobil, DuPont, SABIC, Braskem and Mitsubishi, along with thousands of smaller firms.

In the U.S. there are 170 major chemical companies. They operate internationally with more than 2,800 facilities outside the U.S. and 1,700 foreign subsidiaries or affiliates operating. The U.S. chemical output is $400 billion a year. The U.S. industry records large trade surpluses and employs more than a million people in the United States alone. The chemical industry is also the second largest consumer of energy in manufacturing and spends over $5 billion annually on pollution abatement.

In Europe, especially Germany, the chemical, plastics and rubber sectors are among the largest industrial sectors.^{[citation needed]} Together they generate about 3.2 million jobs in more than 60,000 companies. Since 2000 the chemical sector alone has represented 2/3 of the entire manufacturing trade surplus of the EU. The chemical sector accounts for 12% of the EU manufacturing industry's added value.

The chemical industry has shown rapid growth for more than fifty years. The fastest growing areas have been in the manufacture of synthetic organic polymers used as plastics, fibres and elastomers. Historically and presently the chemical industry has been concentrated in three areas of the world, Western Europe, North America and Japan (the Triad). The European Community remains the largest producer area followed by the USA and Japan.

The traditional dominance of chemical production by the Triad countries is being challenged by changes in feedstock availability and price, labour cost, energy cost, differential rates of economic growth and environmental pressures. Instrumental in the changing structure of the global chemical industry has been the growth in China, India, Korea, the Middle East, South East Asia, Nigeria, and Brazil.

Chemicals Companies

The largest corporate producers worldwide, with plants in numerous countries, are BASF, Dow, Degussa, Eastman Chemical Company, Shell, Bayer, INEOS, ExxonMobil, DuPont, SABIC, Braskem and Mitsubishi, along with thousands of smaller firms.

In the U.S. there are 170 major chemical companies. They operate internationally with more than 2,800 facilities outside the U.S. and 1,700 foreign subsidiaries or affiliates operating. The U.S. chemical output is $400 billion a year. The U.S. industry records large trade surpluses and employs more than a million people in the United States alone. The chemical industry is also the second largest consumer of energy in manufacturing and spends over $5 billion annually on pollution abatement.

In Europe, especially Germany, the chemical, plastics and rubber sectors are among the largest industrial sectors.^{[citation needed]} Together they generate about 3.2 million jobs in more than 60,000 companies. Since 2000 the chemical sector alone has represented 2/3 of the entire manufacturing trade surplus of the EU. The chemical sector accounts for 12% of the EU manufacturing industry's added value.

The chemical industry has shown rapid growth for more than fifty years. The fastest growing areas have been in the manufacture of synthetic organic polymers used as plastics, fibres and elastomers. Historically and presently the chemical industry has been concentrated in three areas of the world, Western Europe, North America and Japan (the Triad). The European Community remains the largest producer area followed by the USA and Japan.

The traditional dominance of chemical production by the Triad countries is being challenged by changes in feedstock availability and price, labour cost, energy cost, differential rates of economic growth and environmental pressures. Instrumental in the changing structure of the global chemical industry has been the growth in China, India, Korea, the Middle East, South East Asia, Nigeria, and Brazil.

Chemical Product Category Breakdown

Sales of the chemical business can be divided into a few broad categories, including basic chemicals (about 35 to 37 percent of the dollar output), life sciences (30 percent), specialty chemicals (20 to 25 percent) and consumer products (about 10 percent).
Basic chemicals, or "commodity chemicals" are a broad chemical category including polymers, bulk petrochemicals and intermediates, other derivatives and basic industrials, inorganic chemicals, and fertilizers. Typical growth rates for basic chemicals are about 0.5 to 0.7 times GDP
Product prices are generally less than fifty cents per pound.
Polymers, the largest revenue segment at about 33 percent of the basic chemicals dollar value, includes all categories of plastics and man-made fibers.
The major markets for plastics are packaging, followed by home construction, containers, appliances, pipe, transportation, toys, and games. The largest-volume polymer product, polyethylene (PE), is used mainly in packaging films and other markets such as milk bottles, containers, and pipe. Polyvinyl chloride (PVC), another large-volume product, is principally used to make pipe for construction markets as well as siding and, to a much smaller extent, transportation and packaging materials. Polypropylene (PP), similar in volume to PVC, is used in markets ranging from packaging, appliances, and containers to clothing and carpeting. Polystyrene (PS), another large-volume plastic, is used principally for appliances and packaging as well as toys and recreation. The leading man-made fibers include polyester, nylon, polypropylene, and acrylics, with applications including apparel, home furnishings, and other industrial and consumer use. The principal raw materials for polymers are bulk petrochemicals.
Chemicals in the bulk petrochemicals and intermediates are primarily made from liquefied petroleum gas (LPG), natural gas, and crude oil. Their sales volume is close to 30 percent of overall basic chemicals.
Typical large-volume products include ethylene, propylene, benzene, toluene, xylenes, methanol, vinyl chloride monomer (VCM), styrene, butadiene, and ethylene oxide. These chemicals are the starting points for most polymers and other organic chemicals as well as much of the specialty chemicals category.

Other derivatives and basic industrials include synthetic rubber, surfactants, dyes and pigments, turpentine, resins, carbon black, explosives, and rubber products and contribute about 20 percent of the basic chemicals' external sales. Inorganic chemicals (about 12 percent of the revenue output) make up the oldest of the chemical categories. Products include salt, chlorine, caustic soda, soda ash, acids (such as nitric, phosphoric, and sulfuric), titanium dioxide, and hydrogen peroxide. Fertilizers are the smallest category (about 6 percent) and include phosphates, ammonia, and potash chemicals.

Life sciences (about 30 percent of the dollar output of the chemistry business) include differentiated chemical and biological substances, pharmaceuticals, diagnostics, animal health products, vitamins, and crop protection chemicals. While much smaller in volume than other chemical sectors, their products tend to have very high prices—over ten dollars per pound—growth rates of 1.5 to 6 times GDP, and research and development spending at 15 to 25 percent of sales. Life science products are usually produced with very high specifications and are closely scrutinized by government agencies such as the Food and Drug Administration. Crop protection chemicals, about 10 percent of this category, include herbicides, insecticides, and fungicides.

Specialty chemicals are a category of relatively high valued, rapidly growing chemicals with diverse end product markets. Typical growth rates are one to three times GDP with prices over a dollar per pound. They are generally characterized by their innovative aspects. Products are sold for what they can do rather than for what chemicals they contain. Products include electronic chemicals, industrial gases, adhesives and sealants as well as coatings, industrial and institutional cleaning chemicals, and catalysts. Coatings make up about 15 percent of specialty chemicals sales, with other products ranging from 10 to 13 percent. Specialty Chemicals are sometimes referred to as "fine chemicals"

Consumer products include direct product sale of chemicals such as soaps, detergents, and cosmetics. Typical growth rates are 0.8 to 1.0 times GDP.

Every year, the American Chemistry Council tabulates the U.S. production of the top 100 basic chemicals. In 2000, the aggregate production of the top 100 chemicals totalled 502 million tons, up from 397 million tons in 1990. Inorganic chemicals tend to be the largest volume, though much smaller in dollar revenue terms due to their low prices. The top 11 of the 100 chemicals in 2000 were sulfuric acid (44 million tons), nitrogen (34), ethylene (28), oxygen (27), lime (22), ammonia (17), propylene (16), polyethylene (15), chlorine (13), phosphoric acid (13) and diammonium phosphates (12).

Chemical industry

The chemical industry comprises the companies that produce industrial chemicals. Central to the modern world economy, it converts raw materials (oil, natural gas, air, water, metals, and minerals) into more than 70,000 different products

Products

Polymers and plastics, especially polyethylene, polypropylene, polyvinyl chloride, polyethylene terephthalate, polystyrene and polycarbonate comprise about 80% of the industry’s output worldwide.^{[citation needed]} Chemicals are used to make a wide variety of consumer goods, as well as thousands inputs to agriculture, manufacturing, construction, and service industries. The chemical industry itself consumes 26 percent of its own output. Major industrial customers include rubber and plastic products, textiles, apparel, petroleum refining, pulp and paper, and primary metals. Chemicals is nearly a $3 trillion global enterprise, and the EU and U.S. chemical companies are the world's largest producers.

Chemical Safety

People use chemicals every day for a wide array of purposes, which can include work and house hold duties. Many of us fail to realize that we are actually handling potentially deadly chemicals when we simply clean the bathroom or wash the car. This brings to mind the reason why chemical safety is so important and why you should always use the chemicals that you own for their intended purpose only. Otherwise, the repercussions could be disastrous to say the least. There are specific things that you and those around you should do when handling, storing or using chemicals in nearly every form. This will ensure that you and everyone else is properly protected from the side effects that these chemicals can have.

Proper storage is first and foremost when it comes to chemicals. Most often chemicals are sold in the type of storage device that they should be kept in. Whether it be glass, plastic or metal, there is a reason why they are kept in such containers. Chemicals can be very reactive with certain components. When this reaction occurs there can be deadly fumes emitted or even a dangerous fire or explosion that is the end result. This is not to mention the fact that many chemicals also have corrosive properties that can eat through certain materials. Check the storage label on the chemicals before changing the container that they are stored in.

Proper labeling is also of great importance. It is possible for there to be a deadly reaction if the wrong chemicals are mixed because of improper labeling. To be safe you should always label the container with all of the contents and add any warnings that go along with it. You can also add the procedures that should be used in case of ingestion or accidental spill. This will add another level of safety that is very important.

Follow all of the instructions as to where in the home or building that you should store all of the chemicals that you use. There are many precautions that need to be taken in this area as chemicals need to be stored in the proper temperature and away from certain other things that they could react with. Bear in mind the climate of the area where you will store any type of chemical and be sure that the area is suitable.

All of these measures will help to keep you and those around you safe when dealing with chemicals.

The Straight Scoop on Uranium

Blame the scientist who “almost” discovered uranium on why this atomic element is named after the seventh planet from the sun. Imagine if the sequence below had taken place in any other way. What would we now be calling the yellowcake that powers nuclear reactors across the world? You would be surprised. This is the story behind uranium’s name.

The word “uranium” has a confusing past, but through no fault of its own. Since the beginning of the sixteenth century, in a silver mining town in an area which is now part of the Czech Republic, miners discovered a black mineral they called “pechblende.” Pitchblende, or uraninite as it is now better known, is a uranium-rich mineral which is also comprised of lead, thorium, radio and rare earths. In the late 19th century, it was from this same northwest Bohemian town where Marie Curie got her pitchblende and isolated radium and polonium from the ore.

European scientists Roentgen, Becquerel, Villard, and others were aggressively experimenting with pitchblende and discovered ionizing radiation, X-rays, beta radiation and gamma rays. Pierre and Marie Curie named the gamma ray phenomenon, attributed to the radium in pitchblende, “radioactivity.” MIT professor of biology Samuel Prescott, who was closely following Madam Curie’s research, began testing those gamma rays on food. He discovered the gamma rays destroyed bacteria in food. From Prescott’s work, food manufacturers discovered they could extend the shelf life of canned goods. Since then, radiation and radioactivity have become an integral part of both the medical profession and the food industry.

Let’s go back about one century. In 1789, Martin Heinrich Klaproth presented his discovery of a “strange kind of half metal” to Berlin’s Royal Academy of Sciences. The German chemist had, on the face of it, isolated uranium oxide from pitchblende. Klaproth suggested this new atomic element (number 92 on the periodic chart) be called “uran.”

Not until 1841 did another European scientist, Eugene-Melchior Peligot, finally isolate uranium as an atomic element. Klaproth was just stabbing in the dark when he tried to identify what “uranium” was. He failed to explain what uranium was, or even to understand it. Nonetheless, his credibility remained intact as a pioneering scientist. Martin Klaproth was later credited for isolating zirconium, chromium and cerium.

Klaproth’s naming ceremony for uranium was a political move, moreso than a scientific christening of the 92nd element. It was because of Dr. Bode. His Royal Academy colleague, German astronomer Johann Elert Bode, had been fuming since England’s William Herschel had discovered the seventh planet. Herschel honored King George III by calling this planet, “the Georgium Sidus (the Georgian Planet). Bode argued the new planet be renamed to conform to the classically mythological names of the other planets, such as Mercury, Mars, Venus, Jupiter and Saturn. Bode chose Uranus, the Greek name for the earliest supreme god.

The Uranus planetary debate went on for about, and was finally settled in 1850, about the same time that a British firm began using uranium in glass to give it a fluorescent yellow or greenish appearance. The point is this: If Klaproth hadn’t contributed to the Uranus-versus-Georgium Sidus debate by naming his “strange half metal” uran, we might be call uranium stocks something else.
About the Author:
James Finch contributes to StockInterview.com and other publications. His archived articles can be found on the internet news website, StockInterview.com, which can be found at http://www.stockinterview.com The above article was a brief excerpt from the upcoming publication, entitled "Investors Guide to Uranium Stocks." For more information about this book, please contact editor@stockinterview.com

Elements as Building Blocks

As you probably saw, the periodic table is organized like a big grid. The elements are placed in specific places because of the way they look and act. If you have ever looked at a grid, you know that there are rows (left to right) and columns (up and down). The periodic table has rows and columns, too, and they each mean something different.

You've got Your Periods...

Periodic Table showing PeriodsEven though they skip some squares in between, all of the rows go left to right. When you look at a periodic table, each of the rows is considered to be a different period (Get it? Like PERIODic table.). In the periodic table, elements have something in common if they are in the same row. All of the elements in a period have the same number of atomic orbitals. Every element in the top row (the first period) has one orbital for its electrons. All of the elements in the second row (the second period) have two orbitals for their electrons. It goes down the periodic table like that. At this time, the maximum number of electron orbitals or electron shells for any element is seven.

...and Your Groups

Periodic Table showing GroupsNow you know about periods. The periodic table has a special name for its columns, too. When a column goes from top to bottom, it's called a group. The elements in a group have the same number of electrons in their outer orbital. Every element in the first column (group one) has one electron in its outer shell. Every element on the second column (group two) has two electrons in the outer shell. As you keep counting the columns, you'll know how many electrons are in the outer shell. There are some exceptions to the order when you look at the transition elements, but you get the general idea.

Two at the Top

Periodic Table showing hydrogen and helium Hydrogen (H) and helium (He) are special elements. Hydrogen can have the talents and electrons of two groups, one and seven. To scientists, hydrogen is sometimes missing an electron, and sometimes it has an extra. Helium is different from all of the other elements. It can only have two electrons in its outer shell. Even though it only has two, it is still grouped with elements that have eight (inert gases).

The elements in the center section are called transition elements. They have special electron rules.

Diamonds Are Forever

Diamonds are still a girl's best friend, right? We love the shiny gems. They are the most popular rocks sold today. But what exactly are they, anyway? Where do they come from? What else are they used for?

Diamonds are a mineral in one of the two crystalline forms of the element carbon. They are the hardest natural substance man knows. Diamonds are sold as gems, and used in industrial applications for smoothing, cutting, and polishing hard materials.

Diamonds are most famous for crystallizing in the common colorless form. They may also be tranlucent to transparent white, yellow, green, blue, or brown. Diamonds have a high refractive index which is why they are so brilliant and sparkly after cutting. The familiar shape of the diamond is the octahedron.

The most brilliant diamonds become gemstones for jewelry and other uses. For those that don't make it to gems, there are other options. There is bort, which is a more poorly crystallized or undesirable color and in fragmentary condition, and carbonados which is gray to black opaque. Bort and carbonados are used as abrasives for the cutting of diamonds and the cutting heads of industrial rock drills.

Diamonds are found in alluvial formations and in volcanic pipes, filled for most of their length with blue ground or kimberlite, and igneous rock consisting primarily of serpentine. Diamond yielding earth is mined by both the open-pit method and by underground mining. After removal to the surface, the soil is crushed and concentrated. Passing the concentrated material in a stream of water over greased tables does the needed sorting. The diamond is largely water repellent and sticks to the grease and the other minerals retain a film of water, which prevents the sticking to the grease. Then the diamonds are removed from the grease, cleaned, and graded for sale and use.

The earliest sources of gem diamonds were India and Borneo. Some famous diamonds are the Great Mogul, Regent, and Pitt. Other famous diamonds include the Hope (blue), Dresden (green) and Tiffany (yellow). In the early 18th century, deposits similar to those in India were found in Brazil, mainly of carbonados. In 1867 a stone found in South Africa was recognized as a diamond. Within a few years began a wild search for diamonds. In 1870-1871, dry diggings including most of the celebrated mines were discovered.

Synthetic diamonds were successfully produced in 1955; a number of small crystals were produced when pure graphite mixed with a catalyst was subjected to pressure of about 1 million lb per sq in. and temperature of the order of 5,000-F (3,000-C). Synthetic diamonds now are extensively used for industry, mainly due to the ease of obtaining and lower cost for them. Diamonds are still very popular and symbolize many things. Their popularity does not seem to be dwindling any time in the near future.

What are Compound Microscopes?

Most of the microscopes used today are compound. A compound microscope features two or more lenses. A hollow cylinder called the tube connects the two lenses. The top lens, the one people look through, is called the eyepiece. The bottom lens is known as the objective lens. Below the two lenses is the stage, with the illuminator below that.

Compound microscopes were among the first magnifying instruments invented. Two Dutch eyeglass makers named Zaccharias and Hans Janssen are credited with making the first compound microscope in 1590 by putting one lens at the top of a tube and another at the bottom of the tube. Their idea was fleshed out by others scientists over the next several centuries, but the basic design remained very similar.

The eyepiece, also known as the ocular lens, is at the top of the compound microscope. It is not adjustable, that is, it only has one strength. Most ocular lenses are 10x, meaning that they magnify objects to ten times their normal size. People look in through the eyepiece through the tube and out through the objective lens.

A compound microscope normally contains several objective lenses. The objective lenses are different lengths, with the longer ones being the strongest. The lenses are situated on a round disk below the tube. Viewers choose which strength lens they want and place it below the tube by turning the disk until the desired lens is in place.

The stage and illuminator are below the objective lens. Specimens are placed over a translucent part of the stage. Light provided by the illuminator shines through the clear part of the stage, making it easier for the viewer to see the magnified details of the specimen. Two adjustment knobs help focus the object on the stage by bringing the lenses and the stage closer together.

Compound microscopes have been around for hundreds of years and are still very useful. A number of scientific disciplines use compound microscopes to discover the wonders of the microscopic world.

Atoms Around Us

If you want to have a language, you will need an alphabet. If you want to build proteins, you will need amino acids. Other examples in chemistry are not any different. If you want to build molecules, you will need elements. Each element is a little bit different from the rest. Those elements are the alphabet to the language of molecules.

Why are we talking about elements? This is the section on atoms.

Let's stretch the idea a bit. If you read a book, you will read a language. Letters make up that language. But what makes those letters possible? Ummm... Ink? Yes! You need ink to crate the letters. And for each letter, it is the same type of ink.

Confused? Don't be. Elements are like those letters. They have something in common. That's where atoms come in. All elements are made of atoms. While the atoms may have different weights and organization, they are all built in the same way. Electrons, protons, and neutrons make the universe go.

If you want to do a little more thinking, start with particles of matter. Matter, the stuff around us, is used to create atoms. Atoms are used to create the elements. Elements are used to create molecules. It just goes on. Everything you see is built by using something else.

You could start really small...
- Particles of matter
- Atoms
- Elements
- Molecules
- Macromolecules
- Cell organelles
- Cells
- Tissues
- Organs
- Systems
- Organisms
- Populations
- Ecosystems
- Biospheres
- Planets
- Planetary Systems with Stars
- Galaxies
- The Universe
.And finish really big.

Wow. All of that is possible because of atoms.

Binding Energy (nuclear binding energy)

The energy equivalent (E = mc^2) of the mass deficiency of an atom.
where: E = is the energy in joules, m is the mass in kilograms, and c is the speed of light in m/s^2

Protein Design: Automated protein discovery and synthesis

In this paper I describe (theoretically) the method(s) of automated protein discovery and synthesis.

1. Protein Folding Problem To solve the protein folding problem we can use Artificial Neural Networks. We will train the networks with natural proteins whose 3D structure and amino acid sequence is known. After that we will test the network with few new artificially designed proteins to check if it works correctly. If it doesn't, we will be changing some of the network's parameter such as training iterations, no of hidden layers, etc. And train the network again.

To check the protein's 3D structure, we need to have a model of actual physical world in the computer model.

2. Simulation of Physical World This is the trickiest part. To simulate the physical world at the atomic level is very difficult. We need to take into account: covalent bonds, spatial & temporal parameters, weak interactions such as hydrogen bonds, dipole interactions, etc. We also need to simulate chemical reactions. This will probably require huge amounts of computing power.

Or perhaps, neural networks can be employed here also as the little inaccuracy produced by a neural network can take care of randomness at quantum level. The neural networks will be used to predict/calculate the magnitude of effect of various forces on an atom/molecule and also how these behave at a grander inter-molecular level.

3. Designing Proteins To design proteins, we will be using Genetic Algorithm method. The random amino-acid sequences will be evolved & tested by converting these sequences into their respective 3D shape by the trained neural network. The best sequences will be retained, while other mutated or crossed-over, etc. The fitness function will work in the simulated physical world. If the protein produced is successful in carrying out our desired unction, then it is fit else it is not. Actually we will assign a fitness level from 0 to 100. Once the final amino acid sequence is determined, it will be sent to the Protein Printer.

4. Protein Printer This is the only hardware part of the whole procedure. It will produce the desired real proteins from the amino acid sequence received from software. It may be able to work in any of the two ways:

* Artificial Ribosome: It will mimic the functionality of the cell to produce proteins. We will generate an mRNA using some assembling mechanism. Then, our artificially designed Ribosome will translate it into a protein which we can use. * Artificial Recombinant DNA: We will assemble a fragment of DNA corresponding to desired amino acid sequence. Then by using some automated means we will introduce the DNA into a colony of E. coli (or some other organism)/ Then E. coli will produce these synthetic proteins in the same way they produce natural proteins in recombinant DNA technology.

5. Conclusion Using this system, we only need to define: "What do we want the protein to do?". All other procedure is automatic. We just need to tell if we want a protein to degrade plastic, convert CO2 into diamond and oxygen, and catalyze/initiate cold fusion, etc. & we will have ready made proteins. It can also help in finding proteins which will help us attain Immortality.

The potential is immense. The only need is its correct use.

Chemicals Products

Chemical Bonds

The attractive forces that hold atoms together in elements or compounds.

Chemical Change

A change in which one or more new substances are formed

Chemical Equation

Description of a chemical reaction by placing the formulas of the reactants on the left and the formulas of products on the right of an arrow.

Chemical Equilibrium

A state of dynamic balance in which the rates of forward and reverse reactions are equal, there is no net change in concentrations of reactants or products while a system is at equilibrium.

Chemicals

Chemical Hygiene Officer (CHO)

A person or employee who is qualified by training or experience to provide technical guidance in the development and implementations of the provisions of a Chemical Hygiene Plan (CHP)

Chemical Hygiene Plan (CHP)

A written program developed and implemented by an employer designating proceedures, equipment, personal protective equipment, and work practices that are capable of protecting employees from the health hazards presented by hazardous chemicals usid in that particular workplace.

Chemicals

Chemical Kinetics

The study of rates and mechanisms of chemical reactions and of the factors on which they depend.

Chemical Periodicity

The variations in properties of elements with their position in the periodic table.

Electrolytic Cells

Electrochemical cells in which electrical energy causes nospontaneous redox reactions to occur.

An electrochemical cell in which chemical reactions are forced to occur by the application of an outside source of electrical energy.

What is uranium? How does it work?

• Uranium is a very heavy metal which can be used as an abundant source of concentrated energy.
• It occurs in most rocks in concentrations of 2 to 4 parts per million and is as common in the Earth's crust as tin, tungsten and molybdenum. It occurs in seawater, and can be recovered from the oceans.
• It was discovered in 1789 by Martin Klaproth, a German chemist, in the mineral called pitchblende. It was named after the planet Uranus, which had been discovered eight years earlier.
• Uranium was apparently formed in supernovae about 6.6 billion years ago. While it is not common in the solar system, today its slow radioactive decay provides the main source of heat inside the Earth, causing convection and continental drift.
• The high density of uranium means that it also finds uses in the keels of yachts and as counterweights for aircraft control surfaces, as well as for radiation shielding.
• Its melting point is 1132°C. The chemical symbol for uranium is U.

The Uranium Atom

On a scale arranged according to the increasing mass of their nuclei, uranium is the heaviest of all the naturally-occurring elements (Hydrogen is the lightest). Uranium is 18.7 times as dense as water.

Like other elements, uranium occurs in several slightly differing forms known as 'isotopes'. These isotopes differ from each other in the number of particles (neutrons) in the nucleus. Natural uranium as found in the Earth's crust is a mixture largely of two isotopes: uranium-238 (U-238), accounting for 99.3% and uranium-235 (U-235) about 0.7%.

The isotope U-235 is important because under certain conditions it can readily be split, yielding a lot of energy. It is therefore said to be 'fissile' and we use the expression 'nuclear fission'.

Meanwhile, like all radioactive isotopes, they decay. U-238 decays very slowly, its half-life being about the same as the age of the Earth (4500 million years). This means that it is barely radioactive, less so than many other isotopes in rocks and sand. Nevertheless it generates 0.1 watts/tonne as decay heat and this is enough to warm the Earth's core. U-235 decays slightly faster.

Energy from the uranium atom

The nucleus of the U-235 atom comprises 92 protons and 143 neutrons (92 + 143 = 235). When the nucleus of a U-235 atom captures a moving neutron it splits in two (fissions) and releases some energy in the form of heat, also two or three additional neutrons are thrown off. If enough of these expelled neutrons cause the nuclei of other U-235 atoms to split, releasing further neutrons, a fission 'chain reaction' can be achieved. When this happens over and over again, many millions of times, a very large amount of heat is produced from a relatively small amount of uranium.

It is this process, in effect "burning" uranium, which occurs in a nuclear reactor. The heat is used to make steam to produce electricity.

Uranium Inside the reactor

uranium fuel is assembled in such a way that a controlled fission chain reaction can be achieved. The heat created by splitting the U-235 atoms is then used to make steam which spins a turbine to drive a generator, producing electricity.
Nuclear power stations and fossil-fuelled power stations of similar capacity have many features in common. Both require heat to produce steam to drive turbines and generators. In a nuclear power station, however, the fissioning of uranium atoms replaces the burning of coal or gas .

The chain reaction that takes place in the core of a nuclear reactor is controlled by rods which absorb neutrons and which can be inserted or withdrawn to set the reactor at the required power level.

The fuel elements are surrounded by a substance called a moderator to slow the speed of the emitted neutrons and thus enable the chain reaction to continue. Water, graphite and heavy water are used as moderators in different types of reactors.

Because of the kind of fuel used (ie the concentration of U-235, see below), if there is a major uncorrected malfunction in a reactor the fuel may overheat and melt, but it cannot explode like a bomb.

A typical 1000 megawatt (MWe) reactor can provide enough electricity for a modern city of up to one million people.

Uranium and Plutonium

a reactor core (most of the fuel), these reactions occur frequently, and in fact about one third of the energy yield comes from "burning" Pu-239.

But sometimes a Pu-239 atom simply captures a neutron without splitting, and it becomes Pu-240. Because the Pu-239 is either progressively "burned" or becomes Pu-240, the longer the fuel stays in the reactor the more Pu-240 is in it.*

* The significance of this is that when the spent fuel is removed after about three years, the plutonium in it is not suitable for making weapons but can be recycled as fuel.

From uranium ore to reactor fuel
Uranium ore can be mined by underground or open-cut methods, depending on its depth. After mining, the ore is crushed and ground up. Then it is treated with acid to dissolve the uranium, which is recovered from solution.

Uranium may also be mined by in situ leaching (ISL), where it is dissolved from a porous underground ore body in situ and pumped to the surface.

The end product of the mining and milling stages, or of ISL, is uranium oxide concentrate (U₃O₈). This is the form in which uranium is sold.

Before it can be used in a reactor for electricity generation, however, it must undergo a series of processes to produce a useable fuel.

For most of the world's reactors, the next step in making the fuel is to convert the uranium oxide into a gas, uranium hexafluoride (UF₆), which enables it to be enriched. Enrichment increases the proportion of the uranium-235 isotope from its natural level of 0.7% to 3 - 4%. This enables greater technical efficiency in reactor design and operation, particularly in larger reactors, and allows the use of ordinary water as a moderator.

After enrichment, the UF₆ gas is converted to uranium dioxide (UO₂) which is formed into fuel pellets. These fuel pellets are placed inside thin metal tubes which are assembled in bundles to become the fuel elements or assemblies for the core of the reactor.

For reactors which use natural uranium as their fuel (and hence which require graphite or heavy water as a moderator) the U₃O₈ concentrate simply needs to be refined and converted directly to uranium dioxide.

When the uranium fuel has been in the reactor for about three years, the used fuel is removed, stored, and then either reprocessed or disposed of underground (see Nuclear Fuel Cycle or Radioactive Waste Management in this series).

Who uses nuclear power?

Who uses nuclear power?

Over 16% of the world's electricity is generated from uranium in nuclear reactors. This amounts to about 2400 billion kWh each year, as much as from all sources of electricity worldwide in 1960. In a current perspective, it is twelve times Australia's or South Africa's total electricity production, five times India's, twice China's and 500 times Kenya's total.

It comes from about 440 nuclear reactors with a total output capacity of about 370 000 megawatts (MWe) operating in 31 countries. About thirty more reactors are under construction and another 40 are planned.

Belgium, Bulgaria, Finland, France, Germany, Hungary, Japan, South Korea, Lithuania, Slovakia, Slovenia, Sweden, Switzerland and Ukraine all get 30% or more of their electricity from nuclear reactors. The USA has over 100 reactors operating, with capacity of almost three times Australia's total, and supplying 20% of its electricity. The UK gets almost a quarter of its electricity from uranium.

Who has and who mines uranium?

Who has and who mines uranium?

Uranium is widespread in many rocks, and even in seawater. However, like other metals, it is seldom sufficiently concentrated to be economically recoverable. Where it is, we speak of an orebody. In defining what is ore, assumptions are made about the cost of mining and the market price of the metal. Uranium reserves are therefore calculated as tonnes recoverable up to a certain cost.

Australia's reasonably assured resources of uranium are 732,000 tonnes U recoverable at up to US$80/kg U (well under the market 'spot' price), Canada's are 345,000 tonnes U. Australia's resources in this category are about 27% of the world's total, Canada's 13%.

Although it has more than any other country, Australia is not the only one with major deposits. Others in order are: Kazakhstan (17% of world total), Canada, USA, South Africa, Namibia, Brazil, Niger and Russia. Many more countries have smaller deposits which could be mined if needed.

Despite being so well-endowed with uranium reserves, political factors mean that Canada is well in front of Australia as the main supplier of uranium to world markets.

Uranium is sold only to countries which are signatories of the Nuclear Non-Proliferation Treaty, and which allow international inspection to verify that it is used only for peaceful purposes. Customer countries for Australia's uranium must also have a bilateral safeguards agreement with Australia. Canada has similar arrangements.

Australian exports in 2005 amounted to over 12,000 tonnes of U₃O₈ valued at nearly A$600 million. This was about 24% of world mine production of uranium. Canada produced almost 14,000 tonnes of U₃O₈ in 2005, about one third of world production and mostly for export.

Other uses of nuclear energy

Many people, when talking about nuclear energy, have only nuclear reactors (or perhaps nuclear weapons) in mind. Few people realise the extent to which the use of radioisotopes has changed our lives over the last few decades.

Using relatively small special-purpose nuclear reactors it has become possible to make a wide range of radioactive materials (radioisotopes) at low cost. For this reason the use of artificially produced radioisotopes has become widespread since the early 1950s, and there are now some 270 "research" reactors in 59 countries producing them.

Radioisotopes

In our daily life we need food, water and good health. Today, radioactive isotopes play an important part in the technologies that provide us with all three. They are produced by bombarding small amounts of particular elements with neutrons.

In medicine, radioisotopes are widely used for diagnosis and research. Radioactive chemical tracers emit gamma radiation which provides diagnostic information about a person's anatomy and the functioning of specific organs. Radiotherapy also employs radioisotopes in the treatment of some illnesses, such as cancer. More powerful gamma sources are used to sterilise syringes, bandages and other medical equipment. About one person in two in the western world is likely to experience the benefits of nuclear medicine in their lifetime, and gamma sterilisation of equipment is almost universal.

In the preservation of food, radioisotopes are used to inhibit the sprouting of root crops after harvesting, to kill parasites and pests, and to control the ripening of stored fruit and vegetables. Irradiated foodstuffs are accepted by world and national health authorities for human consumption in an increasing number of countries. They include potatoes, onions, dried and fresh fruits, grain and grain products, poultry and some fish. Some prepacked foods can also be irradiated.

In the growing of crops and breeding livestock, radioisotopes also play an important role. They are used to produce high yielding, disease-resistant and weather-resistant varieties of crops, to study how fertilisers and insecticides work, and to improve the productivity and health of domestic animals.

Industrially, and in mining, they are used to examine welds, to detect leaks, to study the rate of wear of metals, and for on-stream analysis of a wide range of minerals and fuels.

There are many other uses. A radioisotope derived from the plutonium formed in nuclear reactors is used in most household smoke detectors.

Radioisotopes are used by police to fight crime, in detecting and analysing pollutants in the environment, to study the movement of surface water and to measure water runoffs from rain and snow, as well as the flow rates of streams and rivers.

Other reactors

There are also other uses for reactors. Over 200 small nuclear reactors power some 150 ships, mostly submarines, but ranging from icebreakers to aircraft carriers. These can stay at sea for long periods without having to make refuelling stops . In the Russian Arctic where operating conditions are beyond the capability of conventional icebreakers , very powerful nuclear-powered vessels operate almost year-round, where previously only two months could be used each year.

The heat produced by nuclear reactors can also be used directly rather than for generating electricity. In Sweden and Russia, for example, it is used to heat buildings and to provide heat for a variety of industrial processes such as water desalination. Nuclear desalination is likely to be a major growth area in future.

Uranium Military weapons

Uranium Military Weapons of Chemicals

Both uranium and plutonium were used to make bombs before they became important for making electricity and radioisotopes. But the type of uranium and plutonium for bombs is different from that in a nuclear power plant. Bomb-grade uranium is highly-enriched (>90% U-235, instead of about 3.5%); bomb-grade plutonium is fairly pure (>90%) Pu-239 and is made in special reactors.

Today, due to disarmament, a lot of military uranium is becoming available for electricity production. The military uranium is diluted about 25:1 with depleted uranium (mostly U-238) from the enrichment process before being used.