CHEMISTRY OF CARBIDES

Carbide

carbide,  any of a class of chemical compounds in which carbon is combined with a metallic or semimetallic element. Calcium carbide is important chiefly as a source of acetylene and other chemicals, whereas the carbides of silicon, tungsten, and several other elements are valued for their physical hardness, strength, and resistance to chemical attack even at very high temperatures. Iron carbide (cementite) is an important constituent of steel and cast iron.

Preparation of carbides :

Carbides: Carbides are binary compounds of carbon with elements of lower or about equal electronegativity.

Preparation : Carbides are generally prepared by heating the elements orits oxide with carbon or hydrocarbon at very high temperatures.

 Ca + 2C ——→ BaC₂; 2Li + 2C ——→ Li₂C₂

CaO + 3C ——→ CaC2 + CO

4Li + C₂H₂ ——→ Li₂C₂ + LiH

Carbides are classified into three types on the basis of chemical bonding.

Carbide

Classification of carbides :

Classification of carbides based on structural type is rather difficult, but three broad classifications arise from general trends in their properties. The most electropositive metals form ionic or saltlike carbides, the transition metals in the middle of the periodic table tend to form what are called interstitial carbides, and the nonmetals of electronegativity similar to that of carbon form covalent or molecular carbides.

Ionic carbides

Ionic carbides have discrete carbon anions of the forms C⁴⁻, sometimes called methanides since they can be viewed as being derived from methane, (CH₄); C₂²⁻, called acetylides and derived from acetylene (C₂H₂); and C₃⁴⁻, derived from allene (C₃H₄). The best-characterized methanides are probably beryllium carbide (Be₂C) and aluminum carbide (Al₄C₃). Beryllium oxide (BeO) and carbon react at 2,000 °C (3,600 °F) to produce the brick-red beryllium carbide, whereas pale yellow aluminum carbide is prepared from aluminum and carbon in a furnace. Aluminum carbide reacts as a typical methanide with water to produce methane.Al₄C₃ + 12H₂O → 4Al(OH)₃ + 3CH₄

There are many acetylides that are well known and well characterized. In addition to those of the alkali metals and the alkaline-earth metals mentioned above, lanthanum (La) forms two different acetylides, and copper (Cu), silver (Ag), and gold (Au) form explosive acetylides. Zinc (Zn), cadmium (Cd), and mercury (Hg) also form acetylides, although they are not as well characterized. The most important of these compounds is calcium carbide, CaC₂. The primary use for calcium carbide is as a source of acetylene for use in the chemical industry. Calcium carbide is synthesized industrially from calcium oxide (lime), CaO, and carbon in the form of coke at about 2,200 °C (4,000 °F). Pure calcium carbide has a high melting point (2,300 °C [4,200 °F]) and is a colourless solid. The reaction of CaC₂ with water yields C₂H₂ and a significant amount of heat, so the reaction is carried out under carefully controlled conditions.CaO + 3C → CaC₂ + CO CaC₂ + 2H₂O → C₂H₂ + Ca(OH)₂ Calcium carbide also reacts with nitrogen gas at elevated temperatures (1,000–1,200 °C [1,800–2,200 °F]) to form calcium cyanamide, CaCN₂.CaC₂ + N₂ → CaCN₂ + C This is an important industrial reaction because CaCN₂ finds extensive use as a fertilizer owing to its reaction with water to produce cyanamide, H₂NCN. Most MC₂ acetylides have the CaC₂ structure, which is derived from the cubic sodium chloride (NaCl) structure. The C₂ units lie parallel along the cell axes, causing a distortion of the cell from cubic to tetragonal.

Salt-like (saline) carbides :

Salt-like carbides are composed of highly electropositive elements such as the alkali metals, alkaline earths, and group 3 metals including scandium, yttrium and lanthanum. Aluminum from group 13 forms carbides, but gallium, indium and thallium do not. These materials feature isolated carbon centers, often described as "C⁴⁻", in the methanides or methides; two atom units, "C₂²⁻" in the acetylides ; and three atom units "C₃⁴⁻" in the sesquicarbides.^[  The graphite intercalation compound KC₈, prepared from vapour of potassium and graphite, and the alkali metal derivatives of C₆₀ are not usually classified as carbides.

Methanides

Tungsten carbide drill bits.

Carbides of this class decompose in water producing methane. Two such examples are aluminium carbide Al₄C₃ and beryllium carbide Be₂C.

The reaction of transition metal carbides with water is very slow and is usually neglected. For example, depending on surface porosity, 5–30 atomic layers of titanium carbide are hydrolyzed within 5 minutes at ambient conditions, following by saturation of the reaction.

Acetylides

Sesquicarbides :

The polyatomic ion C₃^4–, sometimes called sesquicarbide, is found in Li₄C₃ and Mg₂C₃. The ion is linear and is isoelectronic with CO₂. The C-C distance in Mg₂C₃ is 133.2 pm. Mg₂C₃ yields methyl acetylene, CH₃CCH, and propadiene , CH₂CCH₂, on hydrolysis which was the first indication that it may contain C₃^4–.

Covalent carbides

The carbides of silicon and boron are described as "covalent carbides", although virtually all compounds of carbon exhibit some covalent character. Silicon carbide has two similar crystalline forms, which are both related to the diamond structure. Boron carbide, B₄C, on the other hand has an unusual structure which includes icosahedral boron units linked by carbon atoms. In this respect boron carbide is similar to the boron rich borides. Both silicon carbide, SiC, (carborundum) and boron carbide, B₄C are very hard materials and refractory. Both materials are important industrially. Boron also forms other covalent carbides, e.g. B₂₅C.

Interstitial carbides :

The carbides of the group 4, 5 and 6 transition metals (with the exception of chromium) are often described as interstitial compounds.  These carbides have metallic properties and are refractory. Some exhibit a range of stoichiometries, e.g. titanium carbide, TiC. Titanium carbide and tungsten carbide are important industrially and are used to coat metals in cutting tools.

The longheld view is that the carbon atoms fit into octahedral interstices in a close packed metal lattice when the metal atom radius is greater than approximately 135 pm:

• When the metal atoms are cubic close packed, (ccp), then filling all of the octahedral interstices with carbon achieves 1:1 stoichiometry with the rock salt structure.
• When the metal atoms are hexagonal close packed, (hcp), as the octahedral interstices lie directly opposite each other on either side of the layer of metal atoms, filling only one of these with carbon achieves 2:1 stoichiometry with the CdI₂ structure.

Intermediate transition metal carbides :

In these carbides, the transition metal ion is smaller than the critical 135 pm, and the structures are not interstitial but are more complex. Multiple stoichiometries are common, for example iron forms a number of carbides, Fe₃C, Fe₇C₃ and Fe₂C. The best known is cementite, Fe₃C, which is present in steels. These carbides are more reactive than the interstitial carbides, for example the carbides of Cr, Mn, Fe, Co and Ni all are hydrolysed by dilute acids and sometimes by water, to give a mixture of hydrogen and hydrocarbons. These compounds share features with both the inert interstitials and the more reactive salt-like carbides

Molecular carbides :

The complex [Au₆C(PPh₃)₆]²⁺, containing a carbon-gold core.

Some metals, such as lead and tin, are believed not to form carbides under any circumstances. There exists however a mixed titanium-tin carbide, which is a two-dimensional conductor. (In 2007, there were two reports of a lead carbide PbC₂, apparently of the acetylide type; but these claims have yet to be published in reviewed journals.)

Calcium carbide :

Calcium carbide is a chemical compound with the chemical formula of CaC₂. Its main use industrially is in the production of acetylene and calcium cyanamide.

The pure material is colorless, however pieces of technical-grade calcium carbide are grey or brown and consist of only 80-85% of CaC₂ (the rest is CaO, Ca₃P₂, CaS, Ca₃N₂, SiC, etc.). Because of presence of PH₃, NH₃, and H₂S, technical-grade calcium carbide has a distinctive smell which some find unpleasant.

Production

Calcium carbide is produced industrially in an electric arc furnace from a mixture of lime and coke at approximately 2000 °C. This method has not changed since its invention in 1888:

CaO + 3 C → CaC₂ + CO

The high temperature required for this reaction is not practically achievable by traditional combustion, so the reaction is performed in an electric arc furnace with graphite electrodes. The carbide product produced generally contains around 80% calcium carbide by weight. The carbide is crushed to produce small lumps that can range from a few mm up to 50 mm. The impurities are concentrated in the finer fractions. The CaC₂ content of the product is assayed by measuring the amount of acetylene produced on hydrolysis. As an example, the British and German standards for the content of the coarser fractions are 295 L/kg and 300 L/kg respectively. Impurities present in the carbide include phosphide, which produces phosphine when hydrolysed.

Crystal structure :

Pure calcium carbide is a colourless solid. The common crystalline form at room temperature is a distorted rock-salt structure with the C₂²⁻ units lying parallel.

Applications :

Production of acetylene

The reaction of calcium carbide with water, producing acetylene and calcium hydroxide, was discovered by Friedrich Wöhler in 1862.

CaC₂ + 2 H₂O → C₂H₂ + Ca(OH)₂

This reaction is the basis of the industrial manufacture of acetylene, and is the major industrial use of calcium carbide.

In China, acetylene derived from calcium carbide remains a raw material for the chemical industry, in particular for the production of polyvinyl chloride. Locally produced acetylene is more economical than using imported oil. Production of calcium carbide in China has been increasing. In 2005 output was 8.94 million tons, with the capacity to produce 17 million tons.

Carbide lamps

Lit carbide lamp

Calcium carbide is used in carbide lamps, in which water drips on the carbide and the acetylene formed is ignited. These lamps were usable but dangerous in coal mines, where the presence of the flammable gas methane made them a serious hazard. The presence of flammable gases in coal mines led to miner safety lamps such as the Davy lamp, in which a wire gauze reduces the risk of methane ignition. However, carbide lamps were still used extensively in slate, copper, and tin mines, where methane is not a serious hazard, but most miner's lamps have now been replaced by electric lamps.

Carbide lamps are still used for mining in some less wealthy countries, for example in the silver mines near Potosí, Bolivia.

Other uses

In the ripening of fruit, calcium carbide is sometimes used as source of acetylene gas, which is a ripening agent similar to ethylene.

Calcium carbide is used in toy cannons such as the Big-Bang Cannon, as well as in bamboo cannons.

Calcium carbide, together with calcium phosphide, is used in floating, self-igniting naval signal flares, such as those produced by the Holmes' Marine Life Protection Association.

In the Netherlands, calcium carbide is still used in a traditional New Year's Eve custom called Carbidschieten (Carbide Shooting). To create an explosion, carbide and water are put in a milk churn with a lid, with ignition usually done with a torch. Some villages in the Netherlands fire multiple milk churns in a row. The custom may be derived from an old pagan religious practice intended to chase off spirits.

tremely hard like diamond and possess very high melting points.

Boron Carbide :

Boron Carbide is one of the hardest materials known, ranking third behind diamond and cubic boron nitride. It is the hardest material produced in tonnage quantities.  Originally discovered in mid 19^th century as a by-product in the production of metal borides, boron carbide was only studied in detail since 1930.

Boron carbide powder (see figure 1) is mainly produced by reacting carbon with B₂O₃ in an electric arc furnace, through carbothermal reduction or by gas phase reactions.  For commercial use B₄C powders usually need to be milled and purified to remove metallic impurities.

In common with other non-oxide materials boron carbide is difficult to sinter to full density, with hot pressing or sinter HIP being required to achieve greater than 95% of theoretical density. Even using these techniques, in order to achieve sintering at realistic temperatures (e.g. 1900 - 2200°C), small quantities of dopants such as fine carbon, or silicon carbide are usually required.

Property
Density (g.cm-3)	2.52
Melting Point (°C)	2445
Hardness (Knoop 100g) (kg.mm-2)	2900-3580
Fracture Toughness (MPa.m-½)	2.9 - 3.7
Young's Modulus (GPa)	450 - 470
Electrical Conductivity (at 25°C) (S)	140
Thermal Conductivity (at 25°C) (W/m.K)	30 - 42
Thermal Expansion Co-eff. x10-6 (°C)	5
Thermal neutron capture cross section (barn)	600

As an alternative, B₄C can be formed as a coating on a suitable substrate by vapour phase reaction techniques e.g. using boron halides or di-borane with methane or another chemical carbon source.

Key Properties:

Boron carbide is characterised by its:

•         Extreme hardness

•         Difficult to sinter to high relative densities without the use of sintering aids

•         Good chemical resistance                                                                                                                                                                                                         Table 1. Typical properties of boron carbide

•         Good nuclear properties

•         Low density

Typical properties for boron carbide are listed in table 1.

.

Applications :

Abrasives

Due to its high hardness, boron carbide powder is used as an abrasive in polishing and lapping applications, and also as a loose abrasive in cutting applications such as water jet cutting. It can also be used for dressing diamond tools.

Nozzles

The extreme hardness of boron carbide gives it excellent wear and abrasion resistance and as a consequence it finds application as nozzles for slurry pumping, grit blasting and in water jet cutters (see figure 2).

Nuclear applications

Its ability to absorb neutrons without forming long lived radio-nuclides make the material attractive as an absorbent for neutron radiation arising in nuclear power plants. Nuclear applications of boron carbide include shielding, and control rod and shut down pellets.

Ballistic Armour

Boron carbide, in conjunction with other materials also finds use as ballistic armour (including body or personal armour) where the combination of high hardness, high elastic modulus, and low density give the material an exceptionally high specific stopping power to defeat high velocity projectiles.

Other Applications

Other applications include ceramic tooling dies, precision toll parts, evaporating boats for materials testing and mortars and pestles.

Aluminium carbide :

Aluminium carbide, chemical formula Al₄C₃, is a carbide of aluminium. It has the appearance of pale yellow to brown crystals. It is stable up to 1400 °C. It decomposes in water with the production of methane.

Structure

Aluminium carbide has an unusual crystal structure that consists of two types of layers. It is based on AlC₄ tetrahedra of two types and thus two types of carbon atoms. One is surrounded by a deformed octahedron of 6 Al atoms at a distance of 217 pm. The other is surrounded by 4 Al atoms at 190–194 pm and a fifth Al atom at 221 pm. Other carbides (IUPAC nomenclature: methides) also exhibit complex structures.

Reactions

Aluminium carbide hydrolyses with evolution of methane. The reaction proceeds at room temperature but is rapidly accelerated by heating.

Al₄C₃ + 12 H₂O → 4 Al(OH)₃ + 3 CH₄

Similar reactions occur with other protic reagents:

Al₄C₃ + 12 HCl → 4 AlCl₃ + 3 CH₄

Preparation

Aluminium carbide is prepared by direct reaction of aluminium and carbon in an electric furnace.

4 Al + 3 C → Al₄C₃

An alternative reaction begins with alumina, but it is less favorable because of generation of carbon monoxide.

2 Al₂O₃ + 9 C → Al₄C₃ + 6 CO

Silicon carbide also reacts with aluminium to yield Al₄C₃. This conversion limits the mechanical applications of SiC, because Al₄C₃ is more brittle than SiC.

4 Al + 3 SiC → Al₄C₃ + 3 Si

In aluminium-matrix composites reinforced with silicon carbide, the chemical reactions between silicon carbide and molten aluminium generate a layer of aluminium carbide on the silicon carbide particles, which decreases the strength of the material, although it increases the wettability of the SiC particles.  This tendency can be decreased by coating the silicon carbide particles with a suitable oxide or nitride, preoxidation of the particles to form a silica coating, or using a layer of sacrificial metal.

An aluminium-aluminium carbide composite material can be made by mechanical alloying, by milling aluminium powder with graphite particles.

Occurrence :

Small amounts of aluminium carbide are a common impurity of technical calcium carbide. In electrolytic manufacturing of aluminium, aluminium carbide forms as a corrosion product of the graphite electrodes.

In metal matrix composites based on aluminium matrix reinforced with metal carbides (silicon carbide, boron carbide, etc.) or carbon fibers, aluminium carbide often forms as an unwanted product. In case of carbon fiber, it reacts with the aluminium matrix at temperatures above 500 °C; better wetting of the fiber and inhibition of chemical reaction can be achieved by coating it with e.g. titanium boride.

Applications :

Aluminium carbide particles finely dispersed in aluminium matrix lower the tendency of the material to creep, especially in combination with silicon carbide particles.

Aluminium carbide can be used as an abrasive in high-speed cutting tools. It has approximately the same hardness as topaz.

Tungsten carbide :

Tungsten carbide (WC) is an inorganic chemical compound (specifically, a carbide) containing equal parts of tungsten and carbon atoms. Colloquially, tungsten carbide is often simply called carbide. In its most basic form, it is a fine gray powder, but it can be pressed and formed into shapes for use in industrial machinery, tools, abrasives, as well as jewelry. Tungsten carbide is approximately three times stiffer than steel, with a Young's modulus of approximately 550 GPa, and is much denser than steel or titanium. It is comparable with corundum (α-Al₂O₃) or sapphire in hardness and can only be polished and finished with abrasives of superior hardness such as cubic boron nitride and diamond amongst others, in the form of powder, wheels, and compounds.

Chemical properties :

There are two well characterized compounds of tungsten and carbon, WC and tungsten semicarbide, W₂C. Both compounds may be present in coatings and the proportions can depend on the coating method.

WC can be prepared by reaction of tungsten metal and carbon at 1400–2000 °C.  Other methods include a patented fluid bed process that reacts either tungsten metal or blue WO₃ with CO/CO₂ mixture and H₂ between 900 and 1200 °C. Chemical vapor deposition methods that have been investigated include: WC can also be produced by heating WO₃ with graphite in hydrogen at 670 °C following by carburization in Ar at 1000 °C or directly heating WO₃ with graphite at 900°C.

• tungsten hexachloride with hydrogen, as a reducing agent, and methane, as the source of carbon at 670 °C (1,238 °F)

WCl₆ + H₂ + CH₄ → WC + 6 HCl

• reacting tungsten hexafluoride with hydrogen, as reducing agent, and methanol, as source of carbon at 350 °C (662 °F)

WF₆ + 2 H₂ + CH₃OH → WC + 6 HF + H₂O

• At high temperatures WC decomposes to tungsten and carbon and this can occur during high-temperature thermal spray, e.g. high velocity oxygen fuel (HVOF) and high energy plasma (HEP) methods.

Oxidation of WC starts at 500–600 °C. It is resistant to acids and is only attacked by hydrofluoric acid/nitric acid (HF/HNO₃) mixtures above room temperature. It reacts with fluorine gas at room temperature and chlorine above 400 °C (752 °F) and is unreactive to dry H₂ up to its melting point.

Physical properties :

Tungsten carbide is high melting, 2,870 °C (5,200 °F), extremely hard (8.5–9.0 Mohs scale, Vickers hardness number = 2242) with low electrical resistivity (~2×10⁻⁷ Ohm·m), comparable with that of some metals (e.g. vanadium 2×10⁻⁷ Ohm·m).

WC is readily wetted by both molten nickel and cobalt.  Investigation of the phase diagram of the W-C-Co system shows that WC and Co form a pseudo binary eutectic. The phase diagram also shows that there are so-called η-carbides with composition (W,Co)₆C that can be formed and the fact that these phases are brittle is the reason why control of the carbon content in WC-Co hard metals is important.

Structure :

α-WC structure, carbon atoms are gray.

There are two forms of WC, a hexagonal form, α-WC (hP2, space group P6m2, No. 187), and a cubic high-temperature form, β-WC, which has the rock salt structure. The hexagonal form can be visualized as made up of hexagonally close packed layers of metal atoms with layers lying directly over one another, with carbon atoms filling half the interstices giving both tungsten and carbon a regular trigonal prismatic, 6 coordination. From the unit cell dimensionsthe following bond lengths can be determined; the distance between the tungsten atoms in a hexagonally packed layer is 291 pm, the shortest distance between tungsten atoms in adjoining layers is 284 pm, and the tungsten carbon bond length is 220 pm. The tungsten-carbon bond length is therefore comparable to the single bond in W(CH₃)₆ (218 pm) in which there is strongly distorted trigonal prismatic coordination of tungsten.

Applications :

Cutting tools for machining

Sintered tungsten carbide cutting tools are very abrasion resistant and can also withstand higher temperatures than standard high speed steel tools. Carbide cutting surfaces are often used for machining through materials such as carbon steel or stainless steel, as well as in situations where other tools would wear away, such as high-quantity production runs. Because carbide tools maintain a sharp cutting edge better than other tools, they generally produce a better finish on parts, and their temperature resistance allows faster machining. The material is usually called cemented carbide, hardmetal or tungsten-carbide cobalt: it is a metal matrix composite where tungsten carbide particles are the aggregate and metallic cobalt serves as the matrix. Manufacturers use tungsten carbide as the main material in some high-speed drill bits, as it can resist high temperatures and is extremely hard.

Ammunition

Tungsten carbide is often used in armor-piercing ammunition, especially where depleted uranium is not available or is politically unacceptable. W₂C projectiles were first used by German Luftwaffe tank-hunter squadrons in World War II. Owing to the limited German reserves of tungsten, W₂C material was reserved for making machine tools and small numbers of projectiles. It is an effective penetrator due to its combination of great hardness and very high density.

Tungsten carbide ammunition can be of the sabot type (a large arrow surrounded by a discarding push cylinder) or a subcaliber ammunition, where copper or other relatively soft material is used to encase the hard penetrating core, the two parts being separated only on impact. The latter is more common in small-caliber arms, while sabots are usually reserved for artillery use.

Nuclear

Tungsten carbide is also an effective neutron reflector and as such was used during early investigations into nuclear chain reactions, particularly for weapons. A criticality accident occurred at Los Alamos National Laboratory on 21 August 1945 when Harry K. Daghlian, Jr. accidentally dropped a tungsten carbide brick onto a plutonium sphere, causing the subcritical mass to go supercritical with the reflected neutrons.

Surgical instruments

It is also used for making surgical instruments meant for open surgery (scissors, forceps, hemostats, blade-handles, etc.) and laparoscopic surgery (graspers, scissors/cutter, needle holder, cautery, etc.). They are much costlier than their stainless-steel counterparts and require delicate handling, but give better performance.

Jewelry

Tungsten carbide, also called cemented carbide, has become a popular material in the bridal jewelry industry due to its extreme hardness and high resistance to scratching.

CHEMISTRY OF CHALCOGENS

DEFINITION OF CHALCOGENS :

The chalcogens are the chemical elements in group 16 (old-style: VIB or VIA) of the periodic table. This group is also known as the oxygen family. It consists of the elements oxygen (O), sulfur (S), selenium (Se), tellurium (Te), the radioactive element polonium (Po), and the synthetic element ununhexium (Uuh).
Although all group 16 elements of the periodic table, including oxygen are defined as chalcogens, oxygen and oxides are usually distinguished from chalcogens and chalcogenides. The term chalcogenide is more commonly reserved for sulfides, selenides, and tellurides, rather than for oxides^] Binary compounds of the chalcogens are called chalcogenides (rather than chalcides; however, this breaks the pattern of halogen/halide and pnictogen/pnictide).
Although the word "chalcogen" is literally taken from Greek words being "copper-former", the meaning is more in line with "copper-ore former" or more generally, "ore-former". These electronegative elements are strongly associated with metal-bearing minerals, where they have formed water-insoluble compounds with the metals in the ores.
Properties of the Group VIA Elements

Element	Symbol	Electron Configuration	Usual Oxidation State	Radius/pm
				Covalent	Ionic (X2-)
Oxygen	O	[He]2s22p4	-2	66	140
Sulfur	S	[Ne]3s23p4	+6, +4, -2	104	184
Selenium	Se	[Ar]4s23d104p4	+6, +4, -2	117	198
Tellurium	Te	[Kr]5s24d105p4	+6, +4, -2	135	221

THE CHALCOGEN

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OXYGEN

The Aurora Borealis: Excited oxygen atoms emit green light.

Classification:	Oxygen is a chalcogen and a nonmetal
Color:	colorless
Atomic weight:	15.9994
State:	gas
Melting point:	-218.3 oC, 54.8 K
Boiling point:	-182.9 oC, 90.2 K
Shells:	2,6
Electron configuration:	1s2 2s2 2p4
Density @ 20oC:	0.001429 g/cm3
Atomic volume:	14.0 cm3/mol
Structure:	bcc: body-centered cubic when solid

8 O 16.00

The chemistry of respiration: Lavoisier carries out an experiment to study the oxygen content of air exhaled from a man's lungs. Lavoisier's wife Marie-Anne makes notes. She also created the engraving from which this image was taken.

Oxygen cylinders.

Discovery of Oxygen :

Author: Dr. Doug Stewart

Oxygen was discovered in 1774 by Joseph Priestley in England and two years earlier, but unpublished, by Carl W. Scheele in Sweden.
Scheele heated several compounds including potassium nitrate, manganese oxide, and mercury oxide and found they released a gas which enhanced combustion.
Priestley heated mercury oxide, focusing sunlight using a 12-inch 'burning lens' - a very large magnifying glass - to bring the oxide to a high temperature. Priestley's lens was smaller than the enormous one used by Antoine Lavoisier in his investigation of carbon.
Totally unexpectedly, the hot mercury oxide yielded a gas that made a candle burn five times faster than normal. Priestley wrote: "But what surprised me more than I can well express was that a candle burned in this air with a remarkably vigourous flame. I was utterly at a loss how to account for it."
In addition to noticing the effect of oxygen on combustion, Priestley later noted the new gas's biological role. He placed a mouse in a jar of oxygen, expecting it would survive for 15 minutes maximum before it suffocated. Instead, the mouse survived for a whole hour and was none the worse for it.
Antoine Lavoisier carried out similar experiments to Priestley's and added to our knowledge enormously by discovering that air contains about 20 percent oxygen and that when any substance burns, it actually combines chemically with oxygen.
Lavoisier also found that the weight of the gas released by heating mercury oxide was identical to the weight lost by the mercury oxide, and that when other elements react with oxygen their weight gain is identical to the weight lost from the air.
This enabled Lavoisier to state a new fundamental law: the law of the conservation of matter; "matter is conserved in chemical reactions" or, alternatively, "the total mass of a chemical reaction's products is identical to the total mass of the starting materials."
In addition to these achievements, it was Lavoisier who first gave the element its name oxygen. ⁽
The word oxygen is derived from the Greek words 'oxys' meaning acid and 'genes' meaning forming.
Before it was discovered and isolated, a number of scientists had recognized the existence of a substance with the properties of oxygen:
In the early 1500s Leonardo da Vinci observed that a fraction of air is consumed in respiration and combustion.
In 1665 Robert Hooke noted that air contains a substance which is present in potassium nitrate [potassium nitrate releases oxygen when heated,] and a larger quantity of an unreactive substance [which we call nitrogen].
In 1668 John Mayow wrote that air contains the gas oxygen [he called it nitroarial spirit], which is consumed in respiration and burning.
Mayow observed that: substances do not burn in air from which oxygen is absent; oxygen is present in the acid part of potassium nitrate [i.e., in the nitrate - he was right!]; animals absorb oxygen into their blood when they breathe; air breathed out by animals has less oxygen in it than fresh air.

Appearance & Characteristics :

Watch steel melt when charcoal (carbon) burns in liquid oxygen. (Liquid oxygen is much more concentrated than the gas. Higher concentrations lead to faster reaction rates.)

Liquid oxygen is pale blue and paramagnetic. Watch it stick to a magnet. (3 min 15 secs and onwards.)

Harmful effects:
O₂ is non-toxic under normal conditions. However, exposure to oxygen at higher than normal pressures, e.g. scuba divers, can lead to convulsions. Ozone (O₃) is toxic and if inhaled can damage the lungs.

Characteristics:
Oxygen in its common form (O₂) is a colorless, odorless and tasteless diatomic gas. Oxygen is extremely reactive and forms oxides with nearly all other elements except noble gases.

Oxygen dissolves more readily in cold water than warm water. As a result of this, our planet's cool, polar oceans are more dense with life than the warmer, tropical oceans.

Liquid and solid oxygen are pale blue and are strongly paramagnetic.

Ozone (O₃), another form (allotrope) of oxygen, occurs naturally in the Earth's upper atmosphere. It is made by the action of ultraviolet light on O₂. Ozone shields us from much of the harmful ultraviolet radiation coming from the sun. In Earth's early atmosphere, before oxygen and hence ozone levels were sufficiently high, the ultraviolet radiation reaching our planet's surface would have been lethal to many organisms.⁽⁵⁾

The reaction with oxygen is one of the critera we use to distinguish between metals (these form basic oxides) and non-metals (these form acidic oxides).

Uses of Oxygen :

The major commercial use of oxygen is in steel production. Carbon impurities are removed from steel by reaction with oxygen to form carbon dioxide gas.

Oxygen is also used in oxyacetylene welding, as an oxidant for rocket fuel, and in methanol and ethylene oxide production.

Plants and animals rely on oxygen for respiration. Pure oxygen is frequently used to help breathing in patients with respiratory ailments.

Abundance & Isotopes :

Abundance earth's crust: 46 % by weight, 60 % by moles
Abundance solar system: 9,000 ppm by weight, 700 ppm by moles
Cost, pure: $0.3 per 100g
Cost, bulk: $0.02 per 100g

Laboratory electrolysis of water. Electrical energy is used to split water. Watch out for the different ways the two gases are collected.

Source:
Oxygen is the most abundant element in the Earth's crust, accounting for almost half of it by mass. More than half of the atoms in the Earth's crust are oxygen atoms. About 86 percent of the mass of Earth's oceans is oxygen - mainly in the form of water.

Oxygen is the third most common element in the Universe, behind hydrogen and helium. It is obtained commercially from liquefied air separation plants. It can be prepared in the laboratory by electrolysis of water.

Isotopes: 13 whose half-lives are known, with mass numbers 12 to 24. Of these, three are stable: ¹⁶O, ¹⁷O and ¹⁸O.

Energies                                                                                                               Specific heat capacity: 0.918 J g-1 K-1 Heat of fusion: 0.444 kJ mol-1 of O2 1st ionization energy: 1313.9 kJ mol-1 3rd ionization energy: 5300.3 kJ mol-1 Heat of atomization: 249 kJ mol-1 Heat of vaporization: 6.82 kJ mol-1 of O2 2nd ionization energy: 3388.2 kJ mol-1 Electron affinity: 140.97875 kJ mol-1

Conductivity Thermal conductivity: 0.02583 W m-1 K-1 Electrical conductivity: S cm-

Radius Atomic radius: 60 pm Ionic radius (2+ ion): pm Ionic radius (2- ion): 126 pm Ionic radius (1+ ion): pm Ionic radius (3+ ion): pm Ionic radius (1- ion): pm

SULFUR

Group 6A Elements
Oxygen	O2
Sulfur	S6
Selenium	Se
Tellurium	Te
Polonium	Po

The group 6A elements are listed in the Table on the side here. This goup of elements are intimately related to our lives. We need oxygen all the time throughout our lives. Did you know that sulfur is also one of the essential elements of life. It is responsible for some of the protein structures in all living organisms. Many industries utilize sulfur, but emission of sulfer compounds is often seen more as a problem than the natural phenomenon. The matallic properties increase as the atomic number increases. The element polonium has no stable isotopes, and the isotope with mass number 209 has the longest half life of 103 years.

Properties of oxygen are very different from other elements of the group, but they all have 2 electons in the outer s orbital, and 4 electrons in the p orbitals, usually written as

s²p⁴

The trends of their properties in this group are interesting. Knowing the trend allows us to predict their reactions with other elements. Most trends are true for all groups of elements, and the group trends are due mostly to the size of the atoms and number of electrons per atom.
The trends are described below:

1. The metallic properties increase in the order oxygen, sulfur, selenium, tellurium, or polonium. Polonium is essentially a metal. It was discovered by M. Curie, who name it after her native country Poland.
2. Electronegativity, ionization energy (or ionization potential IP), and electron affinity decrease for the group as atomic weight increases.
3. The atomic radii and melting point increase.
4. Oxygen differs from sulfur in chemical properties due to its small size. The differences between O and S are more than the differences between other members.

Sulfur - a commodity

Sulfuris recovered by the Frasch process. This process has made sulfur a high purity (up to 99.9 percent pure) chemical commodity in largequantities.

Natural Sources of Sulfur

Most sulfur containing minerals are metal sulfides, and the best known is perhaps pyrite, (FeS₂, known as fools gold because of its golden color). The most common sulfate containing mineral is gypsum, CaSO₄.2H₂O, also known as plaster of paris.

Mining Method - Frasch process

Frasch process force (99.5% pure) sulfur out by using hot water and air. In this process, superheated water is forced down the outer most of three concentric pipes. Compressed air is pumped down the center tube, and a mixture of elemental sulfur, hot water, and air comes up the middle pipe. Sulfur is melted with superheated water (at 170 degrees C under high pressure) and forced to the surface of the earth as a slurry.

Applications

Sulfur is mostly used for the production of sulfuric acid, H₂SO₄. Most sulfur mined by Frasch process is used in industry for the manufacture of sulfuric acid.

Sulfuric acid, the most abundantly produced chemical in the United States, is manufactured by the Contact process.
Most (about 70%) of the sulfuric acid produced in the world is used in the fertilizer industry.
Sulfuric acid can act as a strong acid, a dehydrating agent, and an oxidizing agent. It's applications use these properties.
Sulfur is an essential element of life in sulfur-containing proteins.

Elemental Sulfur

Rhombic and monoclinic sulfur are known as allotropes. The crystals of these have the molecules S₈. In these molecules, S form two S-S bonds. The lone pairs of electrons make the S-S-S bend (108 deg), resulting in S₈ having the shape of a crown.

At 298 K, rhombic sulfur is stable, whereas at at 368 K, monoclinc sulfur is formed. The latter is meta-stable at room temperature for some time. In sulfur vapor, S₈, S₆, and S₂ molecules are present.
What happens at when the solid sulfur melts? The S₈ molecules bread up. When suddenly cooled, long chain molecules are formed in the plastic sulfur which, behave as rubber. Plastic sulfur transform into rhombic sulfur over time.

Reactions of Sulfur

Reading the following reactions, figure out and notice the change of the oxidation state of S in the reactants and products. Common oxidation states of sulfur are -2, 0, 4, and 6.
Sulfur (brimstone, stone that burns) reacts with O₂ giving a blue flame:

S + O₂ = SO₂

SO₂ is produced whenever metalsulfide is oxidized. It is recovered and oxidized further to give SO₃, for production of H₂SO₄. SO₂ reacts with H₂S to form H₂O and S.

2 SO₂ + O₂ = 2 SO₃
SO₃ + H₂O = H₂SO₄ <- a valuable commodity
SO₃ + H₂SO₄ = H₂S₂O₇ <- pyrosulfuric acid

Sulfur reacts with sulfite ions in solution to form thiosulfate,

S + SO₃^2- = S₂O₃^2-,

but the reaction is reversed in an acidic solution.

Sulfuric Acid

Sulfuric acid is produced by the contact process in three steps:

       +O2      +H2SO4         +H2O

  SO2  -->  SO3 ----->  H2S2O7  --->  H2SO4

Applications of sulfuric acid

1. as a strong acid for making HCl and HNO₃.
2. as an oxidizing agent for metals.
3. as a dehydrating agent.
4. for manufacture of fertilizer and other commodities.

Hydrogen Sulfide H₂S

hydrogen sulfide, H₂S is a diprotic acid. The equilibria below.

H₂S = HS^- + H⁺
HS^- = S^2- + H⁺

have been discussed in connection with Polyprotic Acids

Structures of Some Sulfur Compounds

In the DOS version, a Demonstration shows you the rotation of S₈, H₂S, SO₂, SO₃, SF₆, etc. Draw the molecular structures for these substance yourself, so that you will get some sense about the beauty of molecules.

SELENIUM

34 Se 78.96

On average, each brazil nut contains 180 quadrillion selenium atoms. That's 1.8 x 10¹⁷ Se atoms.

Classification:	Selenium is a chalcogen and a nonmetal
Color:	gray or red (crystalline), black or
	red (amorphous)
Atomic weight:	78.96
State:	solid
Melting point:	220 oC, 493 K
Boiling point:	685 oC, 958 K
Shells:	2,8,18,6
Electron configuration:	[Ar] 3d10 4s2 4p4
Density @ 20oC:	4.79 g/cm3
Atomic volume:	16.45 cm3/mol
Structure:	long, helical chains (crystalline hexagonal), Se8
	rings (crystalline monoclinic)
Hardness:	2.0 mohs

Pyrites, shown in the image, are mainly iron sulfide. The 1817 discovery of selenium was in sulfur extracted from pyrites. Photo by Aram Dulyan.

Discovery of Selenium

Selenium lies beneath sulfur in Group 16 of the periodic table. The chemical behavior and reactions of these elements are similar.

It is possible selenium was first observed in about the year 1300 by the alchemist Arnold of Villanova.

Villanova lived from about 1235 to about 1310 and was trained in medicine at the Sorbonne in Paris, becoming physician to Pope Clement V. In the book Rosarium Philosophorum he describes red sulfur or 'sulfur rebeum' which had been left behind in an oven after native sulfur had been vaporized. This may have been one of selenium's red colored allotropes. 

There is no more to be said about selenium's discovery until 500 years later, in 1817. In this year, the eminent Swedish chemist Jöns J. Berzelius had his attention drawn to a red deposit left behind after sulfur was burned in a sulfuric acid factory. ⁽⁴⁾

The factory was actually part owned by Berzelius with his friend the chemist Johann Gahn. ⁽⁵⁾

Writing about the deposit in September 1817, Berzelius informed his friend in London, Dr. Marcet, that the deposit contained the (already known) element tellurium.

In February 1818, however, he let Marcet know he had changed his mind, and told him of his discovery of a new element:

"...what Mr. Gahn and I took for tellurium is a new substance, endowed with interesting properties. This substance has the properties of a metal, combined with that of sulfur to such a degree that one would say it is a new kind of sulfur. The similarity to tellurium has given me occasion to name the new substance selenium." ⁽⁶⁾

To explain Berzelius's name for the new element a little more: 'Tellus' in Latin means earth goddess. Berzelius took selenium from the Greek word 'Selene', meaning moon goddess. (Tellurium had been given its name in 1799 by the German chemist Martin Klaporth, who wrote, "No single element was yet named after the Earth. It needed to be done!) ⁽⁷⁾

Appearance & Characteristics

A look at selenium and its compounds.

Allotropes of selenium.

Harmful effects:
Elemental selenium's oral LD₅₀ (the single dose needed to kill 50% of those exposed) is 6700 mg kg^-1 in rats; this is similar to ethanol, which is 7000 mg kg^-1. These levels are classed as non-toxic. Selenium's legal airborne permissible exposure limit (PEL) is 0.2 mg m^-3 averaged over an 8-hour shift. The EPA describes selenium as not classifiable for human carcinogenicity. Selenium sulfide is a probable carcinogen. Many of selenium's compounds, such as selenates and selenites, are highly toxic. Hydrogen selenide gas (SeH₂) is selenium's most acutely toxic compound.

Characteristics:
Selenium exists in several allotropic forms. The most stable form, crystalline hexagonal selenium, is metallic gray. Crystalline monoclinic selenium is a deep red color. Amorphous selenium is red in powder form and is black in vitreous form.

Gray crystalline 'metallic' selenium conducts electricity better in the light than in the dark (photoconductive) and it can convert light directly into electricity (photovoltaic).

In the same way as sulfur forms sulfides, sulfates, and sulfites, selenium combines with metals and oxygen to form selenides, (such as zinc selanide, ZnSe), selenates, (such as calcium selenate, CaSeO₄), and selenites (such as silver selenite, Ag₂SeO₃).

Although hydrogen selenide gas (SeH₂) is highly toxic, it's unlikely you'll hang around long enough to be poisoned; it has a disgusting smell. Oliver Sacks said, "Hydrogen selenide, I decided, was perhaps the worst smell in the world." 

Uses of Selenium

Selenium is used in the glass industry to decolorize glass and to make red-colored glasses and enamels.

It is used as a catalyst in many chemical reactions.

Selenium is used in solar cells and photocells - in fact the first solar cell was made using selenium. It is also used as a photographic toner.

Selenium is used with bismuth in brasses and as an additive to stainless steel. When selenium is added to iron and copper based metals it improves their machinability.

Selenium sulfide is used in anti-dandruff shampoos.

Despite the toxicity of its compounds, selenium is also an essential trace element for humans and other animals. Without it, the enzyme glutathione peroxidase (GPX), which protects against oxidative damage in cells, could not function. Abnormally low selenium in the diet might increase the risk of cancer. Abnormally high levels of selenium compounds can lead to selenium poisoning.

Plants do not appear to need selenium, but they do need sulfur. When selenium is present in soils, it is used by plants as if it were sulfur, introducing selenium into food chains. In soils with low sulfur content, some plants can have high levels of selenium compounds. Animals that eat these plants may suffer ill-health.

Selenium deficiency in animals can lead to slow growth and reproductive dysfunction.

Abundance & Isotopes

Abundance earth's crust: 50 parts per billion by weight, 10 parts per billion by moles
Abundance solar system: parts per billion by weight, part per billion by moles
Cost, pure: $61 per 100g
Cost, bulk: $5.30 per 100g

Source: Selenium occasionally occurs free in nature, but more often occurs as selenides of iron, lead, silver, or copper. Commercially, selenium is obtained mainly from anode mud waste produced in the electrolytic refining of copper. Brazil nuts are the richest known dietary source of selenium.

Isotopes: Selenium has 24 isotopes whose half-lives are known, with mass numbers 67 to 91. Of these, five are stable: ⁷⁴Se, ⁷⁶Se, ⁷⁷Se, ⁷⁸Se and ⁸⁰Se.

Energies Specific heat capacity: 0.32 J g-1 K-1 Heat of fusion: 6.694 kJ mol-1 1st ionization energy: 940.9 kJ mol-1 3rd ionization energy: 2973.7 kJ mol-1 Heat of atomization: 227 kJ mol-1 Heat of vaporization : 26.32 kJ mol-1 2nd ionization energy: 2044.5 kJ mol-1 Electron affinity: 194.97 kJ mol-1 Reactions & Compounds Reaction with air: vigorous, w/ht ⇒ SeO2 Reaction with 15 M HNO3: mild , ⇒ H2SeO3, NOx Oxide(s): SeO2 Hydride(s): SeH2 Reaction with 6 M HCl: none Reaction with 6 M NaOH: Chloride(s): Se2Cl2, Se4Cl16 Conductivity Thermal conductivity: 0.52 W m-1 K-1 Electrical conductivity: 8 x 106 S m-1

Oxidation & Electrons Shells: 2,8,18,6 Minimum oxidation number: -2 Min. common oxidation no.: -2 Electronegativity (Pauling Scale): 2.55 Electron configuration: [Ar] 3d10 4s2 4p4 Maximum oxidation number: 6 Max. common oxidation no.: 6 Polarizability volume: 3.8 Å3 Radius Atomic radius: 119 pm Ionic radius (2+ ion): pm Ionic radius (2- ion): 184 pm Ionic radius (1+ ion): pm Ionic radius (3+ ion): pm Ionic radius (1- ion): pm

TELLURIUM

127.6

Crystalline tellurium.

Classification:	Tellurium is a chalcogen and a metalloid
Color:	silvery
Atomic weight:	127.60
State:	solid
Melting point:	450 oC, 723 K
Boiling point:	990 oC, 1263 K
Shells:	2,8,18,18,6
Electron configuration:	[Kr] 4d10 5s2 5p4
Density @ 20oC:	6.24 g/cm3
Atomic volume:	20.5 cm3/mol
Structure:	parallel chains
Hardness:	2.3 mohs

Discovery of Tellurium

Tellurium was discovered by Baron Franz Muller von Reichenstein in 1783.

Martin H. Klaproth isolated the element and named it in 1798.

The element name comes from the Latin word 'tellus' meaning Earth.

Appearance & Characteristics

Hubble Telescope Wide Field Camera 3. The crystalline photosensitive surface of the camera's near-infrared detector is composed of mercury, cadmium and tellurium (HgCdTe). (NASA)

Harmful effects:
Tellerium is very toxic and teratogenic (can cause harm to developing embryos). Exposure to as little as 0.01 mg/m² or less in air leads to "tellurium breath", which has a garlic-like odor.

Characteristics:
Tellurium is a rare, silvery-white, brittle, lustrous metalloid. It burns in air with a greenish-blue flame and forms tellurium dioxide (TeO₂). Tellurium is a semiconductor material and is slightly photosensitive. It forms many compounds corresponding to those of sulfur and selenium, the elements above it in the periodic table. Tellurium has radioactive isotopes and is the lightest element to exhibit alpha decay.

Uses of Tellurium :

Tellurium is alloyed with copper and stainless steel to make these metals more workable. It is added to lead to decreases the corrosive action of sulfuric acid and to improve its strength and hardness. Tellurium is used as a coloring agent in ceramics. Tellurium is also used in the electronic industry, for example with cadmium and mercury to form photosensitive semiconductors. It is used in vulcanizing rubber and in catalysts for petroleum cracking and in blasting caps for explosives.

Abundance & Isotopes

Abundance earth's crust: 1 part per billion by weight, 0.2 parts per billion by moles
Abundance solar system:
Cost, pure: $24 per 100g
Cost, bulk: $0.44 per 100g

Source: Tellurium is sometimes found free in nature. More commonly, it is found combined with metals, such as in the minerals calaverite (gold telluride, AuTe₂) and sylvanite (silver-gold telluride). Commercially, tellurium is obtained as a byproduct of electrolytic copper refining.

Isotopes: Tellurium has 33 isotopes whose half-lives are known, with mass numbers 106 to 138. Of these, five are stable: ¹²⁰Te, ¹²²Te, ¹²⁴Te, ¹²⁵Te and ¹²⁶Te.

Energies Specific heat capacity: 0.20 J/gK Heat of fusion: 17.490 kJ mol-1 1st ionization energy: 869.2 kJ mol-1 3rd ionization energy: 2697.7 kJ mol-1 Heat of atomization: 197 kJ mol-1 Heat of vaporization: 52.550 kJ mol-1 2nd ionization energy: 1794.6 kJ mol-1 Electron affinity: 190.16 kJ mol-1

Oxidation & Electrons Shells: 2,8,18,18,6 Minimum oxidation number: -2 Min. common oxidation no.: 0 Electronegativity (Pauling Scale): 2.1 Electron configuration: [Kr] 4d10 5s2 5p4 Maximum oxidation number: 6 Max. common oxidation no.: 6 Polarizability volume: 5.5 Å3

Reactions & Compounds Reaction with air: mild, w/ht ⇒ TeO2 Reaction with 15 M HNO3: mild , ⇒ Te(IV) Oxide(s): TeO2, TeO3 Hydride(s): TeH2 (hydrogen telluride) Reaction with 6 M HCl: none Reaction with 6 M NaOH: none Chloride(s): Te2Cl, Te3Cl2, Te4Cl16

Conductivity Thermal conductivity: 3 W m-1 K-1 Electrical conductivity: 0.0002 x 106 S m-1 Radius Atomic radius: 142 pm Ionic radius (2+ ion): pm Ionic radius (2- ion): 207 pm Ionic radius (1+ ion): pm Ionic radius (3+ ion): 90 pm Ionic radius (1- ion): pm

POLONIUM 84 Po (209)

Marie Curie in 1883, 16 years old. Polonium was the first element she discovered, 15 years later.

Classification:	Polonium is a chalcogen and a metalloid
Color:	silvery-gray
Atomic weight:	(209), no stable isotopes
State:	solid
Melting point:	254 oC, 527 K
Boiling point:	960 oC, 1233 K
Shells:	2,8,18,32,18,6
Electron configuration:	[Xe] 4f14 5d10 6s2 6p4
Density @ 20oC:	9.4 g/cm3
Atomic volume:	22.23 cm3/mol
Structure:	simple cubic
Hardness:

Discovery of Polonium :

Author: Dr. Doug Stewart

Polonium was the first element Marie and Pierre Curie discovered.

They discovered polonium and then radium in 1898, while working in Paris, investigating radioactivity in pitchblende (uranium oxide).

At the time of the discovery they wrote: "We thus believe that the substance that we have extracted from pitchblende contains a metal never known before, akin to bismuth in its analytic properties. If the existence of this new metal is confirmed, we suggest that it should be called polonium after the name of the country of origin of one of us."

In accordance with the Curies' wishes, polonium is named after Poland, the country of Marie Curie's birth.

The dangers of working with radioactive elements were not known when the Curies' made their discoveries. Their laboratory notebooks from this time are so radioactive that they are now stored in a lead-lined case. ⁽¹⁾

Appearance & Characteristics

Part of the uranium decay series. Three isotopes of polonium are produced in nature, either by the decay of radon gas itself or by the decay of atoms resulting from radon's decay. 

Harmful effects:
Polonium is harmful both through its chemical toxicity and its radioactivity. Polonium-210 is an alpha emitter. As such it is very hazardous if swallowed or inhaled. Exposure to polonium increases the risk of getting various cancers.

Characteristics:
Polonium is a rare, silvery-gray, radioactive low-melting metalloid. Polonium readily reacts with dilute acids, but only slightly with alkalis. All of its isotopes are radioactive. ²¹⁰Po emits a blue glow, as the air around it is excited by the decay products. 1 gram of Po emits as many alpha particles as 5 kilograms of radium. The energy released by polonium's alpha decay is considerable and heats the volume around it. The energy released is so large (140 W/g) that a capsule containing about half a gram reaches a temperature above 500  ^oC.

Uses of Polonium :

Polonium is used to eliminate static electricity produced during processes such as rolling paper, wire and sheet metal. However, beta decay sources are more commonly used as they are less dangerous. ²¹⁰Po can be used as an atomic heat source but because of the isotope's short half-life (138.4 days), it doesn't provide power for long-term uses. Polonium is also used in anti-static brushes to eliminate dust on photographic film. It is sealed in brushes to control the radioactive emissions.

Abundance & Isotopes

Abundance earth's crust: Of the order of 1 part per quadrillion.
Abundance solar system: negligible
Cost, pure: per 100g
Cost, bulk: per 100g

Source: Polonium is a very rare element due to the short half-life of all its isotopes. It is found in uranium ores in minute quantities. It can be obtained by bombarding natural bismuth, ²⁰⁹Bi , with neutrons to give ²¹⁰Bi, which then decays to ²¹⁰Po via β decay. Approximately 100 g of polonium is synthesized each year.

Isotopes: Polonium has 29 isotopes whose half-lives are known, with mass numbers 190 to 218. None are stable. The most stable isotope is ²⁰⁹Po, with a half-life of 102 years.

Energies Specific heat capacity: 0.12 J g-1 K-1 0.12 J g-1 K-1 Heat of fusion: 13 kJ mol-1 1st ionization energy: 812 kJ mol-1 3rd ionization energy: kJ mol-1 Heat of atomization: 142 kJ mol-1 Heat of vaporization: 120 kJ mol-1 2nd ionization energy: kJ mol-1 Electron affinity: 180 kJ mol-1

Oxidation & Electrons Shells: 2,8,18,32,18,6 Minimum oxidation number: -2 Min. common oxidation no.: -2 Electronegativity (Pauling Scale): 2.0 Electron configuration: [Xe] 4f14 5d10 6s2 6p4 Maximum oxidation number: 6 Max. common oxidation no.: 4 Polarizability volume: 6.8 Å3

Reactions & Compounds Reaction with air: mild, ⇒ PoO2 Reaction with 15 M HNO3: Oxide(s): PoO2, PoO2 Hydride(s): PoH2 Reaction with 6 M HCl: mild, ⇒ PoCl2 Reaction with 6 M NaOH: none Chloride(s): PoCl2

Radius Atomic radius: 190 pm Ionic radius (2+ ion): pm Ionic radius (2- ion): pm Ionic radius (1+ ion): pm Ionic radius (3+ ion): pm Ionic radius (1- ion): pm Conductivity Thermal conductivity: 0.2 W m-1 K-1 Electrical conductivity: 0.7 x 106 S m-1

Uuh

In studies reported by the Joint Institute for Nuclear Research in Dubna, Russia in 2000/2001, atoms of element 116 were observed to decay following the complete fusion of Ca-48 and Cm-248. No independent confirmation of the experiments has been claimed (an earlier claim to the discovery of elements 116 and 118 at Berkeley was withdrawn)