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Oxygen was discovered in 1774 by Joseph Priestley in England and two years earlier, but unpublished, by Carl W. Scheele in Sweden.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

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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)