Tellurium

Tellurium was discovered in 1782. It is semi-metallic in nature and present in the oxygen group of periodic table.  Tellurium used as semiconductor and photosensitive material.

History and Discovery

In 1700, scientist found a new element in various ores who had both metallic and non-metallic properties. They called that element aurum paradoxum or metallum problematum, meaning problem metal [1]. Officially, it was discovered by Franz Joseph Muller von Reichenstein in 1782. He sent sample to Torbern Bergman in Uppsala, Sweden but he died. In 1798, Martin Heinrich Klaproth confirmed the existence of that element and named that new element tellurium. The word tellurium has been derived from Latin word ‘tellus’ that means ‘Earth’ [2]. In 1960, tellurium was widely used in steel alloys and thermoelectric applications.

Tellurium

Periodic Table ClassificationGroup 16
Period 5
State at 20CSolid
ColorSilvery lustrous gray
Electron Configuration[Kr] 4d10 5s2 5p4
Electron Number52
Proton Number52
Electron Shell2, 8, 18, 18, 6
Density6.24 g.cm-3 at 20°C
Atomic number52
Atomic Mass127.60 g.mol -1
Electronegativity according to Pauling2.10

Occurrence

Tellurium is not very common in the earth crust.  Tellurium is present in about 1 part per billion in the Earth’s crust. The high atomic number and the formation of halides caused it to be lost in the space as a gas during the hot nebular formation of the planet.  Like other metals such as copper, lead, silver or gold, it is often found in uncombined form but mostly present in mineral with gold. It is obtained as a by-product during the refinement of lead or copper. It is mostly found in the ores sylvanite (AgAuTe4 Silver Gold Telluride), calaverite (AuTe2 Gold Telluride) and krennerite (AuTe2 Orthorhombic gold telluride) [3]. The largest producers of tellurium are Canada, Japan and the United States.

Physical Characteristics

Tellurium is semimetallic element present in the oxygen group of the periodic table. It has two allotropes: crystalline and amorphous. Crystalline tellurium looks silvery white and have a metallic luster. Amorphous form exists in black brown powder, which is prepared by precipitating a solution of telluric acid. In molten form, tellurium is corrosive to copper, iron and stainless steel. It has high melting point about 449.51oC and boiling point is 988oC. Its chemical symbol is Te. Tellurium atomic number is 52 and atomic mass is about 127.60 g/mol.

Chemical Characteristics

Tellurium belongs to chalcogen (oxygen family).  It has both properties of metal and nonmetal. When burned in air, it forms tellurium dioxide and gives a greenish blue flame. It is unaffected by water and hydrochloric acid but is dissolved in nitric acid.  It is also treated with concentrated sulfuric acid. Tellurium exist in various oxidation states, -including 2, +2, +4 and +6. Reduction of tellurium metals lead to the formation of tellurides (anion Te2-) and polytellurides.  The halogens forms halides in +2 and +4 states.  Only fluoride forms halides in +6 oxidation sate but others form halides at +2 and +4 state. In compounds, tellurium adopts a polymeric structure consisting of zig zag chain of tellurium atoms.

Significance and Uses

  • Tellurium is widely used in metallurgy (separate metals from their ores) in iron, stainless steel copper and lead alloys.
  • It has high efficiencies for solar cell electric power generators.
  • Tellurium is added to lead to improve its strength and resistance to corrosion.
  • Tellurium is added in rubber which is helpful in curing process, less susceptible to aging and soften the normal rubber.
  • Tellurium is the primary ingredient of blasting caps, explosive caps of TNT detonators etc.
  • Tellurium is used as pigment for glass and ceramics.
  • Bismuth telluride and lead telluride are semiconductor and used in thermoelectric devices for providing electricity or for cooling purpose.
  • Tellurium is used as a catalyst for petroleum cracking, a process of converting larger hydrocarbons to smaller and simpler units, lighter hydrocarbons.
  • Tellurite agar is used to identify the members of Corynebacterium (genus of bacteria, gram + and aerobic) genus.
  • Tellurium is doped with gold and copper in semiconductor applications.

Health effects

In body, tellurium is partly metabolized in the form of dimethyl telluride, which has a garlic like odor. The exhalation of this gas can be indication of tellurium exposure. Prolonged exposure to tellurium may cause abdominal pain, constipation and vomiting.

Isotopes of Tellurium

Naturally occurring tellurium has eight isotopes, out of which six are stable. The stable isotopes are 120Te, 122Te, 123Te, 124Te, 125Te and 126Te. The other two are slightly radioactive 128Te and 130Te. 128Te have long half-life of 2.2 x 1024 years. Tellurium have thirty artificial radioisotopes having atomic masses ranging from 105 to 142.

REFERENCES

[1].  https://www.chemicool.com/elements/tellurium.html

[2]. https://www.britannica.com/science/tellurium

[3]. https://education.jlab.org/itselemental/ele052.html

Nitrogen

Nitrogen is a chemical element with symbol N and atomic number 7. It was first discovered and isolated by Scottish physician Daniel Rutherford in 1772.

History and Discovery

Nitrogen is one of the prehsitric elements. And its use as ammonium chdirde has been known to Herodotus, Middle Ages civilizations and Alchemists. It was known as aqua fortis the strong water. And mixture of hydrochloric and nitric acid was formed that has the ability to dissolve the Nobel metal and the king of metals, gold. The solution was termed as aqua regia (royal water) [1]. Nitrogen was discovered as a novel element by Daniel Rutherford (1722) and he termed it as the noxious air [2]. He found that it was that component of the air that did not support combustion. During the same time, several other chemists, including Joseph Priestly, Wilhelm Scheele and Henry Cavendish carried out various experiments to discover and identify nitrogen and they termed it as phlogisticated air or burnt air. Earlier, Antoine Lavoisier used the term azote (Greek) meaning no life for nitrogen, which later became choke or to suffocate and the term pnictogens (Greek for choke) was assigned to Group 15 due to nitrogen. In 1970, nitrogene (French word) was given to the element by Antoine Chaptal and in 1974 it became nitrogen in English language. Sodium nitrate and potassium nitrate (then termed as saltpeter) were the earliest known compounds of nitrogen that has various industrial, military and agricultural applications.

Nitrogen

Periodic Table ClassificationGroup 15
Period 2
State at 20CGas
ColorColorless gas
Electron Configuration[He] 2s2 2p3
Electron Number7
Proton Number7
Electron Shell2, 5
Density1.25 g.cm-3 at 20°C
Atomic number7
Atomic Mass14.01 g.mol -1
Electronegativity according to Pauling3.04

Occurrence

Nitrogen is a very common element, both on Earth and in the universe. It is ranked as the 7th most abundant element in the universe and is present in the Solar System and the Milky Way. If forms a distinct surface coverage on Pluto. In diatomic form, N2 gas is the most abundant free element in the Earth’s atmosphere and makes about 78% of it, along with oxygen, carbon dioxide etc. Nitrogen is also prevalent in the living systems, as it is part of various vital components of a living organism, including proteins, amino acids, DNA, RNA and in the energy currency of the cell, ATP. It is considered as the 4th most abundant element in the human body and makes about 3% by mass of the body. A specialized cycle keeps the natural balance of nitrogen between the biotic and abiotic component (atmosphere, organic compounds and biosphere) of the ecosystem in order.

Physical Characteristics

Nitrogen is a colorless and odorless gas. Liquid nitrogen resembles water in its appearance, as it is colorless. Nitrogen is a very light gas and is the lightest gas in its group (Group 15). Molecular nitrogen undergoes liquefaction at -195.79C and freezes at -210 C and acquire a beta hexagonal structural assembly [3]. The alpha phase is another allotropic form of nitrogen that it acquires by arranging in a cubic crystal when exposed to temperature lower than -237C. The density of liquid nitrogen is 0.808 g/mL which is about 80.8% denser than water.

Chemical Characteristics

Nitrogen is very reactive. The triple bonds present in N2 are extremely strong and in fact are the second strongest bonds in elemental chemistry. Except for Nobel gases, nitrogen can react with almost every element in the periodic table and forms nitrides.

Significance and Uses

  • Nitrogen is widely used in the manufacturing of nitrates and ammonia that are considered as the key fertilizers all over the world.
  • Nitrogen is used to make adhesives and glues (in the form of cyanoacrylate).
  • Nitrogen is used in to manufacture high quality stainless steel.
  • Nitrogen is used to make high-strength fabric, that can withstand tearing and wearing.
  • Nitrogen is widely used in pharmaceutical industry for the manufacturing of various drugs including antibiotics, and hypertension controlling drugs (nitroglycerin).
  • Nitrogen is used to inflate tires of aircraft and race cars instead of natural air.
  • Nitrogen is used as a coolant or refrigerant and used for cryopreservation purposes of biological tissues, cells and blood.

Health Hazards

Nitrogen is non-toxic in elemental form and at normal atmospheric pressure. When inhaled in an enclosed space or at high partial pressure, it can prove to be very toxic. Such incidents are common in astronauts and scuba divers, as they are exposed to high levels of nitrogen. Inhalation of nitrogen at a partial pressure more than 4 bar can cause severe tissue damages and mental disorders. Nitrogen can displace oxygen and readily dissolves in body fats and blood and cause decompression sickness, which can sometimes prove to be fatal.

Isotopes of Nitrogen

There are two stable isotope of nitrogen, nitrogen-14 and nitogen-15. Nitrogen-14 is very abundant and makes about 99.6% of naturally occurring nitrogen. There isotopes are produced in the stars. Nitrogen-15 was discovered in 1929 by S.M.  Naude. There are ten artificially produced radioactive isotopes of nitrogen, ranging from nitrogen-12 to nitrogen-23.

REFERENCES

[1]. Greenwood and Earnshaw, pp. 406–07

[2]. Weeks, Mary Elvira (1932). “The discovery of the elements. IV. Three important gases”. Journal of Chemical Education. 9 (2): 215. Bibcode:1932JChEd…9..215W. doi:10.1021/ed009p215.

[3]. Gray, Theodore (2009). The Elements: A Visual Exploration of Every Known Atom in the Universe. New York: Black Dog & Leventhal Publishers. ISBN 978-1-57912-814-2.

Niobium

Niobium was discovered by English chemist Charles Hatchett in 1801, Niobium is a soft, grey, ductile element classified amongst the transition elements of periodic table. It is used in making high-grade structural steel.

History and Discovery

Niobium was discovered in 1801 by Charles Hatchett. Initially, it was named Columbium (with symbol Cb) after Columbia; a state of United States. Later, due to confusion, it was considered another element, known as tantalum, and in 1846, German Chemist Heinrich Rose proposed that it was indeed a novel element and named it Niobium. For almost a century Europe and USA used their own name for this element. Later in 1949 at the 15th Conference of the Union of Chemistry, the name Niobium was chosen as the 41st element of the periodic table [1]. The symbol for Niobium is Nb.

Niobium

Periodic Table ClassificationGroup 5
Period 5
State at 20CSolid
ColorGray metallic
Electron Configuration[Kr] 4d4 5s1
Electron Number41
Proton Number41
Electron Shell2, 8, 18, 12, 1
Density8.57 g.cm-3 at 20°C
Atomic number41
Atomic Mass92.91 g.mol -1
Electronegativity according to Pauling1.6

Occurrence

Niobium is not a very rare element. Niobium can be found in nature only in compound state. Mineral columbite is the main source of this element. It is ranked as the 33rd most common element in Earth’s crust. It is present in abundance at the core or center of the earth. Large deposits of Niobium are found in Brazil and Canada and are the largest producers of niobium mineral concentrates [2]. It is also found in Australia, Russia and Nigeria.

Physical Characteristics

Niobium is a soft, lustrous, ductile transition metal. It is greyish in color and is changed to blue when exposed to air and oxygen. It has a cubic crystalline structure. It becomes a superconductor at very low temperatures [3]. Niobium can produce strong magnetic field. It has a high melting point, 2469°C. Despite being a refractory metal, it has a considerably lower density (i.e. 8.57 g/cm3). Niobium becomes resistant to further corrosion, when a layer of oxide is formed on its surface. Atomic number of niobium is 41.

Chemical Characteristics

Niobium is a reactive metal. It readily reacts with carbon, sulfur, nitrogen, and oxygen. It does not react with acids, at room temperature. However, niobium reacts with hot concentrated acids and alkalis. Niobium is attacked by oxidizing agents. And it readily reacts with halogens.

Significance and Uses

  • Niobium is used to increase the refractive index of glass. This allows optical corrective glasses to be made with lighter and thinner lenses.
  • Niobium is used in making commemorative coins.
  • Niobium is used to make special stainless alloys. About 90% of extracted niobium is used in making high grade structural steel [4].
  • Niobium is used to make various super alloys. These alloys are used in making rockets and jet engines, oil rigs and gas turbines. Grinders and beams for buildings are also prepared consuming alloys made using niobium. These alloys show improved strength at low temperatures. Nozzle for Apollo Service module was also made using niobium alloy [5].
  • Niobium carbide is used in the manufacturing of sharp cutting tools.
  • Niobium is also used in manufacturing of surface acoustic wave devices, which are essential components of electronic components [6].
  • Niobium is used to make superconducting magnets for MRI scanners, nuclear magnetic resonance and particle accelerators [3].
  • Niobium is hypoallergenic and is used in making implant devices and prosthetics.
  • Niobium is used for making wide array of shimmery color jewelry.
  • Niobium is used in as a catalysts for making acrylic acid [7].

Health effects

Niobium in elemental form is non-toxic. People rarely are exposed to compounds containing niobium, but some are harmful and should be avoided. Niobium is considered as a potential fire hazard. Niobium dust is toxic and is irritant to skin and eyes. Apart from this, niobium alloys are tested to be physiologically inert and hence harmless. It does not react with human tissues which is why it is used in surgical implants [8]. When inhaled, it is mainly retained in lungs and bones. Inhalation of nitride or pentoxide of niobium in concentration more than 40 mg/m3 of niobium can lead to severe damage to lungs.

Isotopes of Niobium

There are fourteen isotopes of niobium. Natural niobium has one stable isotope, niobium-93. This isotope is extremely light weight and is considered as the lightest nuclide and is susceptible to undergo spontaneous fission. The most stable radioactive isotope of niobium is niobium-92.

REFERENCES

[1].  Rayner-Canham, Geoff; Zheng, Zheng (2008). “Naming elements after scientists: an account of a controversyFoundations of Chemistry10 (1): 13–18. doi:10.1007/s10698-007-9042-1

[2]. Patel, Zh.; Khul’ka K. (2001). “Niobium for Steelmaking”. Metallurgist. 45 (11–12): 477–480. doi:10.1023/A:1014897029026

[3]. Peiniger, M.; Piel, H. (1985). “A Superconducting Nb3Sn Coated Multicell Accelerating Cavity”. Nuclear Science. 32 (5): 3610–3612. doi:10.1109/TNS.1985.4334443.

[4]. Hillenbrand, Hans-Georg; Gräf, Michael; Kalwa, Christoph (2 May 2001). “Development and Production of High Strength Pipeline Steels” (PDF). Niobium Science & Technology: Proceedings of the International Symposium Niobium 2001 (Orlando, Florida, USA).

[5]. S. T. Scheirer, C. R. Cook, “Development of Columbium Alloy Combinations for Gas Turbine Blade Applications”, (Interim Technical Management Report No. 4, prepared under contract AF 33(615)-67-C-1688 by TRW Inc. August 31, 1968).

[6]. Hebda, Chang (2 May 2001). “Niobium alloys and high Temperature Applications” (PDF). Niobium Science & Technology: Proceedings of the International Symposium Niobium 2001 (Orlando, Florida, USA)

 [7]. Nowak, Izabela; Ziolek, Maria (1999). “Niobium Compounds: Preparation, Characterization, and Application in Heterogeneous Catalysis”. Chemical Reviews. 99 (12): 3603–3624. doi:10.1021/cr9800208.

[8]. C. K. Gupta ,A. K. Suri, Extractive Metallurgy of Niobium, (CRC Press, Boca Raton, Florida, 33431, 1994), 3, 22.

 

Chlorine

Chlorine is one of the most reactive gases and belong to the halogen group of periodic table. It was discovered as a distinct element in 1810.

History and Discovery

Chlorine compound, sodium chloride (table salt) has been known by human civilizations since prehistoric times and evidence of use of rock salt have been found from as early as 3000 BC [1]. Elemental chlorine was isolated for the first time in around 1200, when aqua regia was discovered and was used to dissolve Nobel metals, like gold. Chlorine is released when gold is dissolved in aqua regia, however, it was not identified or studied at that time. Later in 1630, Jan Baptist Helmont proposed that chlorine is a gas, and not an oxide of an element. Later, Swedish chemist, Wilhelm Scheele, put forward a detailed study on the novel compound in 1774 and named it as, dephlogisticated muriatic acid air (muriatic acid was name of hydrochloric acid). The discovery of chlorine as an element rather than a compound was made by Sir Humphry Davy in 1810. The name chlorine has been derived from khloros, that is Greek words for pale green color of chlorine gas.

Chlorine

Periodic Table ClassificationGroup 17
Period 3
State at 20CGas
ColorPale yellow-green gas
Electron Configuration[Ne] 3s2 3p5
Electron Number17
Proton Number17
Electron Shell2, 8, 7
Density3.21 g.cm-3 at 20°C
Atomic number17
Atomic Mass35.45 g.mol -1
Electronegativity according to Pauling3.16

Occurrence

Chlorine is an abundant element. It is the 21st most abundant element in the earth’s crust and makes about 126 ppm of Earth’s crust. Chlorine does not occur in free form in nature. It is found as ionic compounds of chloride in the Earth’s crust. The most common mineral is halite (sodium chloride), while other minerals include sylvite (potassium chloride) and carnallite. The dissolved chloride in sea water makes the biggest natural reserve of chloride, including the Dead Sea (Israel) and Great Salt Lake (Utah) [2]. Among halogens, chlorine is the 2nd most abundant element after fluorine. Elemental chlorine is artificially produced from electrolysis of brine at commercial level. Chloride ions are ubiquitous in all living organisms. Inside the cell, chloride ions function to maintain a balanced environment of negative ions against the positive potassium and sodium ions. Elemental chlorine, in very small quantities is produced during an immune response of white blood cells against bacteria by the oxidation of chloride. Chlorine compounds, chlorofluorocarbons are present in high quantities and damage the ozone layer, eventually increasing the global warming. According to an estimate, around 40 million tons of chlorine gas is made each year from brine.

Physical Characteristics

Chlorine is a non-metal. It is a greenish yellow dense gas and has a pungent smell. The boiling and melting point of chlorine is also intermediate between fluorine and bromine. Chlorine has a diatomic molecular structure, that is why its heat of vaporization quite low, which lead to high volatile nature of chlorine. Chlorine becomes colorless at low temperature (-195C) [3]. Chlorine is a very poor conductor of electricity.

Chemical Characteristics

Chlorine is a very reactive non-metal and has a reactivity that is intermediate between fluorine and bromine. Element chlorine is a strong oxidizing agent, however, among halogens, it is still weaker than fluorine but stronger than bromine. Chlorine can readily react with almost all elements in the periodic table to form binary chlorides, except for oxygen, and the Nobel gases (excluding xenon). Chlorine reacts with hydrogen to form hydrogen chloride (HCl), which has a wide range of industrial and laboratory applications. Chlorine oxides are highly unstable, and all are extremely exothermic. Organic compounds of chlorine, C-Cl bonds are quite common and are considered as a common functional group in organic chemistry. Several organochlorides have been found in human, animals, plants as well as bacteria [4].

Significance and Uses

  • Chlorine is abundantly used making synthetic polymer, polyvinyl chloride (PVC). PVC is used in making a wide range of products, including water pipes, car interiors, vinyl flooring, insulations of wire and blood bags.
  • Chlorine is widely used as a disinfectant as it can kill various kinds of bacteria. It is commonly used as a water disinfectant.
  • Chlorine is used in the manufacturing of paper and in textile industry.
  • Chlorine is used in paints.
  • Chlorine is used in making insecticides and pesticides.
  • Chlorine is widely used in pharmaceutical industry for the manufacturing of large range of medicines, chloroform and disinfectants. However, due to the toxic effects of chlorine, its use has been greatly limited.

Health Hazards

Chlorine is highly toxic. It is used to in various chemical weapons. During the First World War, various chlorine-based weapons were used. Chlorine gas was also used in Iraq bombing, 2007 killing many people. Exposure of chlorine gas lead to irritation of eyes and nose, severe breathing difficulty and suffocation. Chloride ions are essential for maintaining a healthy body, and a daily dose of around 3-6 grams of salt is recommended.

Isotopes of Chlorine

Chlorine has two stable isotopes, chlorine-37 and chlorine-35, where chlorine-35 is more abundant (76%). The radioactive isotope, chlorine-36 is the most stable chlorine.

REFERENCES

[1]. “The earliest salt production in the world: an early Neolithic exploitation in Poiana Slatinei-Lunca, Romania”. Archived from the original on April 30, 2011. Retrieved 2008-07-10.

[2]. Greenwood and Earnshaw, p. 795

[3]. Greenwood and Earnshaw, pp. 800–4

[4]. Gordon W. Gribble (1999). “The diversity of naturally occurring organobromine compounds”. Chemical Society Reviews. 28 (5): 335–46. doi:1039/a900201d.

Arsenic

Arsenic is one of the most ancient element known to human history. It is very toxic and can enter the biological system through contaminated water, soil and air.

History and Discovery

Arsenic has a been known since prehistoric time and in the Bronze age it was used to make alloys with bronze. Arsenic was isolated as a distinct compound by Albertus Magnus in 1250. The word arsenic has been originated from zarnik, that means yellow or golden colored in Persian language and from Greek word arsenikon used for male. Later, the Greek word changed to arsenicum in French and finally the English word arsenic was derived from it. Arsenic has a notorious history in regard to its use as a fatal poison and was commonly termed as the poison of the kings and the king of the poisons [1].

Arsenic

Periodic Table ClassificationGroup 15
Period 4
State at 20CSolid
ColorMetallic grey
Electron Configuration[Ar] 3d10 4s2 4p3
Electron Number33
Proton Number33
Electron Shell2, 8, 18, 5
Density5.72 g.cm-3 at 20°C
Atomic number23
Atomic Mass74.92 g.mol -1
Electronegativity according to Pauling2.18

Occurrence

Arsenic is abundant element and occurs in about 1.5ppm of concentration in the earth’s crust. It is categorized as the 53rd most abundant element on earth. However, it is also present in minute quantities in water and atmosphere. Arsenic can exist in nature in its free elemental, as well as combined form. In combined form, it is primarily present in minerals of sulfur. There are various allotropic forms of arsenic, black, yellow and gray. The gray allotrope of arsenic is the most common and is widely used all over the world in various applications [2]. Arsenic is metabolized by some bacteria and certain animals, including goats, hamsters and chicken intake arsenic as an essential dietary nutrient. The largest producers of arsenic include China, Russia, Belgium, USA and Morocco.

Physical Characteristics

Yellow arsenic is waxy and soft. Both yellow and gray arsenic are highly volatile and unstable. They have very low density (5.727 g/cm3). Black arsenic is brittle and have a glassy appearance. It is poor conductor of electricity.

Chemical Characteristics

When exposed to air it is tarnished and forms a golden-bronze layer on its surface that turns black with passage of time [3]. When arsenic is heated in air, it gives a pungent, garlic like odor as arsenic is oxidized to form arsenic trioxide. Certain compounds of arsenic undergo sublimation when exposed to high temperature, around 614 °C [3]. Arsenic reacts with various metals to form arsenides, and the most common oxidation state of arsenic in compounds is +3 and -3. There are various inorganic compounds of arsenic, including arsenic trioxide which is formed by the oxidation of arsenic in the presence of air and water. Arsenic also readily reacts with halogens to form trihalides and pentahalides, for instance arsenic pentafluoride. There are also many organic compounds of arsenic, such as cacodylic acid, which is formed by the methylation of arsenic trioxide, which is highly pungent and toxic compound.

Significance and Uses

  • The use of arsenic for various products was greatly limited due to the increased knowledge of its toxicity. In 2004, an official ban was applied on the use of chromated copper arsenate (CCA) in US and Europe. CCA was widely used as wood preservative in these countries. However, arsenic is still used in other countries of the world for wood preservation and various other applications.
  • Arsenic is used as a feed additive in swine and poultry industry, to improve the growth of animals.
  • Arsenic is used in various medial purposes, such as drugs for treatment of infection and cancer.
  • Arsenic is used to make alloys with lead to provide strength to lead batteries.

Health Hazards

Arsenic is toxic compound. Its toxicity to biological system and environment is greatly enhanced due to its water solubility. The main source of arsenic toxicity includes weathering of arsenic ores and mineral, and volcanic ash. Traces of arsenic pollution have been found in air, soil and water, from where it can be inhaled and ingested [4]. arsenic water pollution has reached to an alarming level in groundwater in Bangladesh and have according to an estimate, it has affected 57 million people living that region [5].

Isotopes of Arsenic

Natural arsenic has only one stable isotope, arsenic-75 [6]. There are around 33 artificially produced isotopes of arsenic, that range in atomic mass from 60 t0 92. The most stable radioactive isotope is arsenic-73, which as a half-life of 80 days.

REFERENCES

[1]. Vahidnia, A.; Van Der Voet, G. B.; De Wolff, F. A. (2007). “Arsenic neurotoxicity – a review”. Human & Experimental Toxicology. 26 (10): 823–32. doi:1177/0960327107084539. PMID18025055.

[2]. Norman, Nicholas C. (1998). Chemistry of Arsenic, Antimony and Bismuth. Springer. p. 50. ISBN 978-0-7514-0389-3.

[3]. Gokcen, N. A (1989). “The As (arsenic) system”. Bull. Alloy Phase Diagrams. 10: 11–22. doi:10.1007/BF02882166.

[4]. Arsenic. The Agency for Toxic Substances and Disease Registry (2009).

[5]. Henke, Kevin R. (28 April 2009). Arsenic: Environmental Chemistry, Health Threats and Waste Treatment. p. 317. ISBN 978-0-470-02758-5.

[6]. Georges, Audi; Bersillon, O.; Blachot, J.; Wapstra, A. H. (2003). “The NUBASE Evaluation of Nuclear and Decay Properties”. Nuclear Physics A. Atomic Mass Data Center. 729: 3–128. Bibcode:2003NuPhA.729….3A. doi:10.1016/j.nuclphysa.2003.11.001.

Cobalt

Cobalt is transition metal and was discovered in pure form in 1735. It is essential part of vitamin B12. Its compounds are known as cobalt blue used to color pottery and glass.

History and Discovery

Cobalt has been used as a coloring agent since Bronze Age. It has been found in Egyptian statuettes and in Persian jewelry of the 3rd millennium BCE. It was also used in China for pottery glazes as early as 618-907 CE. Cobalt in the form of metal was isolated by George Brandt in 1735. In 16th century kobold name was given to the ores that contain copper but in the end found that ores contain poisonous arsenic. Brandt in 1742 determined that in various ores, the blue color is due to the presence of cobalt. The word kobold derived from German word meaning ‘’goblin ore’’ [1].

Cobalt

Periodic Table ClassificationGroup 9
Period 4
State at 20CSolid
ColorBluish gray metal
Electron Configuration[Ar] 3d7 4s2
Electron Number27
Proton Number27
Electron Shell2, 8, 15, 2
Density8.9 g.cm-3 at 20°C
Atomic number27
Atomic Mass58.93 g.mol -1
Electronegativity according to Pauling1.88

Occurrence

Cobalt is widely present on the earth crust, in combination with other minerals and in natural water. It is the 32nd most abundant element present on the earth crust. It is present in the sun, in soil, and in bodies of plants and animals. Cobalt is mostly present in combination with nickel. In compound form, it occurs in minerals of in copper and nickel. Most common minerals of cobalt are cobaltite (cobalt sulfoarsenide mineral), linnaeite (sulfide mineral), skutterudite (series of cobalt and nickel mineral), and smaltite (cobalt, nickel arsenide). Cobalt is produced as a byproduct during the mining and refining nickel, silver, lead, copper and iron. Large deposits of cobalt are present in Canada, Australia, Zambia and Brazil.

Physical Characteristics

Refined or polished cobalt is a silvery white metal and have a faint bluish tinge.  Its physical properties resemble iron and nickel. Cobalt is a ferromagnetic (strongest magnet), up to 11210C. Its specific gravity is 8.9. Pure cobalt is obtained through smelting process, which is hard and lustrous and release vapors of arsenic. Cobalt is a is transition metal. It melts at 1495oC and its boiling temperature is 2927oC [2].

Chemical Characteristics

Cobalt is not very reactive metal. It is stable in air and is unaffected by water. Cobalt reacts slowly with dilute acids.  It combines with oxygen and produce Co3O4 and lose oxygen at higher temperature (900oC) and give monoxide (CoO). Cobalt does not catch fire in air and is resistant to burning, unless it is in powder form. Cobalt is weak reducing metal and protected from oxidation by formation of a passivating oxide film. It reacts with halogens and produce halides. Cobalt does not react with hydrogen and nitrogen gas even at high temperatures. At standard temperature it reacts with mineral acids. Cobalt exist in +2 and +3 oxidation state, although -3 to +5 oxidation states compounds are also formed. Cobalt oxides are antiferromagnetic at low temperature. Cobalt (II) oxide has a rock salt structure. In the form of halides, Cobalt (III) fluoride is used in fluorination reaction.

Significance and Uses

  • Cobalt is primarily used in the manufacturing of magnets.
  • Compounds of cobalt gives deep blue color to glass ceramics and inks.
  • It is also used to prepare wear resistant and high strength alloys.
  • 60Co isotope used as radioactive tracer for the production of gamma rays.
  • Cobalt is main part of Vitamin B12 called Cobalamin and in inorganic form it is a micronutrient for bacteria and algae.
  • Cobalt alloys are used in jet turbines and in gas turbines generator.
  • Salts of cobalt are used to give blue colors to porcelain and pottery.
  • It is used as a catalyst in petroleum and chemical industries.
  • Cobalt-60 is also used to irradiate the food in order to extent its shelf life.
  • Cobalt is used in making powerful magnets and magnetic recording media.

Health effects

Cobalt is necessary for human as it is a part of essential vitamin, B12. It stimulates the production of red blood cells and is widely used to treat anemia. In work place, when workers are exposed to high level of cobalt, they can suffer from lungs infections such as asthma and pneumonia. When plants grow on the soil near cobalt mining area, they take up high amount of cobalt which may eventually affect human health. The daily intake dose of cobalt is around 1mg.

Isotopes of Cobalt

Cobalt has one stable isotope 59Co. Cobalt-60 is a commercially important radioisotope that have a half- life of more than 5.2 years [3]. The isotopes of cobalt range in atomic weight from 50 to 73.

REFERENCES

[1]. Enghag, Per (2004). “Cobalt”. Encyclopedia of the elements: technical data, history, processing, applications. p. 667. I

[2]. Properties and Facts for Cobalt”. American Elements. Retrieved 2008-09-19.

[3]. http://www.chemistryexplained.com/elements/A-C/Cobalt.html

Vanadium

Vanadium is a hard metal and was discovered in 1801. It resists corrosion and is widely used to make various alloys.

History and Discovery

Vanadium was discovered by Andres Manuel del Rio in 1801 in Mexico City. He discovered it in a specimen of vanadite. He sent his samples to the Institute den France for inspection and confirmation, but his letter was lost in a shipwreck. Later, Vanadium was rediscovered by Nils Gabriel Sefstrom Swedish chemist who analyzed the samples of iron from a mine in 1830 [1]. Vanadium was isolated by Sir Henry Enfield Roscoe during reaction of vanadium trichloride with hydrogen gas in 1867. Pure vanadium was produced by reducing vanadium pentoxide with calcium in 1927. Vanadium got his named after ‘Vanadis’ who is the Scandinavian Goddess of beauty, as vanadium form various multicolored beautiful compounds.

Vanadium

Periodic Table ClassificationGroup 5
Period 4
State at 20CSolid
ColorBlue-silver-grey metal
Electron Configuration[Ar] 3d3 4s2
Electron Number23
Proton Number23
Electron Shell2, 8, 11, 2
Density6.11 g.cm-3 at 20°C
Atomic number23
Atomic Mass50.94 g.mol -1
Electronegativity according to Pauling1.63

Occurrence

Vanadium in native form is rare in nature. It is the 20th most abundant element in the earth crust. It is present in crude oil, coal, oil shale (sedimentary rock) and tar sands deposits. Vanadium is detected spectroscopically in sun light and in the light from other stars. Vanadium in the form of their compounds exit in nature. It is mined mostly in South Africa, North West China and Eastern Russia. Vanadium in the form of vanadyl ion (functional group in which triple bond exit between V4+ and O2-) is present in sea water. Some mineral water also contain vanadium ions in high concentration [2].

Physical Characteristics

Vanadium is silver white, medium hard and ductile metal. It has strong structural strength. Vanadium is resistant to corrosion and is stable in alkalis and acids including, sulfuric acid and hydrochloric acid. It mostly exit in combined form with certain other minerals. The atomic number of vanadium is 23, and atomic mass is 50.914 g/mol. Vanadium melts at 1910oC and boils at 3407oC.

Chemical Characteristics

Vanadium is quite reactive with oxygen, nitrogen and carbon at high temperatures. Vanadium in aqueous solution forms different color complexes like lilac [V(H2O)6]2+, green  [V(H2O)6]3+ , blue  [VO(H2O)5]2+and yellowVO3.  Vanadium exist in four adjacent oxidation states, +2 to +5. V(II) compounds are strong reducing agent, V(V) are oxidizing agent and V(IV) compounds mostly exist as vanadyl derivatives that contain a VO2+ center. Vanadium pentoxide is used as a catalyst to produce sulfuric acid. Vanadium also form binary halides (VI4, VCL5, VI5) but they are highly unstable. Vanadium also carry out oxidation of organic substances, as oxidation ethanol to form acetaldehyde. Vanadium is miscible in concentrated sulfuric acid, nitric acid, hydrofluoric acid and aqua regia (mixture of concentrated sulfuric and nitric acid) [3].

Significance and Uses

  • Vanadium is used to form alloys of steel and iron to impart desirable corrosion resistant characteristics to them.
  • In chemical industry, vanadium metal sheets, wires and tubes are widely used.
  • Vanadium steel is strong and provides shock resistance.
  • Vanadium is extensively used as a catalyst, such as vanadium pentoxide.
  • Vanadium pentoxide is also used in ceramics.
  • The vanadium redox batteries are used to efficiently store energy.
  • It is also used to protect steel against rust and corrosion.
  • Vanadium foil is used in cladding of titanium to steel as it has compatibility with both iron and titanium.
  • 51V isotopes is used for NMR spectroscopy.
  • Vanadium in the form of supplement is used in pharmaceutical applications.
  • Vanadium is used as a catalyst in the manufacture process of polyamides, like nylon.
  • It is used for the treatment of prediabetes and diabetes.
  • Vanadium in used for manufacturing of steel alloys for cutting and grinding purposes.

Health effects

Vanadium act like insulin and functions to increase the effects of insulin. It is used to treat low blood sugar, high cholesterol, heart disease and water retention in the body. Vanadium is non-toxic, but when ingested in large amounts, it increases the risk of kidney damage. Vanadium uptake by humans is mainly through food including, like buckwheat, soya beans, olive oil, apples and eggs. Acute effects of vanadium include irritability of throat, eyes and lungs. Prolonged exposure can cause damage to the nervous system, general weakness and skin rashes.

Isotopes of Vanadium

Vanadium has only one stable isotope, 51V and one radioactive isotope, 50V, that has a half-life of 1.5×1017 years.  Twenty-four artificial radioisotopes also have been discovered having atomic mass number ranging from 40 to 65. Most stable in these isotopes is 49V, that has half-life of 330 days.

REFERENCES

[1].  https://education.jlab.org/itselemental/ele023.html

[2]. Rehder, Dieter (2008). Bioinorganic Vanadium Chemistry. Inorganic Chemistry (1st ed.). Hamburg, Germany: John Wiley & Sons, Ltd. pp. 5 & 9–10.

[3]. https://www.britannica.com/science/vanadium

Tungsten

Tungsten was discovered in 1783, it is also known as wolfram as it is an exceptionally strong metal. It has the highest melting and boiling points and alloys of tungsten are used in various high-temperature applications.

History and Discovery

Tungsten has been known since prehistoric times, around 350 years ago when Chinese porcelain used peach colored tungsten pigment. It was isolated as a novel element by Juan Jose and Fausto Elhuyar in 1783, through charcoal reduction of the oxide which was derived from wolframite (ore of tungsten). However, few year earlier, Carl Wilhelm Scheele (1781) investigated and isolated the white oxide and referred it as the new acid, tungstic acid also called scheelite. Th e names tungsten and wolfram have been used since its discovery. In British and America Tungsten is more common, while in Germany and other European countries Wolfram is preferred [1]. The word Tungsten has been derived from Swedish language that means “heavy stone” and wolframite is derived from German “wolf rahm” that means “wolf cream”.

Tungsten

Periodic Table ClassificationGroup 4
Period 6
State at 20CSolid
ColorGrayish white, lustrous
Electron Configuration[Xe] 4f14 5d4 6s2
Electron Number74
Proton Number74
Electron Shell2, 8, 18, 32, 12, 2
Density19.35 g.cm-3 at 20°C
Atomic number74
Atomic Mass183.84 g.mol -1
Electronegativity according to Pauling2.36

Occurrence

Tungsten is an abundant element and is present in about 1.5 parts per million in the earth’s crust. Tungsten do not occur as a free metal. It is as abundant as tin and molybdenum and occur in the form of minerals like scheelite (calcium tungstate), wolframite (tin ore in around granite), huebnertie (manganese) and ferberite (iron). The main producers of tungsten in the world are China, Russia and Portugal.

Physical Characteristics

Tungsten is a transition metal.  Raw tungsten steel-grey metal that is hard and brittle. It has very high melting point about, 3422oC and boiling point is about 5930oC. It has ability to retain its strength at very high temperatures. Its density is also very high, about 19.3 times that of water. There are two allotropic forms of tungsten: polycrystalline tungsten is hard, and which makes it difficult to work with but pure single crystalline tungsten is soft and can be cut with hard steel hacksaw. Tungsten atomic number is 74 and atomic mass number is 183.85 [2].

Chemical Characteristics

Tungsten is not very reactive metal. It does not react with oxygen at room temperature. It also resists attacks by acids and alkalis. The most common oxidation state of tungsten is +6 but it exists in all oxidation states from -2 to +6. In powdered form, tungsten reacts with carbon to produce tungsten carbide upon heating. It is mostly resistant to chemical attack but reacts with chlorine to form tungsten hexachloride [3]. Higher oxidation states compounds of tungsten form oxides and they are relevant to their terrestrial occurrence while mid -level oxidation states compounds form metal clusters and low oxidation states are associated with the formation of CO complexes.

Significance and Uses

  • Tungsten is used to make material with high tensile strength, such as tungsten carbide that has a melting point of 2770o
  • Tungsten is used to make wear resistant abrasives cutting tools, such as knives.
  • Tungsten is used to make heavy metal alloys.
  • Tungsten alloys are widely used in aerospace and automotive industries.
  • Tungsten with steel is used in making hard permanent magnets.
  • Tungsten is widely used in mining, construction, electrical and metal working machinery.
  • Tungsten is used as the filament in incandescent light bulbs.
  • Tungsten (IV) sulfide is a high temperature lubricant and component of catalyst for hydrodesulfurization (remove sulfur from natural gas).
  • Oxides of tungsten are used in manufacturing of ceramic glazes.
  • Tungsten salts are widely used in chemical and tanning industries.
  • Tungsten is also used in printing nozzle for 3D printing.
  • Tungsten is similar with gold in respect of density and is used in jewelry as an alternative of gold.
  • Elemental tungsten due to its high melting point make it useful for cathode ray tube and vacuum tube.

Health Hazards

In biological systems, the action of tungsten is antagonistic to molybdenum, which leads to the inhibition of action of various oxidative enzymes, for instance xanthine oxidase. The effect of tungsten on environment is quite limited. In the workplace, people can be exposed to tungsten via inhalation, skin contact and eye contact but has no harmful health effects.

Isotopes of Tungsten

Tungsten has five stable isotopes: Tungsten-180, tungsten-182, tungsten-183, tungsten-184 and tungsten-186. Thirty artificial radioactive isotopes also have been characterized, in which the most stable is 181W, with the half-life of 121.2 days [4].

REFERENCES

[1]. https://www.britannica.com/science/tungsten-chemical-element

[2]. https://www.thebalance.com/metal-profile-tungsten-2340159

[3]. Daintith, John (2005). Facts on File Dictionary of Chemistry(4th ed.). New York: Checkmark Books. ISBN 0-8160-5649-8.

[4]. Sonzogni, Alejandro. “Interactive Chart of Nuclides”. National Nuclear Data Center: Brookhaven National Laboratory. Archived from the original on 2008-05-22.

Titanium

Titanium is a chemical element with symbol Ti and atomic number 22. It is a lustrous transition metal with a silver colour, low density, and high strength.

History and Discovery

Titanium was discovered by William Gregor in 1791 (Great Britain) [1]. Pure titanium was obtained after a long struggle of 119 years, by Matthew Hunter in 1910.Titanium was used for the first time in making military applications by Soviet Union in 1950s. And later, US defense department recognized the significance of titanium and set motion to commercialization of titanium-based war strategic material.  The element was name Titanium after the Greek Gods, Titans by Martin Klaproth.

Titanium

Periodic Table ClassificationGroup 4
Period 4
State at 20CSolid
ColorSilvery grey-white metallic
Electron Configuration[Ar] 3d2 4s2
Electron Number22
Proton Number22
Electron Shell2, 8, 10, 2
Density4.54 g.cm-3 at 20°C
Atomic number22
Atomic Mass47.87 g.mol -1
Electronegativity according to Pauling1.54

Occurrence

Titanium is an abundant element and is characterized as the 9th most abundant element in the Earth’s crust [2]. Native or elemental titanium is very rare. It exists in form of compounds in minerals. Mostly, it is present as oxides in igneous rocks and sediments. The common minerals of titanium include brookite, sphene (calcium titanium silicate), perovskite, titanite, anatase and rutile (titanium oxide). The minerals of economic importance are ilmenite (iron titanium oxide) and rutile, and around 0.7 million tons were mined in 2011Titanium is also present in ocean water with a concentration of around 4 picomoles [3]. Titanium is present in extraterrestrial environment, including the sun, moon, meteorites and certain type of stars. The main producers of titanium include India, US, New Zealand, China, Canada, South Africa and Norway.

Physical Characteristics

Titanium is a silvery white lustrous metal. it is quite light-weight as compared to other metals. it has an outstanding strength and durability. Titanium is highly resistant to corrosion. The rate of corrosion of titanium is so slow that it will be barely rusted after 4000 years of sea water exposure. When heated, titanium is ductile and malleable. Titanium is insoluble in water. however, it is readily dissolved in concentrated sulfuric and nitric acids.

Chemical Characteristics

Titanium exists in two oxidations states, +4 and +3. The most common oxidation state of titanium in compounds is +4 [4]. the titanium (IV) compounds are usually termed as titanates.  Titanium frequently forms covalent bonds with its compounds. Titanium form titanium oxides with oxygen, and there are three forms of titanium oxide, rutile, anatase and brookite. Titanium forms compound with sulfur.

Significance and Uses

  • Titanium is widely used in making esteem quality steel alloys. It is also used in making of other alloys such as, copper, vanadium and aluminum.
  • Titanium products are widely used in various aerospace, and industrial markets.
  • Titanium in the form of titanium dioxide is used in the making of white colored plastic products, toothpaste, paints, and paper.
  • Titanium is widely used as a strengthening material in the making of fishing rods, helmets, hockey, tennis racket, drill bits etc.
  • Titanium is widely used in automotive applications, such as bikes, sports car and automobile, as it is light weight and has high strength.
  • Titanium is used in making structure of ships, shafts etc. as it is extraordinarily resistant to corrosion by sea water.
  • Titanium is used in making of prosthetic joints for joint replacement therapy , such as hip joint and also as dental implants.
  • Titanium is used in making various electronics including high end cameras, mobiles, and laptops.
  • Nanoparticles of titanium are used as delivery tool for various drugs.

Health Hazards

Titanium is a low toxic metal. However, certain compounds including titanium chlorides are quite corrosive.  In powder from, titanium is readily combustible.

Isotopes of Titanium

There are eighteen isotopes of titanium. There are five stable isotopes present in naturally occurring titanium: titanium-46, titanium-47, titanium-48, titanium-49, and titanium-50.  The most abundant natural isotope is titanium-48 (73.8%). There are also eleven radioactive isotopes of titanium and among them titanium-44 is the most stable. The radioactive isotopes of titanium emit gamma rays and positrons.

REFERENCES

[1]. William Gregor, Beobachtungen und Versuche über den Menakanite, einen in Cornwall gefundenen magnetischen Sand., in Lorenz Crell’s Chemische Annalen, 1791, p40

[2]. Barksdale 1968, p. 732

[3]. “Titanium”. Encyclopædia Britannica. 2006. Retrieved 29 December 2006.

[4]. Greenwood 1997, p. 970

Strontium

Strontium is an alkaline earth metal discovered in 1808 by Humphry Davy. It is a reactive element and has various useful radioactive isotopes.

History and Discovery

Strontium was discovered earlier from mineral of barium (barium carbonate) by Adair Crawford in 1798. Strontium was formally discovered by Humphry Davy in 1808. This element was named after the place where its mineral was discovered, Strontian, a village in Scotland [1]. Strontium was used to produce sugar in the 19th century [2]. Large scale production of strontium was started due to its high demand and use in making television tube (cathode ray tubes) and its demand dramatically decreased with advancements in methods to make display screens and tubes.

Strontium

Periodic Table ClassificationGroup 2
Period 5
State at 20CSolid
ColorSilvery white metallic; with a pale yellow tint
Electron Configuration[Kr] 5s2
Electron Number38
Proton Number38
Electron Shell2, 8, 18, 8, 2
Density2.54 g.cm-3 at 20°C
Atomic number38
Atomic Mass87.62 g.mol -1
Electronegativity according to Pauling0.95

Occurrence

Strontium is quite abundant element and is ranked as the 15th most abundant element in the earth’s crust (about 0.034 %) [3]. Strontium is not present in native elemental form. It exists only in the form of compounds with other metals. the most common minerals of strontium are strontianite (strontium carbonate), celestite (strontium sulfate). Strontium is also present in ocean water with a ratio of 8mg/L. The largest deposits of strontium are found in China, Spain and Mexico.

Physical Characteristics

Strontium is a silvery white metal with a yellowish tint. It is characterized as an alkaline earth metal. It is soft metal and has a density of around 2.64 g/cm3. Strontium burns with a distinctive red color flame. There are three allotropic forms of strontium. Strontium has a considerably high refractive index. Strontium readily dissolves in liquid ammonia and forms a dark blue solution [4].

Chemical Characteristics

Strontium is highly reactive. It reacts with air to forms yellowish strontium oxide. Strontium readily reacts with water to form strontium hydroxide. Strontium reacts with nitrogen at high temperature, above 380°C. In powdered form, strontium undergoes spontaneous ignition (is pyrophoric) when exposed to air. Strontium is stored in kerosene or mineral oil to increase its shelf life and avoid reaction with water or air.

Significance and Uses

  • Strontium is widely used in the manufacturing of aerosol paints.
  • Strontium is used to make fireworks (pyrotechnics) for industrial and military purposes, including oxygen candles, safety matches and explosive bolts.
  • Chloride of strontium are used in manufacturing of medicinal toothpastes for sensitive teeth.
  • Strontium is used to make high quality pottery glaze.
  • Radioactive isotope of strontium (strontium-90) is used in nuclear power plants, spacecrafts and satellites as power generators (radioisotope thermoelectric generator RTGs).
  • Strontium-89 (radioactive isotope) is used in prostate cancer therapy to treat bone pain as the primary ingredient in medicine called Metastron.
  • Strontium is also used in the treatment of osteoporosis (strontium ranelate).

Health Hazards

Strontium is highly similar to calcium, and that is why it is absorbed in bone in place of calcium. This retention in body is mostly harmless as stable forms of strontium do not cause any health hazard. However, radioactive isotope, strontium-90 can lead to bone cancer. The Chernobyl nuclear incident (1986) is the most tragic example of radiation contamination as it polluted around 30,000 km2 of area with greater than 10 kBq/m2 of strontium-90 [5].

Isotopes of Strontium

There are sixteen naturally occurring isotopes of strontium. Naturally occurring strontium has four stable isotopes: strontium-84, strontium-86, strontium-87 and strontium -88. The most abundant stable isotope is strontium-88 (82.58 %). Strontium-87 is thought to be produced during the Big Bang and is also obtained during the radioactive decay of rubidium. The radioactive isotope, strontium-90 is produced as a byproduct of nuclear fission reaction.

REFERENCES

[1]. Murray, W. H. (1977). The Companion Guide to the West Highlands of Scotland. London: Collins. ISBN 0-00-211135-7.

[2]. Heriot, T. H. P (2008). “strontium saccharate process”. Manufacture of Sugar from the Cane and Beet. ISBN 978-1-4437-2504-0

[3]. Turekian, K. K.; Wedepohl, K. H. (1961). “Distribution of the elements in some major units of the Earth’s crust”. Geological Society of America Bulletin. 72 (2): 175–92. Bibcode:1961GSAB…72..175T. doi:10.1130/0016-7606(1961)72[175:DOTEIS]2.0.CO;2.

[4]. Greenwood and Earnshaw, pp. 112–13

[5]. “Chernobyl: Assessment of Radiological and Health Impact, 2002 update; Chapter I – The site and accident sequence” (PDF). OECD-NEA. 2002. Retrieved 3 June 2015