Sulfur

Sulfur has a history that dates to ancient times. It was discovered around 2000 BCE by Chinese. Sulfur has diverse biological significance and have numerous industrial implications.

History and Discovery

Sulfur has been known since prehistoric times and has been used in ancient Greece, Egypt, China and India. In very early times, sulfur was named Torah, and was also mentioned in Bible, by the termed “brimstone”, which means “burning sulfur”. Sulfur was known for its bactericidal activity in Egypt and Greece and was used for fumigation and in medicines and ointments [1]. In China, as early as the 6th century BC, sulfur was known as shiliuhuang and was extracted from pyrite. It was used mainly in black gunpowder by the Chinese. Antoine Lavoisier, in 1777 propose that sulfurs is a distinct element and elemental sulfur was discovered in 1867.

Sulfur

Periodic Table ClassificationGroup 16
Period 3
State at 20CSolid
ColorYellow
Electron Configuration[Ne] 3s2 3p4
Electron Number16
Proton Number16
Electron Shell2, 8, 6
Density2,07 g.cm-3 at 20°C
Atomic number16
Atomic Mass32,06 g.mol -1
Electronegativity according to Pauling2.58

Occurrence

Sulfur is quite abundant on Earth as well as in the universe.  Its ranked 10th in order of abundance among all elements in the universe. Sulfur is created in huge stars and is present in various kinds of meteorites. It is produced during fusion reaction between nucleus of helium and silicon. In the Earth’s crust, sulfur is the 5th most abundant element by mass. It is ubiquitous in volcanic regions and in hot soring areas of the world. The Pacific Ring of Fire is especially known for its abundance in sulfur reserves. Sulfur is also found in native form on earth and is formed because of metabolic activity of anerobic bacteria that degrade sulfate minerals. The most common mineral of sulfur includes, gypsum, pyrite, barite, cinnabar and galena. Sulfur is also released into environment, especially in tropical areas, by the weathering of mineral ores. Currently, sulfur is produced from natural gas, petroleum and fossil reserves. The largest producers of sulfur include China, Canada Japan, Chile and Indonesia. Sulfur is a vital component of all living cells and is embedded in the proteins, DNA, and large variety of enzymes of plants, animals and microbes. Human body is comprised of various forms and compounds of sulfur and is considered as the eight most abundant element by weight in the human body.

Physical Characteristics

Sulfur is a yellow color crystalline non-metal that is solid at room temperature. Sulfur exists in various allotropic forms and have around 30 solid allotropes. It has the highest number of allotropes among all elements. Octasulfur, cycle-S8 is the most common allotrope of sulfur [2]. Sulfur is insoluble in water.

Chemical Characteristics

Sulfur is a reactive metal. It forms compounds with all other elements, except nitrogen, gold, iodine, platinum and the Nobel gases. Upon combustion, sulfur gives out a blue flame and produces sulfur oxide that has pungent odor. Sulfur have various oxidation states, +2, +4 ,+6 and +6 and +4 are more common [3].

Significance and Uses

  • Sulfur is widely used to make fertilizers, such as calcium sulfate.
  • Sulfur is used in various agrochemicals, such as fungicides and insecticides. Dusting of elemental sulfurs in powdered form has been used widely to eliminate the growth of fungus from grapes, and many vegetables. It is also used as insecticide to eliminate ticks and mites from crops and plants.
  • Various compounds of sulfur, especially organo-sulfurs are wieldy used in pharmaceutical industry. A large group of drugs, termed as sulfa drugs are broad spectrum antibacterial sulfonamides. Similarly, penicillin and cephalosporin contain sulfur.
  • Sulfurs is widely used for fumigation purposes.

Health Hazards

Sulfur is non-toxic. However, burning of sulfur can lead to production of sulfur dioxide gas, which at high concentration can lead to damaging effects on eyes, lungs and other tissues. Similarly, other compounds of sulfur, including sulfuric acid is highly corrosive acid and its fumes are damaging to the eyes and nasal linings. Hydrogen sulfide is a highly toxic compounds and resembles cyanide in its toxicity.

Isotopes of Sulfur

There are 25 natural isotopes of sulfur and only four are stable, including sulfur-32, sulfur-33, sulfur-34, and sulfur-36 [4]. There are also various radioactive isotopes of sulfur, among which only sulfur-35 have a relatively ling half-life (85 days) and all others are significantly unstable.

REFERENCES

[1]. Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.

[2]. Steudel, Ralf; Eckert, Bodo (2003). Solid Sulfur Allotropes Sulfur Allotropes. Topics in Current Chemistry. 230. pp. 1–80. doi:10.1007/b12110. ISBN 978-3-540-40191-9.

[3]. Egon Wiberg; Nils Wiberg (2001). Inorganic Chemistry. Academic Press. pp. 513–. ISBN 978-0-12-352651-9.

[4]. https://www.sciencedirect.com/topics/earth-and-planetary-sciences/sulfur-isotopes

 

 

 

 

 

 

SILVER

Silver is one of the seven metals of antiquity and have been considered precious throughout history. Silver coins, ornaments, utensils, jewellery and conductors have been popular since its discovery.

 

History and Discovery

Silver is a prehistoric metal and human have used silver as a primitive form of money along with gold and copper. Its exact discovery has not been known. The name silver is derived from Latin word argentum, that means “white” or “shiny”. Ornaments of silver have been found from ancient civilizations. In Egypt, silver was considered more precious than gold in the 15th century BC [1]. Silver was predominantly present in form of ores, and various techniques were developed as early as in the 4th millennium BC, to extract pure silver. Greeks and Romans also used silver as currency. The rise of Athens was attributed to the nearby mines, which supplied about 30 tons of silver in 300 years (from 600 to 300 BC). Later, Romans became the biggest dealers of silvers and their economy circulated about ten times more silver than the combined amount of silver present in Europe in 800 AD. This trend continued until met a severe downfall with the fall of the Roman Empire. In Middle ages, Central Europe became the heart of silver mining and production. Around the 18th century, America was discovered and soon became the dominant producer of silver. But Spain and China remained the main collector of silver from throughout the world and the mines of the “New World” (America) supported the empire of Spain. But in the 19th century, North America, including Mexico, Canada and Nevada (US) took control over the production and are still the leading producers of silver, along with Peru.

Occurrence

Silver is quite abundant on Earth and is present in about 0.08 ppm in the Earth’s crust. It is predominantly present in the form of minerals and sulfide ores, such as argentite and acanthite. Silver is also present in salty water in the form of chlorargyrite (AgCl), which have high deposits in Wales and Chile. The principal source of silver are ores of lead, copper and lead-zinc. They are prevalent in China, Mexico, Australia and Peru [2]. Currently, silver is produced as a secondary product in the refining of lead, copper and zinc. Silver mines are present in Australia, Poland, Bolivia, Mexico, china, Peru. Recently, one of the largest silver deposits in the world have been discovered in Tajikistan (Central Asia) [3].

Physical Characteristics

Silver is a metallic white lustrous transition metal. It is very soft, malleable and ductile. Silver has the highest electrical conductivity among all metals. It also has exceptionally high thermal conductivity. And has the lowest contact resistant among all metals.

Chemical Characteristics

Silver is a fairly unreactive metal. it has the lowest first ionization energy in among Group 11 elements. It predominantly exists in oxidation state of +1, however, +2 and +3 states also exist. Compounds of silver have strong covalent characters. The electron affinity of silver is 125.6 kL/mol which is higher than hydrogen but lower than oxygen. Silver readily combines with zinc, gold and copper and form alloys. Silver forms organometallic compounds but they are quite unstable. Silver is resistant to reaction with air, even at higher temperatures. This property makes it a noel metal, like gold. It can react with sulfur compounds and forms a black tarnished silver sulfide. It readily dissolves in hot concentrated sulfuric acid, nitric acid and in aqueous cyanide solution.

Significance and Uses

  • Silver is used in solar panels, water storage tanks, tableware, and utensils.
  • Silver is widely used in electronic devices as conductors due to its high electrical conductivity.
  • Silver is used as a colorant for making stained glass.
  • Silver is used for brazing of various metals.
  • Various compounds of silver are used in making of photographic and X-ray film.
  • It is used in the manufacturing of window coatings and specialized mirrors.
  • Silver is used as catalysis in various chemical reactions, such as oxidation relations.
  • Silver is used in the manufacturing of wide range of chemical equipment as it has low chemical reactivity and is resistant to corrosion.
  • Silver nitrate in dilute solution form, is used a common disinfectant and is present as antibiotic coating in catheters, bandages and other medical utensils.
  • Nanoparticles of silver are widely used in for medicinal purposes as antifungal and antibacterial agents.

Heath Hazards

Silver compounds have low toxicity. Silver nitrate have caustic effect and can lead to tissue damage and diarrhea, low blood pressure and paralysis if ingested. Exposure and ingestion of large concentrations lead to accumulation of silver in the body, that lead to bluishness of skin and eyes. Some compounds of silver are highly combustible and require special precautions while handling, these include silver amide, silver oxide etc.

Isotopes of Silver

There are two stable isotopes in naturally occurring silver, silver-107 and solver -109. Both isotopes have almost the same abundance, i.e. silver-107 is 51 %, which is quite rare. There are 28 radioactive isotopes of silver [4].

REFERENCES

[1]. Weeks, pp. 14–19

[2]. Hammond, C. R. (2004). The Elements, in Handbook of Chemistry and Physics (81st ed.). CRC press. ISBN 978-0-8493-0485-9.

[3]. “Why Are Kyrgyzstan and Tajikistan So Split on Foreign Mining?”. EurasiaNet.org. 7 August 2013. Retrieved 19 August 2013.

[4]. “Atomic Weights and Isotopic Compositions for All Elements (NIST)”. Retrieved 11 November 2009.

 

IODINE

Iodine belongs to halogen group in the periodic table. Iodine is an essential mineral commonly found in sea food. Iodine plays an important role in proper functioning of thyroid hormone in the human body.

 

Discovery and History

During Napoleonic Wars, Bernard Courtois, a French chemist discovered Iodine in 1811, while he was extracting sodium and potassium compounds from seaweed ash. He noticed a cloud of violet gas when sulfuric acid was added to the ash and proposed that a distinct element is present in the ashes. Two years later Joseph Louis Gay-Lussac gave the name “iode” to this element, that is the Greek word for ‘’violet colored’’. Iodized table salt was sold first time in Michigan in 1924 [1].

Occurrence

Iodine is the 61st most abundant element on Earth. It is a rare element and does not exist in free -state. About 50mg per metric tons of iodine is present in the sea water [2]. Iodine is the least abundant halogens and is present in only 0.46 ppm (parts per million) of earth crust rocks. The mineral Caliche (sedimentary rock), is the main source of iodine.  Brine (high concentration of iodine in water) is also used for the extraction of iodine and Japan and USA have major deposits of brine. Iodine is also formed in oysters, seaweeds and cod liver. Human body contains iodine in the form of thyroxin produced by the thyroid glands. Iodine is chiefly obtained from sodium iodate (NaIO3) and sodium periodate (NaIO4). Chile and Japan are prime producer of iodine in the world.

Physical characteristics

Iodine is bluish black, shiny solid and has a pungent odor. It is non -metal but show some metallic properties. Iodine is hardly dissolved in water but gives a yellow color solution. Iodine belongs to group halogens, that are the salt formers. Iodine when dissolved in chloroform gives purple color solution. Melting point and boiling point of iodine are highest as compared to other halogens. Its melting point is 113.70C and boiling point is about 184.30C. Atomic number of iodine is 63 and easily attached with organic compounds.  Iodine’s atomic weight is 126.90. The density of iodine is very low, about 4.94g/ml.

Chemical characteristics

Iodine is reactive element among the halogens. It has lowest ionization energy and is easily oxidized. Iodine has various oxidation states, including +1 (iodides), +3, +5 (iodates) and +7 (periodate). and is more stable than bromine and chlorine. Iodine molecules act as Lewis acid with combined with many Lewis bases. Electron affinity of iodine is also similar to other halogen atoms. The simplest compound of iodine is hydrogen iodide. It is a colorless gas that reacts with oxygen to give water and iodine. Iodine does not react with oxygen or nitrogen.  Iodine reacts with nonmetals and forms iodides, for example, silver and aluminum are converted into iodides. The iodide ion is a strong reducing agent and gives up electron easily. Iodide solutions are colorless but gives brownish tint due to oxidation. Iodine reacts with zinc and forms zinc iodide. This reaction is very exothermic and produce violet color vapors of iodine.

Significance and Uses

  • Iodine in form of potassium iodide and alcohol is widely used as a disinfectant.
  • Radioactive isotopes of iodine 131I is used to treat thyroid cancer.
  • Iodine is used as catalyst in the preparation of acetic acid.
  • Iodine is widely used in laboratories to test the presence of starch in solution.
  • Iodine is used in iodoform test to detect the presence of methyl ketones.
  • Iodine has antiviral and antimicrobial action.
  • Erythrosine, an organo-iodine compound is important food coloring agent.
  • Potassium iodide is used to make photographic films.
  • Iodine is added in table salt for nutrition to prevent goiter in the thyroid gland.
  • Tungsten iodide is used to stabilize filaments in light bulbs.
  • Iodine is used in making of dyes and various pharmaceuticals products.
  • Iodine- 129 is used in rainwater studies.
  • Potassium iodide can be used for the treatment of people in a nuclear disaster area.

Health effects

Iodine is an essential part of human diet because the body does not make iodine and it is needed to produce thyroid hormones that regulate the growth and metabolism of the body. Deficiency of iodine may cause enlargement of thyroid gland (goiter) hypothyroidism and intellectual disabilities in infants and children. Goiter disease is also present in animals like dog, cattle, goat and birds. People who consume iodine based food on daily bases can also face various health problems, such as disturbed heart beats and loss of weight.  Pure iodine is dangerous and poisonous if ingested. In U.S, the recommended daily dose of iodine is 110 to 113µg for infants to 12 months, 90 µg for 1 to 18 years old, 220µg for pregnant women and 290µg for lactation. U.S Food and Drug recommends an intake of 150µg per day of iodine for both men and women.

Isotopes of Iodine

Iodine has 34 isotopes with mass numbers ranging from 108 to 141. 127I is the only stable isotope of iodine [3]. The longest-lived isotope is 129I has half-life of 15.7 million years [4]. 125I has half-life of 59 days.

 

References

[1]. https://www.livescience.com/37441-iodine.html

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

[3]. https://www.chemicool.com/elements/iodine.html

[4]. Audi, G.; Bersillon, O.; Blachot, J.; Wapstra, A. H. (2003). “The NUBASE evaluation of nuclear and decay properties” (PDF). Nuclear Physics A729: 3–128.

KRYPTON

Krypton is a rare and inert gas and belong to the Nobel gases. It was discovered by Sir William Ramsay in 1898. It emits unique and sharp spectral lines and is widely used in high speed photography and lasers.

 

History and Discovery

Krypton was discovered by Sir William Ramsay and Morris Travers in 1898, as a residual gas in a chamber after all components of liquid air have been evaporated. Ramsay received the Nobel prize in Chemistry in 1904 for his contributions in the discovery of krypton and other noble gases, except radon. Before its discovery, Ramsay believed that in nature, often element hides in another. And the gap between argon and helium in the periodic table made him look closely within elements. The name krypton has been derived from the Greek word, kryptos, that means “hidden” [1].

 

Occurrence

Krypton is a rare gas. It is present in about 1ppm in the Earth’s atmosphere. Krypton is obtained by fractional distillation of liquid air. Krypton is found in abundance in space. Krypton is also produced during the uranium fission reaction [2].

 

Physical Characteristics

Krypton is colorless and odorless gas. Krypton in solid state is white and have a cubical structure. Krypton gives two distinct lines of green and yellow color in its emission spectrum [3]. Krypton has a density of 3.749 g/L at standard conditions and is almost three times denser than air. Krypton is affected by an external magnetic field and is diamagnetic. Krypton is highly volatile and quickly vaporizes when exposed to water.

 

Chemical Characteristics

Krypton is chemically inactive. However, under extreme conditions, it reacts with fluorine and forms krypton difluoride (KrF2). It is a crystalline white solid and is stable at low temperature. Crystals of krypton hydride can be formed under increased pressure, i.e. above 5GPa. Krypton has two oxidation states [4].

 

Significance and Uses

  • Krypton is widely used in photography. It is used flash lamps for high speed photography.
  • It is used for making energy-efficient fluorescent lamps.
  • Krypton is also used in high-powered flash lamps on airport runways.
  • Krypton is used in making high powered gas lasers and krypton fluoride is used in some lasers.
  • Krypton-85 is used a marker for the detection of nuclear weapon production and research facilities.
  • Krypton is used in making flash signs like luminous neon light signs, that glow with a distinct green yellow light.
  • Krypton emits a shaper red light and is used for making red lasers used in high power laser shows.
  • Krypton fluoride lasers are also used in research involving nuclear fusion reactions for the generation of energy.
  • The international unit of meter in System International use the wavelength of light emitted by krypton isotope, 86Kr, to define distance of one meter.

Health Hazard

Krypton is non-toxic in nature. If inhaled in high concentrations, it can act as an asphyxiant and displace the normal concentration of oxygen in the lungs. This can lead to breathing difficulty and lead to drowsiness and unconsciousness.

Isotopes of Krypton

There are 25 isotopes of krypton, with mass number ranging from 71 to 95. There are six stable isotopes in the naturally occurring krypton. These include krypton-78, krypton-80, krypton-82, krypton-83, krypton-84, krypton-84, krypton-86. Krypton-84 is the most abundant isotope, 57%. There are around thirty artificial isotopes of krypton. Krypton-81 and krypton-85 are radioactive isotopes of krypton but have considerably long half-lives [5].

REFERENCES

[1]. William Ramsay; Morris W. Travers (1898). “On a New Constituent of Atmospheric Air”. Proceedings of the Royal Society of London. 63 (1): 405–408. doi:10.1098/rspl.1898.0051.

[2]. “Krypton” (PDF). Argonne National Laboratory, EVS. 2005. Archived from the original (PDF) on 2009-12-20. Retrieved 2007-03-17.

[3]. “Spectra of Gas Discharges”.

[4]. Kleppe, Annette K.; Amboage, Mónica; Jephcoat, Andrew P. (2014). “New high-pressure van der Waals compound Kr(H2)4 discovered in the krypton-hydrogen binary system”. Scientific Reports. 4. Bibcode:2014NatSR…4E4989K. doi:10.1038/srep04989.

[5]. http://www.rsc.org/periodic-table/element/36/krypton

 

 

 

BISMUTH

Bismuth is a post-transition metal and was discovered in 1753. It is a highly useful metal, and is component of various cosmetics, alloys and medications.

 

Discovery and History

Bismuth has been known to human civilizations since 1400, but due to close resemblance, it was confused with lead and tin. Bismuth is among the earliest ten elements to be discovered and used by humans. It was discovered as a distinct element in 1753 by Claude Geoffroy the Younger. Bismuth form white oxide layer when exposed to air and its name has been derived from the Latin word “bisemutum” that means “white substance” and dates to 1600s [1].

Occurrence

Bismuth is ubiquitously distributed on the Earth, both in pure and mineral form. The primary source of bismuth is its mineral, bismuthinite, or bismuth sulfide (Bi2S3) [2]. Bismuth is also obtained as a by-product in the refining processes of silver, lead, tin and gold ores. The leading producers of bismuth include Japan, Canada, Mexico, Peru and Bolivia [3].

Physical characteristics

Bismuth is a pinkish white metal. It is solid at room temperature and is lustrous and brittle in nature. Bismuth is a highly diamagnetic element. Bismuth develops a layer of oxide on its surface, that shows various colors from blue to yellow and shows iridescent (shine surface that gives different colors when angel of illumination is changes) property. Bismuth has an exceptionally low thermal conductivity. And high resistance to electric field. It has a low melting point, 271.4 C. Liquid bismuth expands like water when exposed to freezing temperature and forms a crystalline structure. Upon combustion, bismuth gives a blue flame. Its boiling point is 2847C. Bismuth is a very dense metal, 9.78 g/cm3. Bismuth is scarcely radioactive and decay to thallium, but it has an extremely long half-life (20 billion billion years). But due to its extremely long half-life, it is considered stable. Bismuth acts as semiconductor when present in extremely thin layers.

Chemical characteristics     

Bismuth is a stable element. It does not react with air at standard temperature, but forms bismuth oxide at high temperature. Bismuth form pentavalent and trivalent compounds, and trivalent being more common. It reacts with various acids, such as nitric acid, hydrochloric acid and sulfuric acid. Among halogens, bismuth reacts with fluorine (at 500C) to form bismuth (V) fluoride.

 Significance and Uses

  • Bismuth is widely used in medication as bismuth subsalicylate (pink bismuth), (Bismol), for curing diarrhea and upset stomach and as bibrocathol to treat eye infections.
  • Bismuth (Bismuth oxychloride) is used in cosmetic industry for the manufacturing of nail colors, eyeshadows and lipsticks. It gives the products the required shine and gloss.
  • Bismuth is used in the manufacturing of rubber and synthetic fiber.
  • Bismuth is used in nuclear reactors.
  • Bismuth is widely used as pigments in paints and is non-toxic alternative of cadmium based paints and colors.
  • Bismuth has replaced lead in various alloys, due to its low toxicity compared to lead.

Health Hazards

Bismuth has a very low toxicity. The low solubility of bismuth salts in water is the primary cause of low toxicity. However, long term exposure to high levels of bismuth can lead to development of bismuth line on the gums. Bioaccumulation of bismuth has also been known and the biological half-life of bismuth has been reported to be less than a week, but it can persist in the kidneys for many years of individuals given medication with bismuth compounds [4].

Isotopes of Bismuth

There is one natural isotope of bismuth, bismuth-209. It is one of the heaviest isotopes in nature. Bismuth is considered non-radioactive due to its extraordinary long alpha decay half-life [5]. Various artificial isotopes of bismuth have been synthesized. Various isotopes of bismuth are also formed during the radioactive disintegration chains of different metals, such as uranium-233, radium and actinium [6].  

REFERENCES

[1]. Harper, Douglas. “bismuth”. Online Etymology Dictionary.

[2]. Hammond, C. R. (2004). The Elements, in Handbook of Chemistry and Physics (81st ed.). Boca Raton (FL, US): CRC press. p. 4–1. ISBN 0-8493-0485-7.

[3]. Anderson, Schuyler C. “2016 USGS Minerals Yearbook: Bismuth” (PDF). United States Geological Survey. Retrieved 1 July 2016.

[4] Fowler, B.A. & Sexton M.J. (2007). “Bismuth”. In Nordberg, Gunnar. Handbook on the toxicology of metals. Academic Press. pp. 433 ff. ISBN 978-0-12-369413-3.

[5]. Carvalho, H. G.; Penna, M. (1972). “Alpha-activity of209Bi”. Lettere al Nuovo Cimento. 3 (18): 720. doi:10.1007/BF02824346.

[6]. Loveland, Walter D.; Morrissey, David J.; Seaborg, Glenn T. (2006). Modern Nuclear Chemistry. p. 78. ISBN 978-0-471-11532-8.

TIN

 

Tin has been known since ancient times, and has characteristics feature of superconductor. Tin belongs to the carbon family and is widely used in container to form protective layer.

Discovery and History

Tin has been used since ancient civilization, mostly in the form of alloy as early as 3000 BC. After 600BC, Pewter purified metallic tin alloy which contained 85-90% tin and consisted of copper, antimony and lead. It was used for making flatware (tableware). Chinese mined tin in the province of Yunnan in about 700 BC. Mediterranean peoples from British Isles (group of islands) were mining tin around 300-200 BC. British scientist Robert Boyle published his experimental description on the oxidation of tin in 1673. The symbol of Tin ‘’Sn’’ was derived from its Latin word Stannum [1]. Tin was used in making toys in early 1800s.

 

Occurrence

Tin is present in igneous rocks of earth’s crust. It is the 49th most abundant element on Earth. Tin is not a native element, and mostly extracted from its ore Cassiterite (SnO2), reduced with coal in a furnace [2]. Tin is found in the ‘’tin belt’’ which stretches from South China, Thailand, Burma to Malaysia and then Indonesia.

Physical characteristics

Tin is soft, pliable and ductile in nature. Tin retains its color due to formation of stannic oxide protective film on the surface via reaction with oxygen of the air. Tin is used as oxidation resistant coating material due to its low melting point. Tin has atomic number 50 and its atomic weight is 118.69. Tin has two allotropes forms: β-tin is silvery white soft metal, and at low temperature it transforms into less dense α- tin metal which is gray in color and has diamond cubic structure [2]. Tin melts at low temperature about 2310 C. Its boiling point is very high, about 22600C. Tin produce a distinct sound, known as the ‘Tin cry’’ when it is bent.  During winter, it changes from one allotropic form to another [3].

 

 

Chemical characteristics

Tin is resistant to corrosion and at room temperature it is unaffected by water and oxygen. But with increase in temperature, tin reacts with oxygen and forms oxides. Tin does not react with dilute acids but is easily dissolved in concentrated acids. Tin reacts with halogens to form compounds like tin chloride and tin bromide. When oxygen is present in a solution, tin act as a catalyst in the chemical reaction. Tin compounds usually occur in the divalent state (Sn2+) and tetravalent state (Sn4+). Tin forms halides, Sn (IV) halides include, SnF4, SnCl4, SnBr4 and SnI4. Tetrafluoride is polymeric (poly means many, mer means parts), and others compounds are volatile. Sn (II) halides are SnF2, SnCl2, SnBr2 and SnI2. These all are polymeric solids. Tin forms many oxides. SnO2 is amphoteric in nature and sulfides of tins exits in both +2 and +4 states.

Uses and Significance

  • Tin is used in alloys with lead as solders to attach metal wires with electrical devices.
  • Tin is also used in the manufacturing of various alloys such as, bronze, pewter, phosphor bronze.
  • Tin oxide is used for making ceramic bodies opaque.
  • Organic tin compounds are used as biocides and fungicides.
  • Tin is helpful to coat other metals to prevent them from corrosion.
  • For the prevention of barnacles in ships and boats, tin compounds are used as anti-fouling (a compound that retards the growth of algae and marine organisms).
  • Tin chloride is used as a powerful reducing agent.
  • Steel container plated with tin is used for preservation of food.
  • Tin powder is also used in the making of paper, inks and spray.
  • Electroplating of tin is very common
  • Tin foil is also used to wrap candies tobacco etc.
  • Tin chromate is used as coloring agent for porcelain.
  • Tin valve and piping is helpful in maintaining purity in water.
  • Tin salt sprays are used to produce electrically conductive coating on glass.
  • Tin is used in making various toothpaste.
  • It is used in various textiles industries.

 

Health effects

Tin is non -toxic element but some compounds of tin are toxic in nature. Mostly, using tin utensil can have adverse effects on health. Tin inhalation can cause problem such as nausea, diarrhea, vomiting and cramps [4].

 

Isotopes of Tin

Tin have ten stable isotopes. The isotopes contain atomic masses of 112, 114 to 120, 122 and 124. The most abundant isotopes are 120Sn, 118 Sn and 116Sn, and least abundant isotope is 115Sn. Tin also has 29 unstable isotopes with atomic masses from 99 to 137. 126 Sn isotopes has half -life of 230,000 years.

 

Reference

  1. https://www.britannica.com/science/tin
  2. https://en.wikipedia.org
  3. Http://www.chemistryexplained.com/elements/T-Z/Tin.html
  4. http://www.chemistryexplained.com/elements/T-Z/Tin.html

 

FLUORINE

 

Fluorine is highly reactive and the most electronegative element in the periodic table. It was discovered by Andre-Marie Ampere in 1810.

 

Discovery and History

The high reactivity and corrosive nature of fluorine led to delay in the discovery and isolation of fluorine as a distinct element. Several early experiments with fluorine caused serious health hazards to the scientist. However, in 1810 André-Marie Ampère, proposed that the unknown component of hydrofluoric acid is a distinct element that is analogous to chlorine. Later, Humphry Davy named the substance fluorine from fluoric acid and the suffix “-ine” used for all halogens [1]. Long before its formal discovery, fluorine minerals were used in various laboratory and industrial settings. It was used for the smelting of ores and was named fluo, which is the Latin word for flow. The name fluorine was derived from its characteristic of flowing minerals. The symbol of fluorine is F, which is also derived from its Latin name which later became, fluorum.

Occurrence

Fluorine is the 13th most abundant element on Earth and is present in about 600 ppm by mass in the Earth’s crust [2]. In universe, fluorine is the 24th most abundant element. Fluorine in elemental form quickly reacts with vapors present in the atmosphere and thus its elemental existence is almost nil. Fluorine does not exist in pure form but is present combined with minerals. In nature, the main source of fluorine is fluorite, cryolite and fluorapatite. Fluorite is the most abundant and China and Mexico are the largest suppliers of fluorite.

Physical Characteristics

Fluorine is a pale yellow, diatomic gas. It has a pungent smell. At lower temperature (-188°C), fluorine can condense to bright yellow liquid. There are two solid forms of fluorine, α- and β-fluorine. β-fluorine is transparent and soft while α- fluorine is hard and opaque, and form at -220 °C and -228 °C, respectively. It is highly flammable. Fluorine can be fluorescent under certain conditions. In liquid form, fluorine is readily soluble in liquid oxygen and ozone [3].

Chemical Characteristics

Fluorine is highly reactive gas. It is highly corrosive in nature. Fluorine has the second highest electron affinity. Fluorine has the highest electronegativity among all elements. It has the third highest first ionization energy among all elements, that is why it is almost impossible to remove electrons from its valence shell. Fluorine forms very strong bonds with other elements. Cold fluorine gas can react with unreactive substances, such as glass, powdered steel. Water and wood can catch fire when exposed to pressurized fluorine. Alkali metals reacts vigorously with fluorine and can cause explosions. Gases, such as sulfur dioxide and hydrogen sulfide readily combine with fluorine. Fluorine reacts explosively with hydrogen gas. Oxygen does not react with fluorine at room temperatures. Halogens also readily react with fluorine [4].

Significance and Uses

  • The largest consumption of fluorine is in the making of UF6 for nuclear fuel chain (fluorination of uranium tetrafluoride).
  • A huge proportion of fluorine is used in the preparation of SF6 (an inert dielectric compound) that is used in making high-voltage transformers and circuit breakers.
  • Various compounds of fluorine are used in various electronics, such as in cleaning equipment.
  • Fluoride in the form of inorganic compounds is used in glass etching and steel picking.
  • Organofluorides are used in making refrigerant gases (Freons) and surfactants. They are used as solvents, propellants and in air conditioning systems.
  • Fluorine is used to make agrichemicals, including fungicides and herbicides.
  • Fluorination of water has been done since 1940s to fight tooth decay.
  • Fluorine-19, the stable isotope is used in magnetic resonance imaging (MRI).
  • Fluroine-18 is used as a radioactive tracer in various tomographic techniques to detect tumors.
  • Fluorine is present in various pharmaceuticals, such as Lipitor, that is a cholesterol reducing drug. Similarly, Seretide, an asthma prescription is a widely used drug and contain fluorinated compound called the, fluticasone which is as the active agent.
  • Fluorine is used in making of various steroids and antibiotics.

Health Hazards

Fluorine in elemental form is highly toxic to living organisms. It is considered more dangerous than hydrogen cyanide. Fluorine causes significant irritation in the respiratory system and eyes at about 100 ppm. Ingestion of fluorine can lead to liver and kidney damage. Ingestion of fluorine in concentration above 25ppm and inhalation of 1000 ppm is considered as potentially lethal dose [5]. Hydrofluoric acid can lead to severe tissue damage upon inhalation, ingestion or direct contact. If enter in blood, it can adversely react with magnesium and calcium and lead to life threatening situation. Fluorine is also part of a notorious group of compounds, known as the chlorofluorocarbons (CFCs), which damage the ozone layer and are released in to the environment via various anthropogenic activities.

 Isotopes of Fluorine

In nature, there is only one stable isotope of fluorine, fluorine-19. It is quite abundant and its magnetogyric ratio is quite high. There are seventeen artificial isotopes of fluorine, and their mass number range from 14 to 31. Fluorine-18 is the most stable artificial isotope. All other isotopes of fluorine are radioactive [6].

 

REFERENCES

[1]. Ampère 1816

[2]. Jaccaud et al. 2000, p. 384

[3]. http://www.elementalmatter.info/fluorine-properties.htm

[4]. Riedel & Kaupp 2009.

[5]. Keplinger & Suissa 1968.

[6]. National Nuclear Data Center & NuDat 2.1, Fluorine-19.

Helium

Helium is the second most abundant element in the universe and is primarily produced because of radioactive decay. It was discovered in 1868 by Jules Janssen and Norman Lockyer.

 History and Discovery

The discovery of helium is linked with its primary sources, the sun. In the solar eclipse of 1868, several scientists studied the spectral lines coming from the sun and observed the presence of distinct and unknown yellow lines. These lines were named “Helium” by Norman Lockyer, from the word Helios, which is the name of Greek God of the Sun [1]. Helium was isolated by William Ramsay in 1895. Helium was used as lifting gas for air craft in World War I and in World War II, helium was produced widely and used in welding of shielded arc and lifting gas.

Helium

Periodic Table ClassificationGroup 18
Period 1
State at 20CGas
ColorColorless gas
Electron Configuration1s2
Electron Number2
Proton Number2
Electron Shell2
Density0.18 g.cm-3 at 20°C
Atomic number2
Atomic Mass4.00 g.mol -1
Electronegativity according to PaulingN/A

Occurrence

Helium is present rarely on the Earth. However, it is the 2nd most abundant element in the universe [2]. The atmospheric content of helium on the Earth is only 5.2 ppm [3]. There is a continuous production of helium on the Earth (via radioactive decay), but it readily escapes the Earth’s atmosphere and enter the space. Helium is the most abundant gas in the Earth’s heterosphere (the layer of atmosphere around 80km above Earth). The primary source of helium on Earth are minerals of thorium and uranium, which emit alpha particles (helium nuclei). About 3000 metric tons of helium are produced annually in the lithosphere (the layer of Earth including upper mantle and the crust) [4]. USA has been the biggest producer of helium since its discovery, but till 2012, most of the helium reserves have been exhausted and now account for 30% of the world’s helium supply. New reserves of helium have been discovered in North America. Russia and Qatar have also developed helium production plants.

Physical Properties

Helium is a colorless and odorless gas. It has the lowest melting point among all elements. Its boiling point is close to absolute zero. Helium is hardly soluble in water and is in fact, the least soluble monoatomic gas. Helium is present in plasma state on Earth. The atomic state of helium is predominant in the outer earth and space. There is a considerable difference between the two states of helium. The plasma state of helium has high electrical conductivity. Helium is also highly affected by the magnetic field and interacts with Earth’s magnetosphere to from aurora. Helium remains in liquid form at absolute zero, at standard pressure. This occurs due to quantum mechanics as the zero-point energy of helium is significantly high and does not allow freezing. Helium solidifies at -272C at a pressure of 2.5 MPa (25 bar). The refractive index of solid and liquid helium is almost the same and it is difficult to distinguish between the two states. Helium can solidify at higher temperature under high pressure.

Chemical Properties

Helium is an inert (Nobel) gas under all standard conditions [5]. In plasma form or subjected to electron bombardment, helium can from unstable compounds with certain metals including, sulfur, iodine, phosphorus and tungsten. These compounds are termed as excimers. Helium can exist in molecular ion form, such as HeH+, which is a highly stable but reactive form of helium. Other compounds, known as Van der Waals compounds of helium are also formed with lithium and cryogenic helium gas. Under high pressure, however, helium can form various compounds, for example, helium-nitrogen clathrate (He(N2)11).

Significance and Uses

  • Helium is widely used as coolant in magnetic resonance imaging (MRI) for medical scanning.
  • It is used in supersonic wind tunnels.
  • Helium is used in arc welding processes.
  • Helium is used as an ideal gas for filling of balloons, and airships, as it is lighter than air and is non-flammable.
  • Helium is used to detect cracks in buildings with high-vacuum rooms, and high-pressure containers, as it can efficiently diffuse through solids.
  • Helium is used as carrier gas in chromatography technique.

Health Hazards

Helium is non-toxic but can lead to poisonous effects if inhaled in high concentrations, it can lead to suffocation and death. Minor inhalation of helium can lead to change in the voice (becomes reedy) of individual, as sound waves travel faster in helium as compared to air. This effect is short term and not dangerous [6].

Isotopes of Helium

There are nine isotopes of helium, and only two are stable, helium-3 and helium-4. The two stable isotopes occur in a ratio of 1:1000,000 in the Earth’s atmosphere. The most abundant isotope (helium-4) is produced because of alpha decay of radioactive elements, such as uranium. Helium-3 is present in scarce amount on Earth and is formed by beta decay of tritium. Helum-4 is highly stable isotope. It is also formed as a result of Big Bang nucleosynthesis. Helim-3 is found in abundance in stars. It is also present on the surface of the Moon surface. Artificial isotopes of helium are also present, including helium-6, helium- and helium-8, where helium-7 and helium-8 are produced during various nuclear reactions [7].

 

 

REFERECNES

[1]. Harper, Douglas. “helium”. Online Etymology Dictionary.

[2]. Emsley, John (2001). Nature’s Building Blocks. Oxford: Oxford University Press. pp. 175–179. ISBN 978-0-19-850341-5.

[3]. “The Atmosphere: Introduction”. JetStream – Online School for Weather. National Weather Service. 2007-08-29. Archived from the original on January 13, 2008. Retrieved 2008-07-12.

[4]. Morrison, P.; Pine, J. (1955). “Radiogenic Origin of the Helium Isotopes in Rock”. Annals of the New York Academy of Sciences. 62 (3): 71–92. Bibcode:1955NYASA..62…71M. doi:10.1111/j.1749-6632.1955.tb35366.x.

[5]. Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5.

[6]. Ackerman, M. J.; Maitland, G. (1975). “Calculation of the relative speed of sound in a gas mixture”. Undersea Biomed Res. 2 (4): 305–10. PMID 1226588. Retrieved 2008-08-09.

[7]. “Helium Gas Safety & Data Sheet”. bouncetime

 

 

Copper

Copper is relative inexpensive metal and widely used since old civilization. It is an excellent conductor of heat and electricity. Copper resist corrosion and is widely used in making various alloys.

 Discovery and History

Copper has been known from prehistoric times and Neolithic humans used copper as stones as early as 8000 BCE. Copper was the first metal that was smelted from ores in 5000 BC and later used in pottery in North Africa. Early societies used it in place of gold and silver for making decorative items and ornaments [1]. Later, bronze which is an alloy of copper and tin was introduced between 3500 to 2500 BC in West Asia and Europe. In the Temple of King Sa’H-Re in Abusir, copper tubes for conveying water were used in 2750 BC. The name copper has been originated from Cyprium, which is Latin for metal of Cyprus. The term copper was introduced for the first time in 1530. The symbol of copper is Cu, derived from cuprum.

Copper

Periodic Table ClassificationGroup 11
Period 4
State at 20CSolid
ColorRed-orange metallic luster
Electron Configuration[Ar] 3d10 4s1
Electron Number29
Proton Number29
Electron Shell2, 8, 18, 1
Density8.96 g.cm-3 at 20°C
Atomic number29
Atomic Mass63.55 g.mol -1
Electronegativity according to Pauling1.90

Occurrence

Copper is widely present in many parts of world in combined state and free state. In combined form it exit as chalcocite (sulfide mineral), chalcopyrite (copper +iron sulfide), bornite (copper+ iron ore), cuprite (oxide mineral), malachite (copper carbonate) and azurite (copper carbonate) [2]. It is also present in the ashes of sea weeds and in sea corals. Copper is also present in human liver. In invertebrates it is present in many mollusks and arthropods. Andes Mountains in Chile is the greatest known deposit of copper mineral. Other major producers are Peru, China and the United States. Commercially copper is produced through smelting, followed by electrodeposition from sulfate solutions.

Physical characteristics

Fresh Copper has pinkish color but soon convert into reddish orange color due to direct exposure with oxygen. Copper oxidizes in the air and exhibit green color that’s why roof of building looks green. Copper is flexible and soft due to which it can be stretched into wires easily [3]. Copper dissolve in a mixture of hydrogen peroxide and hydrochloric acid to form copper chloride. Copper’s atomic number is 29 and its atomic mass is 63.54g/mol. Its melting point is 1083oC and boiling point is 2595oC. Copper is very dense in nature its density at 20oC is 8.9 g/cm3. Copper is biostatic in nature that means no bacteria and other forms of life can grow on it. Various alloys of copper also have antimicrobial properties.

Chemical characteristics

Copper has low chemical activity, it slowly reacts with oxygen and form a layer of brown black copper oxide that protects the underlying metal from further corrosion. Copper compound exist in two oxidation state +1 and +2. +2 compounds are blue in color. +1 compounds are white in color. They are weak oxidizing agent. Copper (I) compounds are weak reducing agents they react with air and make copper (II) compounds. They are not dissolve in water. Copper (II) are stable in air than copper(I) compounds. Gases are soluble in molten copper helpful in mechanical and electrical properties of solidified metal. Copper forms many alloys by mixing with other metals, most common alloys are brass and bronze.

Isotopes

Copper has 29 isotopes 63Cu and 65Cu are stable in nature. Other isotopes are radioactive in nature, 67Cu has half-life 61.83 hours.

Uses and significance

  • Copper is frequently used in wires, as it is an excellent conductor of electricity [4].
  • Copper has been used in making sculpture, it was also used in the construction of Statue of Liberty.
  • Copper is also used in photographic techniques.
  • Copper is used as fungicide in agriculture.
  • Copper is very important in countless types of electrical equipment.
  • Electrical devices rely on copper wiring due to its inherent properties.
  • Copper is corrosion resistant and present in weatherproof architectural materials.
  • Various alloys of copper are widely used in making jewelry.
  • Copper is used in textile industry for making of antimicrobial protective fabrics.
  • In past, copper chloride has been used to treat fever, arthritis and sciatica.

Dietary recommendations.

U.S recommended dose of copper is about 1.4 to 2.1mg per kg body mass.

Health effects

Copper is helpful in facilitating iron uptake that’s why its deficiency can lead to anemia[3]. Too much copper in diet also cause various problems. Human get copper mostly in the form of food and vitamin supplements. Various genetic disorders can affect the ability of body to use copper properly. Intake of copper is helpful to prevent cardiovascular diseases and osteoporosis. Copper enables body to make red blood cells.

 

References

  1. http://www.rsc.org/periodic-table/element/29/copper
  2. https://en.wikipedia.org/wiki/Copper#Chemical
  3. https://www.britannica.com/science/copper
  4. https://www.ehow.com/info_8601006_10-uses-copper.html

 

 

 

 

 

Iron

Iron is a chemical element with symbol Fe and atomic number 26. It is a metal in the first transition series. It is by mass the most common element on Earth, forming much of Earth’s outer and inner core. It is the fourth most common element in the Earth’s crust.

Discovery and History

The first signs of use of iron come from artifacts of the Sumerians and Egyptians dated to around 4000 B.C.E. They prepared tips of spears, daggers, and ornaments from iron recovered from meteorites. Because meteorites fall from the sky, some linguists have conjectured that the English word iron (Old English īsern), which has cognates in many northern and western European languages, derives from the Etruscan aisar, which means “the gods”[1]. Some have linked the iron in meteorites to a verse in the Quran(57:25) that says, “… and We sent down iron in which is incredible strength and many benefits for mankind.” Iron objects have been found in Egypt around 3500 BC. They contain about 7.5% nickel, which indicates that they were of meteoric origin. The ancient Hittites of Asia Minor, today’s Turkey, were the first to smelt iron from its ores around 1500 BC and this new, stronger, metal gave them economic and political power.

Iron

Periodic Table ClassificationGroup 8
Period 4
State at 20CSolid
ColorMetallic - gray
Electron Configuration[Ar] 3d6 4s2
Electron Number26
Proton Number26
Electron Shell 2, 8, 14, 2
Density7.87 g.cm-3 at 20°C
Atomic number26
Atomic Mass55.84 g.mol -1
Electronegativity according to Pauling1.83

Occurrence

Iron is an abundant element in the universe; it is found in many stars, including the sun. Iron is the fourth most abundant element in the earth’s crust, of which it constitutes about 5% by weight, and is believed to be the major component of the earth’s core. Iron is found distributed in the soil in low concentrations and is found dissolved in ground water and the ocean to a limited extent. It is rarely found uncombined in nature except in meteorites, but iron ores and minerals are abundant and widely distributed [2].
The principal ores of iron are hematite (ferric oxide, Fe 2O 3) and limonite (ferric oxide trihydrate, Fe 2O 3·3H 2O). Other ores include siderite (ferrous carbonate, FeCO 3), taconite (an iron silicate), and magnetite (ferrous-ferric oxide, Fe 3O 4), which often occurs as a white sand. Iron pyrite (iron disulfide, FeS 2) is a crystalline gold-colored mineral known as fool’s gold. Chromite is a chromium ore that contains iron. Lodestone is a form of magnetite that exhibits natural magnetic properties [3]. Its abundance in rocky planets like Earth is due to its abundant production by fusion in high-mass stars, where it is the last element to be produced with release of energy before the violent collapse of a supernova, which scatters the iron into space.

Physical properties

The Latin name for iron is ferrum, which is the source of its atomic symbol, Fe atomic number 26; atomic weight 55.845; melting point about 1,535°C; boiling point about 2,750°C and specific gravity of 7.87 at 20°C. Like the other group 8 elements, iron exists in a wide range of oxidation states, −2 to +7, although +2 and +3 are the most common. Fresh iron surfaces appear lustrous silvery-gray, but oxidize in normal air to give hydrated iron oxides, commonly known as rust.

Chemical Characteristics

In chemical terms, it is classified as a transition metal. It is located in period 4 of the periodic table, situated between manganese and cobalt. In addition, it lies at the top of group 8 (former group 8B). Iron, cobalt, and nickel have a number of similar properties and were once grouped together as group 8B [4]. Iron has the ability to form variable oxidation states differing by steps of one and a very large coordination and organometallic chemistry.

Significance and Uses

  • Iron (III) oxide: ferric oxide, or red iron oxide (Fe2O3): This compound corresponds to iron rust. Its mineral form, known as hematite, is mined as the main ore of iron and is used in the production of iron in a blast furnace.
  • Iron (III) oxide-hydroxide, hydrated iron oxide, or yellow iron oxide (FeO(OH)): This solid material has colors ranging from yellow through dark brown to black. It occurs naturally as minerals goethite, feroxyhyte, limonite, and lepidocrocite. It is used in aquarium water treatment as a phosphate binder.
  • Iron (II) sulfate or ferrous sulfate (FeSO4): It is commonly used additive that is found in various foods. Ferrous sulphate is used to treat anemia (iron-deficiency).
  • The most common use of iron is in manufacturing of steel, that has various attractive properties and uses.
  • Cast iron (with 3-5% carbon) is used for making pumps, pipes and valves.
  • Iron and steel are widely used in civil engineering and construction.

Health Effects

Iron can be found in meat, whole meal products, potatoes and vegetables. The human body absorbs iron in animal products faster than iron in plant products. Iron may cause conjunctivitis, choroiditis, and retinitis if it contacts and remains in the tissues. Chronic inhalation of excessive concentrations of iron oxide fumes or dusts may result in development of a benign pneumoconiosis, called siderosis. Inhalation of excessive concentrations of iron oxide may enhance the risk of lung cancer development in workers exposed to pulmonary carcinogens [5].

Isotopes of Iron

Naturally occurring iron consists of four isotopes: 5.85 percent of slightly radioactive 54Fe (half-life >3.1×1022 years), 91.75 percent of stable 56Fe, 2.12 percent of stable 57Fe, and 0.28 percent of stable 58Fe. In addition, it appears that the naturally occurring radioactive isotope 60Fe, with a half-life of 1.5 million years, is now extinct, but it can be produced synthetically. Much of the past work on measuring the isotopic composition of iron centered on determining 60Fe variations due to processes accompanying nucleosynthesis (that is, through meteorite studies) and ore formation.

Reference

1. Rick McCallister and Silvia McCallister-Castillo, A-AL Etruscan Glossary (1999). Retrieved October 22, 2011.
2. http://www.rsc.org/periodic-table/element/26/iron
3. https://www.infoplease.com/encyclopedia/science-and-technology/chemistry/compounds-and-elements/iron/natural-occurrence
4. http://www.newworldencyclopedia.org/entry/Iron#cite_note-0
5. https://www.lenntech.com/periodic/elements/fe.htm#ixzz5SuB3O600