Thorium

Thorium is a slightly radioactive element discovered by Jons Jacob Berzelius in 1828.  It is three times more abundant than Uranium and can meet the world’s energy demand by being used as fuel in nuclear reactors.

History and Discovery

In 1828, an amateur mineralogist named Morten Thrane Esmark found a black mineral on an island, in Norway. He found it interesting and sent it to his father who forwarded it to, a Swedish chemist, Jons Jacob Berzelius for examination. Berzelius managed to isolate the new element thorium from the sent mineral sample. Thorium was named after a Germanic god called Thor, the god of thunder, and the mineral from which it was extracted is now known as thorite [1]. Thorium was isolated in metallic form in 1914 by Dutch entrepreneurs Lodewijk Hamburger and Dirk Lely Jr. It was used in gas mantle till electricity became extensively available around the world. Thorium was discovered to be radioactive in 1898 by German Chemist Gerhard Carl Schmidt and by Marie Curie, two years after the discovery of radioactivity of uranium. Radioactive decay of natural thorium is the major contributor to Earth’s internal heat.

Thorium

Periodic Table ClassificationGroup n/a
Period 7
State at 20CSolid
ColorSilvery, often with black tarnish
Electron Configuration[Rn] 6d2 7s2
Electron Number90
Proton Number90
Electron Shell2, 8, 18, 32, 18, 10, 2
Density11.72 g.cm-3 at 20°C
Atomic number90
Atomic Mass232.04 g.mol -1
Electronegativity according to Pauling1.30

Occurrence

Thorium is a primordial element which still naturally occurs in large quantities in the earth’s crust. It is found in minor quantity in soil and most rocks. It is a radioactive element, like uranium, but it is three times more abundant than uranium. Around four-fifth of the thorium produced at the time of formation of the earth still exists due to its long half-life [2]. It is the rarest primordial element in the universe as it is only produced during r process (core collapse supernovae).

Physical Characteristics

Thorium is a bright silvery metal and corrodes to black when exposed to air. Thorium belongs to actinide series in the periodic table. It makes a face centred cubic crystal structure at room temperature. Thorium is very ductile in pure form and can be forged easily. It is malleable and its bulk modulus is same as that of tin. Thorium is less dense when compared to its nearby elements, for instance, uranium. It has a high melting point of 1750 degree centigrade and a fifth highest boiling point (of 4788°C) amongst the known elements’ boiling point.   It is represented by symbol Th and has atomic number 90.

Chemical Characteristics

Thorium is a weakly radioactive element. It is electropositive element with four valence electrons. It is quite reactive and can ignite spontaneously if exposed to air when in finely divided form and forms thorium dioxide. Thorium dioxide is used as refractory material because it has very high melting point. At standard temperature and pressure thorium reacts with water.

Significance and Uses

  • Thorium element is abundant and can satisfy world’s energy demand. It is used as fuel in nuclear reactors, as a replacement for uranium.
  • Thorium is used as alloying element in making welding electrodes.
  • It is added to improve the mechanical strength of magnesium.
  • Thorium samples are purified to extract daughter nuclides which are used in cancer therapy.
  • Thorium is used to mark impurities in evacuated tubes because of its reactivity with air.

Health Effects

Thorium decays relatively slowly and emits alpha radiation which cannot penetrate human skin. Hence exposure to small amount of thorium is considered safe. The decay products of thorium include radium and radon which are dangerous radionuclides. Exposure to thorium containing dust can cause lung cancer, blood cancer, and liver or pancreas diseases.  Some thorium compounds are toxic.  Thorium metal ignites spontaneously in air so should be handled with care.

Isotopes of Thorium

All isotopes of thorium are unstable. Th-232 is the most stable and abundantly existing isotope with half-life approximately equal to the age of universe and about three times the age of the earth. Thirty radioisotopes of thorium have been characterized, their atomic masses ranging from 209 to 238.  The half-life of these radioactive isotopes ranges from thousands of years to less than ten minutes [3].

REFERENCES

[1]. Thomson, T. (1831). A System of Chemistry of Inorganic Bodies. 1. Baldwin & Cradock and William Blackwood. p. 475

[2]. Audi, G.; Bersillon, O.; Blachot, J.; et al. (2003). “The NUBASE evaluation of nuclear and decay properties” (PDF). Nuclear Physics A. 729: 3–128. Bibcode:2003NuPhA.729….3A. doi:10.1016/j.nuclphysa.2003.11.001. Archived from the original (PDF) on 24 July 2013.

 [3]. Ikezoe, H.; Ikuta, T.; Hamada, S.; et al. (1996). “alpha decay of a new isotope of 209Th”. Physical Review C. 54 (4): 2043–2046. Bibcode:1996PhRvC..54.2043I. doi:10.1103/PhysRevC.54.2043.

 

 

Rubidium

Rubidium was discovered in 1861. It belongs to the alkali metals group of the periodic table. It is an abundant and highly reactive metal.

History and Discovery

Rubidium was discovered by Gustav Kirchhoff and Robert Bunsen in 1861. They used the technique known as flame spectroscopy to isolate the new element. The name rubidium has been derived from Latin word rubidus, that means deep red. It was given to the element as it emitted bright red lines in its emission spectrum [1]. The radioactivity of rubidium was discovered in 1908. The discoverers of Bose-Einstein condensate in 1995 used rubidium-87 and were awarded the Nobel prize in Physics in 2001 [2].

Rubidium

Periodic Table ClassificationGroup 1
Period 5
State at 20CSolid
ColorGrey white
Electron Configuration[Kr] 5s1
Electron Number37
Proton Number37
Electron Shell2, 8, 18, 8, 1
Density1.63 g.cm-3 at 20°C
Atomic number37
Atomic Mass85.47 g.mol -1
Electronegativity according to Pauling0.82

Occurrence

Rubidium is an abundant element and is ranked as the twenty-third most abundant element in the Earth’s crust [3]. In mostly occurs in the form of minerals, including carnallite, leucite, zinnwaldite and pollucite. The commercial production of rubidium is carried out from lepidolite, which contains up to 3.5% of rubidium [4]. Various minerals of potassium and potassium chlorides also contain significant amounts of rubidium. Rubidium is also present in the sea water, with a concentration of 125 µg/L which is much lower as compared to potassium. The largest producers of rubidium which have large deposits of the metal include Canada and Italy [5].

Physical Characteristics

Rubidium is silvery-white metal. It belongs to the alkali metal group of elements and is soft in nature. Rubidium is ductile metal and used for various purposes. It has a melting point of 39.3°C and have a low density, 1.532 g/cm3.

Chemical Characteristics

Rubidium is a very reactive metal. It undergoes rapid oxidation in the presence of air. Rubidium burns with a purplish flame like potassium. It has high electropositive values and is ranked second among the stable alkali metal. Rubidium reacts vigorously with water and can lead to the ignition of hydrogen gas that is produced during the reaction. It also undergoes spontaneous ignition when exposed to air. It forms alloys with iron, gold, sodium and potassium and forms amalgams with mercury. The ionization energy of rubidium is very low, 406 kJ/mol. The most common and widely used compound of rubidium is rubidium chloride (RbCl). There are various oxides of rubidium, and forms superoxide when excess amount of oxygen is present.

Significance and Uses

  • Rubidium is widely used in the manufacturing of electronic devices.
  • Rubidium is used to make purple colored fireworks.
  • It is used to make thermoelectric generators.
  • Rubidium is used in making photocells, oscillators and vacuum tubes.
  • Rubidium is used in the manufacturing of special type of glass.
  • Rubidium-82 is used for medical purposes for the diagnosis of various diseases, such as myocardial perfusion and detection of brain tumors.

Health Effects

Rubidium is a non-toxic metal. Rubidium has no biological role. However, the similarity of charge between rubidium and potassium ions makes it taken up by the cells in similar ways as potassium. Due to its vigorous reaction with water and its ability to spontaneously catch fire, the handling and storage of rubidium is quite challenging. Rubidium can be taken up by cells of the body, but these ions are not poisonous. The biological half-life of rubidium is around 50 days.

Isotopes of Rubidium

There are two naturally occurring isotopes in rubidium, rubidium-85 is the stable and more abundant isotope, while rubidium-87 is the radioactive isotope. Rubidium-87 has a half-life of around 49 billion years and is considered as a primordial nuclide. There are twenty-four artificial radioactive isotopes of rubidium, which have a half-life of less than 90 days [6].

REFERENCES

[1]. Weeks, Mary Elvira (1932). “The discovery of the elements. XIII. Some spectroscopic discoveries”. Journal of Chemical Education. 9 (8): 1413–1434. Bibcode:1932JChEd…9.1413W. doi:10.1021/ed009p1413.

[2]. Levi, Barbara Goss (2001). “Cornell, Ketterle, and Wieman Share Nobel Prize for Bose-Einstein Condensates”. Physics Today. Physics Today online. 54 (12): 14. Bibcode:2001PhT….54l..14L. doi:10.1063/1.1445529.

[3]. Butterman, William C.; Brooks, William E.; Reese, Jr., Robert G. (2003). “Mineral Commodity Profile: Rubidium” (PDF). United States Geological Survey. Retrieved 2010-12-04

[4]. Wise, M. A. (1995). “Trace element chemistry of lithium-rich micas from rare-element granitic pegmatites”. Mineralogy and Petrology. 55 (13): 203–215. Bibcode:1995MinPe..55..203W. doi:10.1007/BF01162588.

[5]. Teertstra, David K.; Cerny, Petr; Hawthorne, Frank C.; Pier, Julie; Wang, Lu-Min; Ewing, Rodney C. (1998). “Rubicline, a new feldspar from San Piero in Campo, Elba, Italy”. American Mineralogist. 83 (11–12 Part 1): 1335–1339. Bibcode:1998AmMin..83.1335T. doi:10.2138/am-1998-11-1223.

[6]. Audi, Georges; Bersillon, O.; Blachot, J.; Wapstra, A. H. (2003). “The NUBASE Evaluation of Nuclear and Decay Properties”. Nuclear Physics A. Atomic Mass Data Center. 729 (1): 3–128. Bibcode:2003NuPhA.729….3A. doi:10.1016/j.nuclphysa.2003.11.001.

Rhodium

Rhodium was discovered in 1803 and is member of the platinum group. It is widely used as a catalytic converter in automobile fuel industry that was introduced by Volvo in 1976.

History and Discovery

Rhodium was discovered by William Hyde Wollaston in 1803 [1]. He extracted the element from its platinum ore that was obtained from South Africa. The name rhodium has been derived from Greek word rhodon that means rose. The new element was named rose due to the colored compound formed when compound of rhodium with chloride was dissolved in aqua regia.

Rhodium

Periodic Table ClassificationGroup 9
Period 5
State at 20CSolid
ColorSilvery white metallic
Electron Configuration[Kr] 4d8 5s1
Electron Number45
Proton Number45
Electron Shell2, 8, 18, 16, 1
Density12.41 g.cm-3 at 20°C
Atomic number45
Atomic Mass102.91 g.mol -1
Electronegativity according to Pauling2.28

Occurrence

Rhodium is extremely rare metal and is present in around 0.0002 parts per million of the earth’s crust. The minerals and ores of rhodium are not very common in the earth’s crust. Rhodium is found in free form in nature, as well as alloyed with gold and other metals of the platinum group. Rhodium is also present in form of minerals such as rhodplumsite and bowieite. Commercially, rhodium is extracted from its nickel and platinum ores. It is also produced as a byproduct during the uranium-235 fission reaction, however, this extraction is complex and makes the commercial level production through this method fairly impossible. The largest producers of rhodium include South Africa, Canada and Russia, where large natural deposits of rhodium metal are present [2].

Physical Characteristics

Rhodium is silvery white transition metal. It has a shiny appearance and is hard in nature. Rhodium is a Nobel metal and belongs to the group of exquisite metals (including rhodium, platinum, osmium, ruthenium, osmium and iridium) termed as the platinum group metals (PGMs). Rhodium is a durable metal. It does not react with oxygen even at high temperatures. That is why, rhodium is highly resistant to corrosion and tarnishing. Its density is lower than platinum, around 12.41 g/cm3. Rhodium has a higher melting point as compared to platinum, 1964 °C and have a boling point of 3695 °C.

Chemical Characteristics

Rhodium is non-reactive, chemically inert element. It is not dissolved in dilute nitric acid, or aqua regia [3]. Rhodium does not react with oxygen and that is why is resistant to tarnishing in air. The most common oxidation state of rhodium is +3, but other oxidation states such as +0 and +6 also exist. Compounds of rhodium resemble compounds of platinum. It does not form volatile compounds. It reacts with halogens to form halides with various oxidation states. Most of the oxides of rhodium are stable.

Significance and Uses

  • About 80% of rhodium that is extracted worldwide is used as a catalytic converter, for the catalytic conversion of harmful gases from automobile exhaust, carbon monoxide and hydrocarbon into less harmful gasses, such as nitrogen, carbon dioxide and water. Rhodium specifically reduced the amount of NOx in the automobile exhaust.
  • Rhodium is widely used in making of various alloys to improve the hardness and corrosion resistance of other metals, such as platinum and palladium.
  • Rhodium is also used to increase the beauty and shine of various precious metals, such as silver is lined with rhodium to make it resistant to tarnishing and discoloration and it is used as an additional attractive plating on gold to enhance its shine and appearance.
  • Rhodium is used in nuclear power plants.
  • Rhodium is used in glass industry, to produce flat panel glass and fiber glass.
  • Rhodium is used in making ornaments and jewellery.

Health Hazards

In free form, rhodium is an inert and nontoxic metal. However, various compounds and salts f rhodium have moderate toxicity. Individuals working with rhodium are subjected to occupational hazards and the Occupational Safety and Health Administration (OSHA) has determined the permissible exposure dose of for rhodium in the workplace at 0.1 mg/m3 over an eight-hour work duration. In mice, the median lethal dose of soluble rhodium has been estimated to be around 198mg/ kg of body weight.

Isotopes of Rhodium

Rhodium has one naturally occurring isotope, rhodium-103. The most stable radio-isotope is rhodium-101 that has a half-life of around 3.3 years. There are twenty other radioactive isotopes of rhodium that are artificially produced and have atomic masses that range from 92.926 u to 116.925 u. most of these isotopes have half-life of less than an hour, expect rhodium-100 that has a half-life of 20.8 hours and rhodium-35.36 with a half-life of 35.36 hours [4].

REFERENCES

[1]. Hammond, C. R. (2004). “The Elements”. Handbook of Chemistry and Physics (81st ed.). CRC press. ISBN978-0-8493-0485-9.

[2]. “Platinum-Group Metals” (PDF). Mineral Commodity Summaries. United States Geological Survey. January 2007.

[3]. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 1113. ISBN 0-08-037941-9.

 [4]. Audi, G.; 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.

 

Roentgenium

Roentgenium is an artificial and highly radioactive element that was synthesized in 1994. It is used for various research purposes.

History and Discovery

Roentgenium is a synthetic element. It was predicted by Mendeleev and was named eka-gold and later was named as the element 111.  Roentgenium was synthesized for the first time by a team of researchers led by Sigurd Hofmann in December 1994, who worked at GSI Helmholtz Centre for Heavy Ion Research, Germany. The bismuth-209 isotope was bombarded by nickel-62 which produced a single atom of isotope roentgenium-272 [1]. Using the same approach, three more atoms of roentgenium-272 were produced in 2002. The name roentgenium has been given in honor of the contributions of physicist Wilhelm Rontgen who discovered the X-rays.

Roentgenium

Periodic Table ClassificationGroup 11
Period 7
State at 20CSolid (predicted)
ColorSilvery (predicted)
Electron Configuration[Rn] 5f14 6d9 7s2 (predicted)
Electron Number111
Proton Number111
Electron Shell2, 8, 18, 32, 32, 17, 2 (predicted)
Density28.70 g.cm-3 at 20°C (predicted)
Atomic number111
Atomic Mass282.00 g.mol -1 (most stable isotope)
Electronegativity according to Paulingn/a

Occurrence

Roentgenium does not occur in nature. It is a synthetic element and is highly unstable. It has a half-life of few seconds.

Physical Characteristics

Roentgenium exists in solid form under normal conditions. It crystallizes in a cubic structure that has a body-centered symmetry. Roentgenium is thought to be a very dense metal and would have a density of round 28.7 g/cm3. But due to its unstable nature, the physical properties can only be predicted or theoretically calculated. Further, creating that much sample of roentgenium to study its physical and chemical properties is difficult as it would undergo spontaneous decay within minutes.

Chemical Characteristics

Roentgenium is extremely radioactive element. The chemical characteristics of roentgenium have not been studied in detail yet [2]. it has been predicted to be a Nobel element. And based in the oxidation’s states of other member of its group (Group 11), the common oxidation states of roentgenium include +5 and +3. It has similar reactivity as gold but is thought to produce more stable and diverse compounds. Roentgenium has also been predicted to form compounds with hydrogen, ammonia and phosphine. The extremely short half-lives of isotopes and highly volatile nature of its compounds make the statistically significant chemical analysis a challenge. Further, the rate of production of roentgenium isotopes need to be at least one per week in order to obtain considerable amount of the element.

Significance and Uses

  • Roentgenium and its various isotopes are used for research and laboratory purposes.

Health Effects

Roentgenium is highly radioactive and require special handling precautions.

Isotopes of Roentgenium

There are nine isotopes of roentgenium. There are no stable or natural isotopes of roentgenium. The isotopes are produced by decay of heavier element or by the fusion of nuclei of light elements. The isotopes have atomic masses: 272, 274, 278, 279, 280, 281, 282, 283, and 286. [3]. The heavier isotopes are more stable as compared to lighter ones. And the heaviest isotope, roentgenium-282 has a half-life of 2.1 minutes. All the isotopes undergo decay through alpha decay or spontaneous fission, and none of the isotopes undergoes beta decay.

REFERENCES

[1]. Hofmann, S.; Ninov, V.; Heßberger, F. P.; Armbruster, P.; Folger, H.; Münzenberg, G.; Schött, H. J.; Popeko, A. G.; Yeremin, A. V.; Andreyev, A. N.; Saro, S.; Janik, R.; Leino, M. (1995). “Production and decay of 269110″. Zeitschrift für Physik A. 350 (4): 277. Bibcode:1995ZPhyA.350..277H. doi:10.1007/BF01291181

[2]. Düllmann, Christoph E. (2012). “Superheavy elements at GSI: a broad research program with element 114 in the focus of physics and chemistry”. Radiochimica Acta. 100 (2): 67–74. doi:10.1524/ract.2011.1842.

[3]. Sonzogni, Alejandro. “Interactive Chart of Nuclides”. National Nuclear Data Center: Brookhaven National Laboratory. Retrieved 2008-06-06.

 

 

Cerium

Cerium was discovered in 1803 by Berzelius and Hisinger, and independently by Klaproth in the same year. It is widely used in ferrocerium lighters. Its readily forms oxide in the air and form CeO2.

History and Discovery

Cerium was discovered in 1803 by Jons Jakob Berzelius and Wilhelm Hisinger in Sweden, and independently by Martin Heinrich Klaproth in Germany in the same year. It was named by Berzelius after the dwarf planet Ceres, which was discovered two years earlier. Ceres is the name of a Roman Goddess of agriculture, crop, fertility and relationships [1]. Berzelius and Hisinger investigated the chemical properties of the new element and successfully prepared various salts of cerium. They proposed that cerium had two oxidation states: one formed colorless salts and other formed yellowish-red ones. Klaproth named its oxide ockroite due to its yellow color. In 1825, Carl G. Mosander worked with Berzelius to prepare metallic cerium. He isolated cerium from its chloride with the help of using potassium.

Cerium

Periodic Table ClassificationGroup n/a
Period 6
State at 20CSolid
ColorSilvery white
Electron Configuration[Xe] 4f1 5d1 6s2
Electron Number58
Proton Number58
Electron Shell2, 8, 18, 19, 9, 2
Density6.70 g.cm-3 at 20°C
Atomic number58
Atomic Mass140.12 g.mol -1
Electronegativity according to Pauling1.12

Occurrence

Cerium is the most abundant element among all lanthanides. It is present in concentration of about 66ppm in the earth crust. It is more abundant than lead and tin. Its content in the soil varies between 2 and 150ppm, and in 1.5 parts per trillion in sea water [2]. It is also present in minerals of the monazite (reddish brown phosphate mineral) and bastnasite (family of 3 carbonate-fluoride) groups. Cerium is easily extracted from its ores because in it is the only lanthanide that can acquire the oxidation state of +4 in an aqueous solution.

Physical Characteristics

Cerium is a silvery-white metal. It is soft and can be cut with a knife. Cerium ductile in nature and its hardness and elasticity are like silver. It is present in between lanthanides and actinides in the periodic table. It is present in four allotropic forms α, β, δ and γ. Both β and γ quite stable at room temperature. α- cerium is stable below -150oC. It has high density about 8.16g/cm3. δ- cerium exist above 726oC. Cerium chemical symbol is Ce, its atomic number is 58. Cerium atomic weight is 140.116. Cerium melting point is 795OC. Its boiling point is 3443oC. Its density at room temperature is about 6.770 g/cm3.  Liquid cerium is dense at atmospheric pressure than its solid form.

Chemical Characteristics

Cerium is tarnished when exposed to air. At 150OC, cerium readily burns in air and form pale yellow cerium (IV) oxide which is also known as ceria. With hydrogen gas it is reduced to cerium (III) oxide. It is pyrophoric (when it is ground the resulting product spontaneously catch fire). It efficiently conducts electricity. It is electropositive metal that reacts with water to form cerium (III) hydroxide and hydrogen gas. Cerium also reacts with halogens and forms trihalides. It is readily dissolved in dilute sulfuric acid and forms a colorless Ce3+ ions. Cerium is the only lanthanide element which exist in +4 oxidation state. Cerium has strong oxidizing property and oxidizes hydrochloric acid to produce chlorine gas. Cerium (IV) salts are used in cerimetry titrations [3].

Significance and Uses

  • Cerium dioxide is used as a catalyst for the combustion of thorium dioxide.
  • Cerium is used in making alloys with various metal, such as aluminum and iron. It is also used in the manufacturing of stainless steel as precipitation hardening agent.
  • Cerium is used as catalyst in the reduction of nitrogen oxides to nitrogen gas.
  • Cerium sulfide has replaced cadmium in red pigments for toys and household wares.
  • Cerium is also used in the manufacturing of flat screen television.
  • It is widely used as a catalyst to refine petroleum.
  • Cerium in the form of oxides is widely used in incandescent lanterns.
  • It is also used to make carbon arc lights, that are used in the motion pictures for studio lighting and projector lights.
  • Cerium oxides are used to refine and polish glass surface.
  • It is used as flint in cigarette and gas lighter.
  • Cerium oxides in the nanopowder form is mixed in diesel fuel to reduce the emission of fumes and is also helpful in improving the overall performance of automobile engines.

Health Effects

Cerium is present in many household equipment including energy saver bulbs, colored television, fluorescent lamps and glasses. Cerium is dangerous in work place because the cerium gas can be inhaled with the air and prolonged exposure can lead to embolism of lung. There is no biological role of cerium, however some studies have reported that certain salts of cerium can stimulate metabolic activity of the body.

Isotopes of Cerium

Cerium has four natural isotopes 136Ce, 138Ce, 140Ce and 142Ce, and only cerium-140 is theoretically the most stable and abundant isotope. The unstable isotopes have half-lives of: 136Ce has ˃ 3.8×1016 years, 138Ce has ˃ 1.5×1014 years. 142Ce has 5×1016 years. There are thirty-five radioactive artificial isotopes of cerium, which range in atomic masses from 119 u to 157 u.

REFERENCES

[1]. https://en.wikipedia.org/wiki/Cerium#History

[2]. Emsley, John (2011). Nature’s Building Blocks: An A-Z Guide to the Elements. Oxford University Press. pp. 120–125. ISBN978-0-19-960563-7.

[3]. Greenwood and Earnshaw, pp. 1238–9

 

 

 

Tennessine

Tennessine is a synthetic element that was discovered in 2010. It is highly radioactive and unstable element.

History and Discovery

According to the Mendeleev’s nomenclature of undiscovered elements, Tennessine was named as eka-astatine or element-117. Tennessine was synthesized by collaboration of an American-Russian team led by Yuri Oganessia working in Dubna, Russia in 2010 [1]. The Joint Institute for Nuclear Research (JINR) worked with Oak Ridge National Laboratory (ORNL) in Oak Ridge, Tennessee, USA, and successfully created the element 117. They bombarded berkelium-249 (element 97) with calcium-48 (element20) nuclei. The initial bombardment was carried out for 70 days produced an atom of Tennessine-293. Tennessine is considered as the most recent element to be discovered. Its name was confirmed as Tennessine by International Union of Pure and Applied Chemistry in 2016, after the state of USA, Tennessee. Its first isotope was created in 2011. It symbol is Ts.

Tennessine

Periodic Table ClassificationGroup 17
Period 7
State at 20CSolid (predicted)
ColorSemimetallic (predicted)
Electron Configuration[Rn] 5f14 6d10 7s2 7p5 (predicted)
Electron Number117
Proton Number117
Electron Shell2, 8, 18, 32, 32, 18, 7 (predicted)
Density7.1 - 7.3 g.cm-3 at 20°C (predicted)
Atomic number117
Atomic Mass294.00 g.mol -1 (most stable isotope)
Electronegativity according to Paulingn/a

Occurrence

Tennessine is an artificial element and does not exist in nature. In 2012, seven more atoms of tennessine were produced by the Dubna team [2].

Physical Characteristics

Tennessine is predicted to be a solid under normal conditions. Only minute amount of Tennessine-294 have been produced so far and it is not enough to carry out analysis of its physical and chemical characteristics.

Chemical Characteristics

The chemical characteristics of tennessine is not well studied yet. It is expected to be a volatile metal and would acquire low oxidation states in compounds, -1, +1, +3 and +5. Tennessine belongs to the group 17 of the periodic table which is comprised of the halogens. Due to its relativistic effects, it is presumed to differ in properties from the halogens. However, its ionization energy, boiling point and melting point will follow the trend of the halogens.

Significance and Uses

  • Tennessine is used for research purposes.

Health Hazards

Tennessine is a radioactive element and requires special precautions with handling and storage.

Isotopes of Tennessine

There are two isotopes of tennessine, tennessin-293 and Tennessine-294. They are unstable and unnatural. Tennessin-293 ha a half-life of only 14 milliseconds, while tennessine-294 has a half-life of 78 milli-seconds.  Tennessine-293 decay through emission of alpha particles into element moscovium, the element 115 [3].

REFERENCES

[1]. Yu. Ts. Oganessian et al., Phys. Rev. Lett., 2010, 104, 142502, 4 pages.

[2]. “Russian scientists confirm 117th element”. Sputnik. 2012-06-25. Retrieved 2012-07-05

 [3]. Office of Science, Nations Work Together to Discover New Element,.

 

 

 

Silicon

Silicon is the second most abundant element in earth’s crust. It was discovered in 1823 by Jöns Jacob Berzelius. Silicon has tremendous uses including manufacturing of ceramic, glass, synthetic polymers and is an essential part of integrated circuits.

History and Discovery

Compounds of silicon were used long before the discovery of silicon. Antoine Lavoisier (1787) tried reducing silica, an oxide of silicon, to isolate silicon but failed. Sir Humphry Davy, in 1808 named the element silicium but also failed to isolate the element. The element was given its present name, silicon, by Thomas Thomson in 1817. Gay Lussac and Thenard successfully prepared impure amorphous silicon in 1811 but they did not characterize it as a new element. In 1823, silicon was finally prepared in pure form by Jöns Jacob Berzelius and hence given credit for its discovery [1]. Crystalline form of silicon was prepared, 31 years later, by Deville in 1854.

Silicon

Periodic Table ClassificationGroup 14
Period 3
State at 20CSolid
ColorCrystalline, reflective with bluish-tinged faces
Electron Configuration[Ne] 3s2 3p2
Electron Number14
Proton Number14
Electron Shell2, 8, 4
Density2.33 g.cm-3 at 20°C
Atomic number14
Atomic Mass28.09 g.mol -1
Electronegativity according to Pauling1.90

Occurrence

Silicon is the second most abundant element present in the earth’s crust. It is the seventh most abundant element in the universe. Silicon is formed through the oxygen-burning process in stars. Silicon reacts with oxygen to make silicon dioxide or silicates. Silicate minerals make up over 90% of earth’s crust. Silicon is rarely found in pure form. Group of minerals composed of silicon and oxygen are named silica. Silica is mostly found in crystalline state. Silicon minerals make up 90% of the earth’s crust and it can be used industrially in its naturally occurring form which makes it cheap and easily available raw material.

Physical Characteristics

Silicon is a brittle and hard crystalline solid. It has blue-grey metallic lustre. Silicon, in comparison with neighbouring elements in the periodic table, is unreactive. The symbol for silicon is Si with atomic number 14. It has a very high melting and boiling point. At standard conditions silicon also makes a giant covalent structure like other group 14 elements of periodic table do.

Chemical Characteristics

At room temperature, pure silicon acts as an insulator. Silicon is a semiconductor at standard temperature and pressure. Silicon is inert in crystalline form at low temperatures. Its conductivity increases with high temperature. Silicon readily reacts with oxygen [2]. It reacts with air above 900-degree centigrade. Melted silicon becomes very reactive and has to be stored in unreactive, refractory material to avoid any chemical reaction.

Significance and Uses

  • Silicon minerals are used as structural compounds for instance as clays, silica sand, building mortar, stucco and building stones.
  • Silicon minerals are used in making concrete.
  • Silica is used to make fire brick (refractory brick) which is used in lining of furnace.
  • It is used in making whiteware ceramics such as soda lime glass and porcelain.
  • Silica is used in making optical fibre which has vast uses in telecommunications and computer networking.
  • It is used in making fibreglass and glass wool which are used for structural support and thermal insulation.
  • Silicon is used in making mechanical seals and waterproofing.
  • Waxes and high-temperature greases are made using silicon.
  • For medical purposes, silicon is used in breast implants and contact lenses.
  • Silicon is used in making superalloys.
  • Silicon is used for making silicon wafers which has wide applications in the semiconductor industry.
  • Silicon is also essential for human beings i.e. skin, nail, hair and bone density of human beings depends on the amount of silicon present.
  • Synthetic polymers called silicones are produced using silicon.
  • Solar cells, semiconductors detectors, transistors and other semiconductor devices used in computer industry are made using silicon.
  • Silicon is a crucial part of integrated circuits (ICs) which have vital importance in our electronic appliances, for instance, computers and cell phones [5].
  • Free silicon is used for casting of aluminium and steel refining industry.

Health Effects

Silicon is slightly hazardous. If crystalline silica is inhaled, it may lead to lung disease such as asthma or inflammation in upper lobes of lungs. Exposure of elemental silicon can cause eye or skin irritation.

Isotopes of Silicon

Silicon has three stable isotopes; Si-28, Si-29 and Si-30. Of these three naturally occurring isotopes Si-28 is the most abundant as it is produced in stars as well as during nuclear fusion reaction. The remaining two isotopes of silicon form only 7% of the naturally occurring silicon. So far twenty radioisotopes of silicon have been characterized. Most of these radioisotopes have half-life of few seconds only. Unstable isotopes of silicon decay to form aluminium or phosphorus isotopes.

REFERENCES

[1]. Weeks, Mary Elvira (1932). “The discovery of the elements: XII. Other elements isolated with the aid of potassium and sodium: beryllium, boron, silicon, and aluminum”. Journal of Chemical Education. 9 (8): 1386–1412.

[2]. Voronkov, M. G. (2007). “Silicon era”. Russian Journal of Applied Chemistry. 80 (12): 2190. doi:10.1134/S1070427207120397

[3]. Rahman, Atta-ur- (2008-09-24). “Silicon”. Studies in Natural Products Chemistry. 35. p. 856. ISBN 978-0-444-53181-0

[4]. Jugdaohsingh, R. (Mar–Apr 2007). “Silicon and bone health”. The Journal of Nutrition, Health and Aging. 11 (2): 99–110. PMC 2658806.

[5]. Cheung, Rebecca (2006). Silicon carbide microelectromechanical systems for harsh environments. Imperial College Press. p. 3. ISBN 978-1-86094-624-0

 

 

 

Rutherfordium

Rutherfordium was discovered in 1964 and resynthesized in 1969. It is an artificially prepared radioactive element. Due to its instability, it has no major commercial use, only used in laboratories for certain research purposes.

History and Discovery

Rutherfordium is a synthetic element. It was synthesized for the first time in 1964 by the team of scientists at Dubna, Russia, which was led by Georgy Flerov. They bombarded plutonium with neon ions. The scientist thought that they found isotope-259.  This discovery was not accepted and resynthesized in 1966, but they confirmed the 1964 results. The Dubna researchers suggested the name kurchatovium (ku) after Igor Kurchatov, Russian nuclear physicist.  In 1969, University of California, Berkeley led by Albert Ghiorso successfully synthesized rutherfordium through bombarding a californium target with carbon-12 and carbon-13 ions. They named the element rutherfordium after Earnest Rutherford, father of nuclear physics and nuclear chemistry [1].

Rutherfordium

Periodic Table ClassificationGroup 4
Period 7
State at 20CSolid (predicted)
ColorUnknown
Electron Configuration[Rn] 5f14 6d2 7s2
Electron Number104
Proton Number104
Electron Shell2, 8, 18, 32, 32, 10, 2
Density23.20 g.cm-3 at 20°C (predicted)
Atomic number104
Atomic Mass261.00 g.mol -1
Electronegativity according to Paulingn/a

Occurrence

Rutherfordium is not naturally present in the earth’s crust. It is prepared by bombardment and decay of heavy isotopes. Mostly made by bombarding plutonium-242 with accelerated neon ions and sometime bombarding californium-249 with carbon ions.

Physical Characteristics

Rutherfordium is solid under normal condition. It isa very heavy element and has a high density of about 23.2 g/cm3. Its melting point is very high, about 2100oC. Boiling point of rutherfordium is also very high, 5500oC. Its chemical symbol is Rf and the atomic number is 104. Its atomic weight is 267 g/mol [2].

Chemical Characteristics

Rutherfordium is a synthetic and highly reactive element. It is the first transactinide element (super heavy). Its chemical properties resemble with elements of group 4 in the periodic table especially Hafnium (Hf). Rutherfordium most stable oxidation state is +4 but it is expected to be able to form less stable compounds at +3 oxidation state. Rutherfordium reacts with halogens to form tetrahalides RfX4, they hydrolyzed when contact with water and form oxyhalides RfOX2. Tetrahalides are volatile solids and they exist in vapor phase as monomeric tetrahalides molecules.

Significance and Uses

  • Rutherfordium is artificially prepared in the laboratory that is why it has no commercial use.
  • It is used in laboratories for some research purpose.
  • It is highly reactive so it is used for nuclear power and some medicinal purpose.

Health Effects

Rutherfordium is a very unstable element. It has an extremely short half-life. That’s why still there is no known effect on human health.

Isotopes of Rutherfordium

Rutherfordium has no stable isotopes, the isotopes are also not naturally occurring. Radioactive isotopes of rutherfordium have been synthesized in the laboratory by fusing two elements and through decay of heavier elements. Lighter isotopes like 253Rf and 254Rf have shorter half-life about 50µs.256Rf and 258Rf and 260Rf are stable around 10 ms. While, 255Rf, 257Rf, 259Rf and 262Rf have half-life between 1 and 5 seconds. 261Rf have 1 minute, 265Rf have 1.5 minutes and 263Rf have 10 minutes half-life respectively. Heaviest isotope 267Rf have half -life of about 1.3 hours [3].

REFERENCES

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

[2]. https://en.wikipedia.org/wiki/Rutherfordium

[3]. Sonzogni, Alejandro. “Interactive Chart of Nuclides”. National Nuclear Data Center: Brookhaven National Laboratory. Retrieved 2008-06-06.

 

Darmstadtium

Darmstadtium is an artificial element that was synthesized in 1994. It is highly radioactive and a very dense metal.

History and Discovery

Darmstadtium is a synthetic element. It was predicted by Mendeleev and was named eka-platinum and later was named as the element 110. It was created by Peter Armbruster and Gottfried Munzenberg in 1994 in GSI Helmholtz Centre for Heavy Ion Research, Germany. The lead-208 was bombarded by nickel-62 which produced a single atom of isotope darmstadtium-269 [1]. Using the same approach, nickel-64 ions were bombarded and nine atoms of darmstadtium-271 were produced. The name darmstadtium has been given in honor of the city where it was discovered, Darmstadt, Germany [2]. Its symbol is Ds.

Darmstadtium

Periodic Table ClassificationGroup 10
Period 7
State at 20CSolid (predicted)
ColorUnknown
Electron Configuration[Rn] 5f14 6d8 7s2 (predicted)
Electron Number110
Proton Number110
Electron Shell2, 8, 18, 32, 32, 16, 2 (predicted)
Density34.80 g.cm-3 at 20°C (predicted)
Atomic number110
Atomic Mass281.00 g.mol -1 (most stable isotope)
Electronegativity according to Paulingn/a

Occurrence

Darmstadtium do not occur in nature as it is highly unstable and have a half-life of few seconds.

Physical Characteristics

Darmstadtium exists in solid form under normal conditions. It crystallizes in a cubic structure that has a body-centered symmetry. Darmstadtium is a very dense metal and would have a density of round 34.8 g/cm3. But due to its unstable nature, the physical properties can only be predicted or theoretically calculated. Further, creating that much sample of darmstadtium to study its physical and chemical properties is difficult as it would undergo spontaneous decay with in minutes.

Chemical Characteristics

Darmstadtium is extremely radioactive element. The chemical characteristics of darmstadtium have not been studied in detail yet [3]. The extremely short half-lives of isotopes and highly volatile nature of its compounds make the statistically significant chemical analysis a challenge. Further, the rate of production of darmstadtium isotopes need to be at least one per week in order to obtain considerable amount of the element. However, it reacts with fluorine to form a relatively less volatile compound, darmstadtium hexafluoride

Significance and Uses

  • Darmstadtium and its various isotopes are used for research and laboratory purposes.

Health Hazards

Darmstadtium is highly radioactive and require special handling precautions.

Isotopes of Darmstadtium

There are nine isotopes of darmstadtium. There are no stable isotopes of darmstadtium. As it is a synthetic element, none of its isotopes are stable. The isotopes are produced by decay of heavier element or by the fusion of two light nuclei. The isotopes have atomic masses: 267, 296,270,271, 279,280 and 281 [4]. The heavier isotopes are more stable as compared to lighter ones. And the heaviest isotope, darmstadtium-281 has a half-life of 11 seconds, while darmstadtium-279 has a half life of 0.18 seconds. All the isotopes undergo decay through alpha decay or spontaneous fission, and none of the isotopes undergoes beta decay.

REFERENCES

[1]. Hofmann, S.; Ninov, V.; Heßberger, F. P.; Armbruster, P.; Folger, H.; Münzenberg, G.; Schött, H. J.; Popeko, A. G.; Yeremin, A. V.; Andreyev, A. N.; Saro, S.; Janik, R.; Leino, M. (1995). “Production and decay of 269110″. Zeitschrift für Physik A. 350 (4): 277. Bibcode:1995ZPhyA.350..277H. doi:10.1007/BF01291181

[2]. Griffith, W. P. (2008). “The Periodic Table and the Platinum Group Metals”. Platinum Metals Review. 52 (2): 114–119. doi:10.1595/147106708X297486.

[3]. Düllmann, Christoph E. (2012). “Superheavy elements at GSI: a broad research program with element 114 in the focus of physics and chemistry”. Radiochimica Acta. 100 (2): 67–74. doi:10.1524/ract.2011.1842.

[4]. Sonzogni, Alejandro. “Interactive Chart of Nuclides”. National Nuclear Data Center: Brookhaven National Laboratory. Retrieved 2008-06-06.

 

Zinc

Zinc is a prehistoric metal and have been used as copper alloy since 3rd millennium BC. It is widely used today to make electrical batteries and for zinc electroplating.

History and Discovery

The history of zinc dates to around 3rd millennium BC, when it was used in the form of alloy with copper (brass). Traces of use of brass in various applications have been found in ancient civilizations of Iraq, Kalmykia, UAE, West Indies, Iran and Syria [1]. In the 9th century AD, a distillation process was developed in Rajasthan to obtain pure zinc. The commercial production of zinc started in the 12th century. Zinc was discovered in pure form in 1746 by German Chemist Andreas Sigismund Marggraf. And a detailed insight in the electrochemical properties of zinc was presented by Luigi Galvani and Alessandro Volta in 1800.Zinc was named as white snow by Alchemists, as it burned in air to form white compound. The word zinc has been derived from German word Zinke that means tooth.

Zinc

Periodic Table ClassificationGroup 12
Period 4
State at 20CSolid
ColorSilver-gray
Electron Configuration[Ar] 3d10 4s2
Electron Number30
Proton Number30
Electron Shell2, 8, 18, 2
Density7.13 g.cm-3 at 20°C
Atomic number30
Atomic Mass65.39 g.mol -1
Electronegativity according to Pauling1.65

Occurrence

Zinc is quite abundant element and is ranked as the 24th most abundant element in the Earth’s crust (around 75ppm). It is found in soil and sea water.  Mostly, it is present in the form of ores and minerals of copper and lead. The most common ore of zinc is sphalerite (zinc sulfide) and contains around 60% of zinc [2]. It is also found in other minerals such as hemimorphite (zinc silicate), smithsonite (zinc carbonate) and wurtzite (zinc sulfide). Australia, USA and Iran and Canada are the largest producers of zinc in the world.

Physical Characteristics

Zinc is a whitish-blue shiny metal. It is hard and brittle at standard temperature. And gives a specific sound when bent like tin. It becomes malleable when temperature is increased from 100C to 150 [3]. The boiling and melting points of zinc are lower as compared to other members of the d-block elements and are 907°C and 419.5°C, respectively. Zinc is a good electrical conductor and have diamagnetic properties.

Chemical Characteristics

Zinc is a reactive metal. The surface of zinc metal is quickly tarnished as it reacts with the carbon dioxide present in the air and forms a layer of zinc carbonate. It is a strong reducing agent [4]. Zinc readily reacts with non-metals, acids and alkalis. In the presence of concentrated sulfuric acid, the upper passivating layer of zinc carbonate is dissolved and lead to the release of hydrogen gas. The most common oxidation state of zinc is +2. It burns with a greenish blue flame and forms zinc oxide.

Significance and Uses

  • Zinc is widely used for plating that imparts corrosion resistant upper layer to various metals.
  • Zinc is used to make various alloys, such as brass.
  • Zinc is used in the manufacturing of batteries.
  • Zinc compounds, such as zinc gluconate and zinc carbonate are used as dietary supplements.
  • Zinc is used as the active compound in various antidandruff shampoos.
  • Zinc sulfide is used to make luminescent dyes and paints.

Health Hazards

Zinc is considered as a biologically important element. It is required for the proper growth and development of human fetus. Deficiency of zinc in children lead to delayed or retarded growth. About two billion people in the world suffer from zinc deficiency that lead to various disorders and ailments [5].

Isotopes of Zinc

There are five stable isotopes of zinc, these include zinc-64, zinc-66, zinc-67, zinc68 and zinc-70. The most abundant natural isotope is zinc-64, which has an abundance of 48.63% [6]. It has a half life of around 4.3×1018 year, which is so high that it is considered almost stable. There are several dozen radioactive isotopes of zinc.

REFERENCES

[1]. Kharakwal, J. S. & Gurjar, L. K. (December 1, 2006). “Zinc and Brass in Archaeological Perspective”. Ancient Asia. 1: 139–159. doi:10.5334/aa.06112. Archived from the original on December 3, 2013.

[2]. Lehto 1968, p. 822

[3]. Heiserman 1992, p. 123

[4]. CRC 2006, pp. 8–29

[5]. Prasad, A. S. (2003). “Zinc deficiency : Has been known of for 40 years but ignored by global health organisations”. British Medical Journal. 326 (7386): 409–10. doi:10.1136/bmj.326.7386.409. PMC 1125304. PMID 12595353.

[6]. NNDC contributors (2008). Alejandro A. Sonzogni (Database Manager), ed. “Chart of Nuclides”. Upton (NY): National Nuclear Data Center, Brookhaven National Laboratory. Archived from the original on May 22, 2008. Retrieved September 13, 2008