Neodymium

Neodymium was discovered in 1885. It is a highly reactive metal and gets tarnished in air. It is main component of didymium glass. NIB magnets (Neodymium, iron and boron) are used in computers, cell phone, motors and audio system.

History and Discovery

Neodymium was discovered in 1885 by Baron Carl Auer Von Welsbach in Vienna. He separated neodymium and praseodymium from didymium through fractional crystallization of the double ammonium nitrate tetrahydrates from nitric acid, also following the separation through spectroscopic analysis. The word neodymium has been derived from Greek word neos meaning new and didymos meaning twin [1]. In 1950s, commercial purification of neodymium was carried out through ion-exchange method by Lindsay Chemical Division.  In 1927, commercially neodymium was used as glass dyes. Neodymium glasses were made in 1930s which has a reddish or orange tinge due to the traces of praseodymium but now a days the glass color is purple due to fractional crystallization technology.

Neodymium

Periodic Table ClassificationGroup n/a
Period 6
State at 20CSolid
ColorSilvery white
Electron Configuration[Xe] 4f4 6s2
Electron Number60
Proton Number60
Electron Shell2, 8, 18, 22, 8, 2
Density7.01 g.cm-3 at 20°C
Atomic number60
Atomic Mass144.24 g.mol -1
Electronegativity according to Pauling1.14

Occurrence

Neodymium is an abundant metal and is present in about 3.8 mg/kg in the Earth crust. It is twice more abundant than lead in the igneous rocks of the earth crust and half as plentiful as copper [2]. Neodymium is not found in free or elemental form in nature. It is usually found in the form of ores like monazite and bastnaesite. It belongs to the rare earth metal group, but it is not rare in nature. The largest deposits of neodymium in the world are present in China, United States, Brazil, India, Sri Lanka and Australia.

Physical Characteristics

Neodymium is a bright silver metal with a luster. It exists in two allotropic forms: α and β. Neodymium has chemical symbol is Nd with atomic number 60. Its atomic weight is 144.24. Its density at room temperature is about 7.01 g/cm3. Its melting point is 1024OC. Boiling point of neodymium is about 3074OC.

Chemical Characteristics

Neodymium is a very reactive lanthanide rare earth metal. Neodymium metal tarnishes in the air. It burns readily to form neodymium (III) oxide at about 150OC. It is quite electropositive element and reacts slowly with cold water and quickly with hot water to form neodymium (III) hydroxide. Neodymium reacts vigorously with all halogens. Neodymium dissolves in dilute sulfuric acid and form solution that contain the lilac Nd (III) ions. Neodymium mostly exist in +3 oxidation state and it forms compounds including halides, oxides, sulfides, nitrides, hydroxide phosphide and sulfate. Neodymium compounds have colors which depend upon the type of lighting. Most of its salts are pale purple in color.

Significance and Uses

  • Neodymium-doped glasses are used in laser that emit infrared with wavelength between 1047 and 1062 nanometer.
  • Neodymium is also a component in alloys to make high strength magnets.
  • Neodymium magnets are used in microphone, head phone, loudspeakers and computer hard disks.
  • It is also used with other substrate crystals in Nd-YAG laser (Neodymium: yttrium aluminum garnet). This laser emits infrared at a wavelength of 1064 nanometer, used in solid-state laser.
  • Neodymium acts in similar way as Ca2+ in promoting the growth of plants.
  • Neodymium is used to determine the age relationships of rocks and meteorites.
  • Neodymium is used in the electric motors of hybrid automobiles.
  • It is also used to color the glass in different shades ranging from pure violet with the help of wine red and warm gray.
  • Neodymium is also used to remove green color from glass due to iron contaminants.
  • Neodymium is component of didymium that is used to make welder’s and glass blower’s goggles.
  • Neodymium glass is used in incandescent light bulbs.
  • Its salts are used as colorant for enamels.
  • Neodymium is used in the manufacturing of house hold equipments like color television, lamp, energy saving lamps and glasses.

Health Hazards

Neodymium has no biological role. The salts and neodymium dust is very irritating for eyes. Compounds of neodymium are toxic if they are soluble in water but non-toxic if they are insoluble. It is dangerous in working areas and considered as a work hazard as it can cause lung embolisms due to long term exposure [3].

Isotopes of Neodymium

Neodymium has five stable isotopes: 142Nd, 143Nd, 145Nd, 146Nd and 148Nd. Neodymium-142 is the most abundant isotope. It has two radioactive isotopes: 144Nd and 150Nd. 144Nd have half-life of 2.29×1015 years and 150Nd has half -life of 7×1018 years approximately.

REFERENCES

[1]. John Emsley (2003). Nature’s building blocks: an A–Z guide to the elements. Oxford University Press. pp. 268–270.

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

[3]. https://www.lenntech.com/periodic/elements/nd.htm

Germanium

Germanium was discovered in 1886. Germanium belongs to the carbon family and used in electronic devices as semiconductors.

History and Discovery

Dmitri Mendeleev predicted the existence of germanium in 1869. He named it eka-silicon and proposed some of the physical characteristics of the then unknown metal. Later, Clemens Winker discovered germanium from a rare mineral known as argyrodite, in 1886 [1]. He successfully prepared various compounds with chlorides, fluorides, and dioxide, and found that the physical characteristics of germanium were almost the same as were predicted by Mendeleev. Germanium became a significant element in 1945, when it was used semiconductor in electronic, mostly diodes. Its symbol is Ge. In 1948, germanium transistors were developed and opened new horizons in electronic industry. In 2000, around 80% of the world’s germanium production was used in the manufacturing of optical fibers communication networks and laser infrared night vision systems. The element was named germanium by Winkler that was derived from Germania (Germany), his homeland.

Germanium

Periodic Table ClassificationGroup 14
Period 4
State at 20CSolid
ColorGrayish-white
Electron Configuration[Ar] 3d10 4s2 4p2
Electron Number32
Proton Number32
Electron Shell2, 8, 18, 4
Density5.32 g.cm-3 at 20°C
Atomic number32
Atomic Mass72.64 g.mol -1
Electronegativity according to Pauling2.01

Occurrence

Germanium is an abundant element and is ranked as the 15th most abundant element in the earth’s crust. Germanium is not present in its pure elemental form and mostly extracted from its zinc ore (sphalerite) and from ores of copper, silver and lead. Extremely pure germanium crystals are obtained through a technique termed as zone refining which produces semiconductor grade germanium with an impurity of 1 in 1010. This germanium semiconductors are considered as one of the purest substance that are ever made. The largest producers of germanium in the world are China, USA and Russia.

Physical Characteristics

Germanium is a whitish grey lustrous metalloid (semi-metallic) [2]. It is hard and brittle in nature.  there are two allotropic forms of germanium, alpha and beta germanium. Alpha germanium is hard lustrous metal with a cubic crystal structure that resembles diamond. The beta germanium is silvery and soft metal [3].

Chemical Characteristics

Pure germanium reacts slowly with air at higher temperature, around 250°C and forms germanium oxide. It does not react with dilute acids but is easily dissolved in hot concentrated acids. Germanium reacts with halogens to form compounds like germanium tetrachloride and germanium tetraiodide. Germanium reacts with oxygen and form two oxides, germanium monoxide and germanium dioxide. It also reacts with sulfur, selenide and telluride to form disulfide GeS2, germanium diselenide GeSe2 and germanium telluride (GeTe) [4].

Significance and Uses

  • Germanium has been widely used for manufacturing of semiconductors.
  • It is used to make light emitting diodes (LEDs).
  • Germanium is used in making optical fibers and solar cells.
  • Germanium is used to make nanowires, which is one of its most recent usages.
  • Various organic compounds of germanium have diverse applications, such as tetraethyl germane is widely used in the study of organometallic chemistry.

Health Hazards

Germanium and most of its natural compound are non-toxic as they are insoluble in water. certain synthetic salts of germanium are toxic and have shown to have damaging effects on kidneys. Germanium is not a biologically significant element and have no role in the bodies of plants or animals.

Isotopes of Germanium

Germanium have five naturally occurring isotopes: germanium-70, germanium-72, germanium-73, germanium-74 and germanium-76. Germanium-74 is the most abundant isotope while germanium-76 is the least abundant. Germanium-76 is also radioactive, however, it is not highly radioactive as it has a half-life of 1.78 ×1021 years. There are also twenty-seven artificially prepared radioisotopes of germanium. Their mass number range from 58 to 89.

REFERENCES

[1]. Winkler, Clemens (1887). “Germanium, Ge, a New Nonmetal Element”. Berichte der Deutschen Chemischen Gesellschaft (in German). 19 (1): 210–211. doi:10.1002/cber.18860190156. Archived from the original on December 7, 2008.

[2]. Emsley, John (2001). Nature’s Building Blocks. Oxford: Oxford University Press. pp. 506–510. ISBN 978-0-19-850341-5.

[3]. Holleman, A. F.; Wiberg, E.; Wiberg, N. (2007). Lehrbuch der Anorganischen Chemie (102nd ed.). de Gruyter. ISBN 978-3-11-017770-1. OCLC 145623740

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

Flerovium

Flerovium is a synthetic element that was discovered in 1999. It is an extremely heavy and highly radioactive element.

History and Discovery

Dmitri Mendeleev predicted the presence of element 114 and it was named eka-lead. Flerovium was synthesized team of physicist led by Yuri Oganessian working in Joint Institute of Nuclear Research (JINR), Dubna, Russia in 1999 [1]. They bombarded plutonium-244 with calcium-48 nuclei and produced a single atom of flerovium-289. Long before its discovery, isotopes of flerovium were predicted to lie in the island of stability (elements from copernicium (112) to oganesson (118)) and with magic number of protons and neutrons, 114 protons and 184 neutrons, after helium-4, oxygen-16, calcium-48 and lead-208 which are the other members of the doubly magic atomic numbers group. It has proposed that the nuclei flerovium would be oblate and deformed which would attribute to its resistance to spontaneous fission. And in 1972, it was proposed that flerovium-298 would have a half-life of around a year. The name of the newly discovered element was confirmed as flerovium by International Union of Pure and Applied Chemistry in 2012, after the Flerov Laboratory of Nuclear Reactions of JINR, Russia. The laboratory was named after a Russian physicist, Georgy Flyorov [1]. Its symbol is Fl.

Flerovium

Periodic Table ClassificationGroup 14
Period 7
State at 20CGas (predicted)
Colorn/a
Electron Configuration[Rn] 5f14 6d10 7s2 7p2 (predicted)
Electron Number114
Proton Number114
Electron Shell2, 8, 18, 32, 32, 18, 4 (predicted)
Density14.00 g.cm-3 at 20°C (predicted)
Atomic number114
Atomic Mass289.00 g.mol -1 (most stable isotope)
Electronegativity according to Paulingn/a

Occurrence

Flerovium is an artificial element and does not exist in nature. Around 90 atoms of flerovium have been made so far.

Physical Characteristics

Flerovium is predicted to be a gas at room temperature. Flerovium is expected to be present in the carbon group (Group 14) along with silicon, tin, germanium and lead. Its boiling point is lower than other members of group 14, around -60 °C. and the melting points is also low, i.e. 70 °C. Flerovium crystallize in a cubical structure that is face-centered.  In solid form, flerovium is an extremely heavy element and would have a density of 14 or 22 g/cm3 [2].

Chemical Characteristics

The outermost shell of flerovium is filled and thus makes it a noble (inert) metal. It has no electron affinity. The first ionization energy of flerovium is expected to be the highest among its group. And its electronegativity would also be higher than lead.  It is also the heaviest member of the periodic table. Other members of carbon group form stable compounds with +4 oxidation state but flerovium is expected to be most stable at an oxidation state of +2 and 0 [2]. Due to its short life span and relativistic effects, the chemical and physical characteristics of this element are yet to be determined.

Significance and Uses

  • Flerovium does not have any industrial applications is used for research purposes.

Health Hazards

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

Isotopes of Flerovium

There are seven isotopes of flerovium, however some of them are still unconfirmed. They are unstable and unnatural. The most stable isotope of flerovium is flerovium-289 with a half-life of 2.6 seconds [3].

REFERENCES

[1]. Welsh, J. (2 December 2011). “Two Elements Named: Livermorium and Flerovium”. LiveScience. Retrieved 2 December 2011

 [2]. Hoffman, Darleane C.; Lee, Diana M.; Pershina, Valeria (2006). “Transactinides and the future elements”. In Morss; Edelstein, Norman M.; Fuger, Jean. The Chemistry of the Actinide and Transactinide Elements (3rd ed.). Dordrecht, The Netherlands: Springer Science+Business Media. ISBN 1-4020-3555-1.

[3]. Oganessian, Yu. Ts.; et al. (2000). “Synthesis of superheavy nuclei in the 48Ca + 244Pu reaction: 288114″ (PDF). Physical Review C. 62 (4): 041604. Bibcode:2000PhRvC..62d1604O. doi:10.1103/PhysRevC.62.041604.

 

Protactinium

Predicted by Dimitri Mandeleev in 1871, protactinium was first identified in 1913 by Kasimir Fajans and Oswald Helmuth Gohring. It is highly radioactive, scarce and toxic element and has no uses outside scientific research. 

History and Discovery

Dmitri Mendeleev predicted the existence of an element between uranium and thorium, in 1871 [1]. William Crookes, in 1900, successfully isolated protactinium as a radioactive material but instead of classifying it as a new element he called it uranium-X. In 1913, Kasimir Fajans and Oswald Helmuth Gohring identified an isotope of protactinium and named it brevium. The isotope he studied was protactinium-234 and had short half-life of 6.7 hours. Four years later a more stable isotope of protactinium was discovered by Otto Hahn and Lise Meitner. They chose a name Proto-actinium for this element indicating actinium (Ac-89) is a decay product of this new element. The international union for pure and applied chemistry finalized the name protactinium for this new element, in 1949. They credited Otto Hahn and Meitner for its discovery. Aristid von Grosse first isolated protactinium in elemental form in 1934.

Protactinium

Periodic Table ClassificationGroup n/a
Period 7
State at 20CSolid
ColorBright, silvery metallic luster
Electron Configuration[Rn] 5f2 6d1 7s2
Electron Number91
Proton Number91
Electron Shell2, 8, 18, 32, 20, 9, 2
Density15.40 g.cm-3 at 20°C
Atomic number91
Atomic Mass231.04 g.mol -1
Electronegativity according to Pauling1.50

Occurrence

Protactinium is one of the rarest naturally occurring elements. Two isotopes of protactinium Pa-231 and Pa-234 occur naturally. Pa-231 is the most abundant naturally occurring isotope. Protactinium is found in uraninite ore (uraninite is a uranium rich ore). Both isotopes of protactinium are a decay product of different isotopes of uranium. It is homogeneously dispersed in water and most natural minerals. Protactinium is also produced in nuclear reactors. In 1961, 127g of 99.9% pure protactinium (Pa-231) was produced by United Kingdom atomic energy authority which remained a substantial supply for the whole world for scientific studies on protactinium. These 127 grams of protactinium was obtained by a 12 stage procedure which processed 60 tonnes of waste material, costing 0.5 million USD.

Physical Characteristics

Protactinium is dense silver grey metal. It is highly radioactive. It belongs to actinide series and is placed between thorium and uranium in the periodic table. The melting point of protactinium is lower than thorium but higher than that of uranium. Protactinium crystallizes and forms a rigid tetragonal structure at room temperature. It is paramagnetic and becomes superconductor at low temperatures. The atomic number of protactinium is 91 and is represented by symbol Pa.

Chemical Characteristics

Protactinium has bright metallic lustre and tarnishes slowly in air. Protactinium reacts with air, oxygen, acids and water vapours. It does not react with alkalis. Protactinium oxides easily convert to hydroxides and form various salts.

Significance and Uses

Protactinium is rare and expensive mineral and due to its high radioactivity and toxicity it is not used much but for scientific research. It has very few uses, such as:

  • It can be used as a tracer in paleoceanography and geology. Protactinium relative concentration in water and minerals is used for radiometric dating of sediments [2].
  • Protactinium-233 with half-life of 27 days is used as neutron poison because of its high cross section for neutron capture.

Health Hazards

Protactinium is highly radioactive element. Isotopes of protactinium decay via both alpha and beta decay. Protactinium-231 contributes to the radio-toxicity of the spent fuel in nuclear reactors. It is also a toxic element. Protactinium is naturally present in most natural products and minerals in small amount [3]. When inhaled through air or ingested with food or water a small amount of protactinium is absorbed by the body which goes to kidney, liver and bones. Due to its radioactivity it promotes cancer in human body.

Isotopes of Protactinium

Protactinium has twenty-nine known radioisotopes. Apart from Pa-231 all the isotopes are unstable with half-life ranging from few days to less than few seconds. Protactinium-231 is known to have the longest half-life amongst the isotopes of protactinium. It is a decay product of uranium and has a half-life of 32,760 years. Protactinium isotopes form actinium and uranium isotopes as decay products [4].

REFERENCES

[1]. Emsley, John (2003) [2001] “Protactinium”. Nature’s Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 347–349.

[2]. Hammond, C. R. The Elements, in Handbook of Chemistry and Physics (81st ed.). CRC press.  ISBN 0-8493-0485-7.

 [3]. Articles “Protactinium” and “Protactinium-231 – thorium-230 dating” in Encyclopaedia Britannica, 15th edition, 1995, p. 737.

[4]. Audi, G.; Bersillon, O.; Blachot, J.; Wapstra, A. H. (2003). “The NUBASE evaluation of nuclear and decay properties” , Nuclear Physics A. 729:3.  doi:10.1016/j.nuclphysa.2003.11.001.

Neptunium

Neptunium, successor of uranium, was discovered in 1940 by Edwin McMillan and Philip H. Abelson while they were trying to isolate fission products. It is used in plutonium production which in turn has uses in military and is also used as fuel to provide electricity to space craft.

History and Discovery

Dmitri Mendeleev published the periodic table with the then undiscovered elements indicated by dashes. At that time uranium was the last known element with atomic number 92. In Mendeleev version of periodic table and the later version of periodic table, published in 1913 by Kasimir Fajans, an empty space after uranium indicated a possibility of a new element. After the discovery of induced radioactivity in 1933, Enrico Fermi with his group members bombarded all the known elements with neutron and analysed that heavier elements would disperse the energy by emitting gamma ray.  Emission of gamma ray results in beta decay which in turn increases the atomic number of the resulting isotope. Neutron bombardment on uranium resulted in an isotope with atomic number 93. Fermi published his finding of element 93 but several objections were raised against his claims. Fermi was not able to prove his claim as he was unable to isolate the newly created element.  In 1934, Odolen Koblic claimed to have extracted a small amount of element 93 from pitchblende but when analysed it turned out to be a mixture of two other metals. In 1938, Yvette Cauchois and Horia Hulubei claimed to have discovered element 93 via spectroscopy but the claim was denied since it was then believed that element 93, if it exist at all, would not exist naturally. In 1939, Edwin McMillan decided to bombard uranium with cyclotron to separate fission products. McMillan observed a new half-life of 2.3 days in his sample results. This element had chemistry like no other known element and was similar to uranium. McMillan, with the help of Philip H. Abelson, finally isolated the source and published their results in 1940 [1].  Element 93 was named after the planet Neptune following the naming of Uranium which was named after the planet Uranus, Neptune being the next planet after Uranus in our solar system.

Neptunium

Periodic Table ClassificationGroup n/a
Period 7
State at 20CSolid
ColorSilvery metallic
Electron Configuration[Rn] 5f4 6d1 7s2
Electron Number93
Proton Number93
Electron Shell2, 8, 18, 32, 22, 9, 2
Density20.20 g.cm-3 at 20°C
Atomic number93
Atomic Mass237.00 g.mol -1
Electronegativity according to Pauling1.36

Occurrence

Neptunium is only present on earth as an intermediate decay product of other elements. Because of its relatively short half-life any primordial element would have decayed by now. Neptunium isotopes are found naturally in uranium ores which are formed as decay products from the transmutation reactions. Majority of currently existing neptunium on earth is produced in artificial nuclear reactions.

Physical Characteristics

Neptunium is a radioactive metal. It has hard silvery metallic appearance and tarnishes with the exposure to air. It occurs in three allotropic forms. Neptunium has 7 electrons in its valence shell. Neptunium is the first transuranic element and belongs to actinide series. All transuranic elements are unstable and decay into other elements radioactively. It is represented by the symbol Np and has atomic number 93.

Chemical Characteristics

Neptunium is a very reactive metal. It is pyrophoric at room temperature in powdered form. Neptunium is the heaviest actinide that has the ability to lose valence electrons in stable compound. The most stable state in solid form is +4 while +5 valence state is preferred in solution form [2].

Significance and Uses

Neptunium has very few uses, such as:

  • Neptunium is used to for the detection of high energy neutrons.
  • It is used as precursor in production of plutonium, which in turn is used as fuel in thermal generators to provide electricity to spacecraft and has applications in military [3].
  • Neptunium is fissionable and theoretically can be used to make nuclear explosives or can be used as fuel in fast neutron reactors.

Health Hazards

There are 20 radioisotopes of neptunium that have been characterized. The atomic weight of isotopes of neptunium range from 225.0339 u to 244.068 u [4]. Np-237 has the longest half-life of about 2.14 million years. Np-236 and Np-235 also have half-life of a year or more than a year. The rest of the known isotopes have half-life ranging from few days to less than few minutes.

Isotopes of Neptunium

There are seven isotopes of palladium, out of which only one is unstable. Palladium-107 is the most stable radio-iso tope and it have a half life of 6.5 million years. There are eighteen radioactive isotopes of palladium that are artificially produced and have atomic masses that range from 90.94 to 122.93. most of these isotopes have half-life of less than half an hour, expect palladium-101, palladium-109, and palladium-112.

REFERENCES

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

[2]. Hindman J. C. 1968, “Neptunium”, in C. A. Hampel (ed.), The encyclopedia of the chemical elements, Reinhold, New York, pp. 434.

[3]. Lemire, R. J. et al.,Chemical Thermodynamics of Neptunium and Plutonium, Elsevier, Amsterdam, 2001

[4]. Lee, J.; Mardon, P.; Pearce, J.; Hall, R. (1959). “Some physical properties of neptunium metal II: A study of the allotropic transformations in neptunium”. Journal of Physics and Chemistry of Solids. 11 (3–4): 177–181. doi:10.1016/0022-3697(59)90211-2.

Gallium

Gallium was discovered in 1971. It is a non-radioactive, non-toxic metal and have been widely used a semiconductors in various electronic devices.

History and Discovery

Dmitri Mendeleev predicted the existence of gallium in 1971 and named it eka-aluminum. He also predicted some properties of gallium which were later confirmed when the element was discovered [1]. In 1875, Paul Emile Lecoq de Boisbaudran a French chemist discovered gallium through spectroscopically analysis of sphalerite and observing violet lines in spectrum [2]. The element was named gallium after the Latin word Gallia that was the native land of French discoverer.

Gallium

Periodic Table ClassificationGroup 13
Period 4
State at 20CSolid
ColorSilvery blue
Electron Configuration[Ar] 3d10 4s2 4p1
Electron Number31
Proton Number31
Electron Shell2, 8, 18, 3
Density5.91 g.cm-3 at 20°C
Atomic number31
Atomic Mass69.72 g.mol -1
Electronegativity according to Pauling1.81

Occurrence

Gallium is a not an abundant element and is present in about 16.9ppm in the Earth’s crust. It Gallium does not exist in elemental form in nature. Gallium is also found in scarce amount in minerals such as gallite [3]. It is mainly present in the ores of aluminum (bauxite) and zinc. Commercially, gallium is produced through smelting of various ores, including bauxite and some ores of zinc sulfide.

Physical Characteristics

Gallium is bluish silver metal. It is soft at STP (standard temperature and pressure), while at low temperature, it acquires a brittle solid state. Gallium also exists in liquid form at temperature more than 29.76 °C. When gallium is held in the hand, it melts and is re-solidified when placed back from hand. Thus, the melting point of gallium is 29.76 °C. Melting point of gallium is used as reference point of temperature set by the International Bureau of Weights and Measures (BIPM). The boiling point of gallium is also unique, as it is the only element that has the greatest difference or ratio between boiling and melting point as the boiling point is 2399°C which is about eight times higher than its melting point. At high temperature, gallium has a low vapor pressure. Gallium belongs to the boron group (Group 13) along with aluminum, indium, thallium and nihonium. When gallium is frozen, it expands as it solidifies, so care should be taken while storing gallium at low temperatures. In solid form, it forms an orthorhombic crystal shape with eight atoms that form a unit cell. Gallium also tends to supercool below its freezing point into gallium nanoparticles. Gallium is a high-density liquid.

Chemical Characteristics

Gallium is not a very reactive metal. the most common oxidation state of gallium is +3. Gallium dissolves in strong acids and alkalis. Gallium dissolves in water and aqueous solutions of gallium have hydrated gallium ions. Gallium hydroxide is an amphoteric compound and it can dissolve in alkalis and produce salts of gallium. Gallium reacts with ammonia to form gallium nitride at high temperature. It can also react with antimony, arsenic, and phosphorus to form binary compounds. Gallium reacts with halides, including fluorine, chlorine and iodine to form stable compounds. Gallium does not react with water and air at room temperature due to the formation of a surface layer of oxide. However, at higher temperatures, gallium reacts with air to form gallium oxide.

Significance and Uses

  • Gallium is widely used in electronic industry. It is part of infrared circuits, microwaves and efficient switching circuits.
  • Gallium is used as a non-toxic alternative to mercury in thermometers.
  • Gallium is widely used to make alloys that have low melting points.
  • Gallium is used in the manufacturing of semiconductors. Gallium nitride semiconductors produce violet and blue light-emitting diodes and lasers.
  • Gallium is used to make high quality jewellery, for instance gadolinium gallium stones that have brilliant looks that resembles diamonds or gemstones.

Health Hazards

Gallium is not a biologically significant element. It is a non-toxic element. Gallium is also a non-radioactive metal.

Isotopes of Gallium

There are thirty-one isotopes of gallium and their atomic mass range from 56 to 86. The stable isotopes are only two, gallium-69 and gallium-71. The most abundant isotope is gallium-69 and makes around 71% in abundance. All other isotopes of gallium are unstable and radioactive. Gallium-67 and gallium-68 are the commercially important radioactive isotopes of gallium. The lighter isotopes (below gallium-69) decay through emission of positron termed as beta plus emission while the heavier isotopes undergo electron emission that is termed as beta minus decay.

REFERENCES

[1]. Ball, Philip (2002). The Ingredients: A Guided Tour of the Elements. Oxford University Press. p. 105. ISBN 978-0-19-284100-1

[2]. de Boisbaudran, Lecoq (1835–1965). “Caractères chimiques et spectroscopiques d’un nouveau métal, le gallium, découvert dans une blende de la mine de Pierrefitte, vallée d’Argelès (Pyrénées)”. Comptes Rendus. 81: 493. Retrieved 2008-09-23.

[3]. “The distribution of gallium, germanium and indium in conventional and non-conventional resources – Implications for global availability (PDF Download Available)”. ResearchGate. doi:10.13140/rg.2.2.20956.18564. Retrieved 2017-06-02

Ruthenium

Ruthenium was discovered in 1884 and is member of the platinum family. It is a rare metal and widely in electrical industry and in making alloys.

History and Discovery

Like other six members of the platinum-group, ruthenium has been known since pre-Columbian American time. It was used by European chemist in mid-16th to 18th century.  Ruthenium was discovered by Karl Ernst Claus, a Russian scientist, in 1884 [1]. The was a member of Russian Academy of Science and discovered the element at Kazan State University in Russia. The element was named after the Latin name of Claus’s homeland, Ruthenia.

Ruthenium

Periodic Table ClassificationGroup 8
Period 5
State at 20CSolid
ColorSilvery white metallic
Electron Configuration[Kr] 4d7 5s1
Electron Number44
Proton Number44
Electron Shell2, 8, 18, 15, 1
Density12.37 g.cm-3 at 20°C
Atomic number44
Atomic Mass101.07 g.mol -1
Electronegativity according to Pauling2.20

Occurrence

Ruthenium is a rare metal. It is ranked as the seventy-fourth most abundant element and is present around 100 parts per trillion in the Earth’s crust. The elemental or free form of ruthenium is very rare. Mostly, ruthenium is present in minerals and ores, especially with platinum, in minute amounts. Each year around 12 tons of ruthenium are extracted and the estimated reserves of ruthenium in the world are about 5000 tons [2]. Commercially, ruthenium is extracted as a by-product during the smelting process of ores of platinum, copper and nickel.  The largest reservoirs of ruthenium are in South Africa, Canada and USA [3].

Physical Characteristics

Ruthenium is a whitish silver transition metal. It belongs to the platinum group of metals and in the 8th group of periodic table. Unlike the other members of the group that have two electrons in their outmost shell, ruthenium has only one valence electron. Ruthenium is resistant to tarnishing by air until exposed to high temperatures.

Chemical Characteristics

Ruthenium is a fairly inert metal as it does not react with most of the elements of the periodic table. It is resistant to attack by acids; however, it dissolves in mixture of alkalis. Ruthenium also reacts with halogens at high temperatures. The most common oxidation states of ruthenium are +2, +3 and +4, however, the compounds of ruthenium can exist in a long range of oxidation state, -2 and from 0 to +8. Ruthenium is oxidized to form ruthenium oxide and is further oxidized to form ruthenium tetroxide, which is a volatile and strong oxidizing agent. Oxide of ruthenium resemble osmium tetroxide. With chalcogens, ruthenium forms dichalcogenides, which have properties of diamagnetic semiconductors. Ruthenium reacts with halogens, and forms ruthenium halides, such as ruthenium reacts with fluorine to form hexafluoride, pentafluoride and tetrafluoride and form ruthenium trichloride with chlorine.

Significance and Uses

  • Ruthenium is most widely used in electronic industry. Ruthenium dioxide is used to make thick-film chip resistors.
  • Ruthenium is widely used to make alloys with other metals. For example, it used to provide strength and increase the corrosion resistance of palladium and platinum.
  • Ruthenium tetroxide is used as staining and fixing agent in electron microscopy of organic tissues and compounds.
  • Ruthenium is widely used as catalyst in various industrial applications.

Health Hazards

Ruthenium is a non-toxic metal. It has not known biological role. However, various compounds of ruthenium, including ruthenium oxide, are quite toxic, and some are considered as poisonous and carcinogenic. Ruhtneium-106, is a radioactive isotope and has a long half-life. It was part of the nuclear testing program that started in second World War (1945) and ended in 1980s. Due to its stable nature, ruthenium can pose danger to human health for long time period, for decades and even centuries [4].

Isotopes of Ruthenium

There are seven isotopes in naturally occurring ruthenium. They are also the stable isotopes. There are thirty-four unstable or radioactive isotopes of ruthenium, among them the most stable isotope is ruthenium-106 that has a half-life of 373.59 days. Most of the isotopes undergo decay through electron capture and beta emission.

REFERENCES

[1]. Weeks, Mary Elvira (1932). “The discovery of the elements. VIII. The platinum metals”. Journal of Chemical Education. 9 (6): 1017. Bibcode:1932JChEd…9.1017W. doi:10.1021/ed009p1017.

[2]. Emsley, J. (2003). “Ruthenium”. Nature’s Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 368–370. ISBN 978-0-19-850340-8.

[3]. “Platinum–Group Metals” (PDF). U.S. Geological Survey, Mineral Commodity Summaries. January 2007. Retrieved 2008-09-09

[4]. https://www.lenntech.com/periodic/elements/ru.htm

 

Praseodymium

Praseodymium was discovered in 1841. It is strong paramagnetic in nature. It is widely used as an alloying agent with magnesium to make high strength metal. It is found in house equipment like color television, lamps and glasses.

History and Discovery

Praseodymium was discovered in 1841 by Carl Gustav Mosander through extraction from a rare earth oxide called ‘didymium’ which was obtained from the residues called ‘lanthana’. In 1869, didymium got the symbol ‘Di’ in Mendeleev’s first periodic table. In 1885 Baron Carl Auer von Welsbach separated ‘didymium’ into two different elements, one named praseodymium and other called neodymium. The word praseodymium has been derived from Greek prasinos meaning green, and didymos means twins as it produces green color compounds. Pure metallic form of praseodymium was first synthesized in 1931 [1].

Praseodymium

Periodic Table ClassificationGroup n/a
Period 6
State at 20CSolid
ColorGrayish white
Electron Configuration[Xe] 4f3 6s2
Electron Number59
Proton Number59
Electron Shell2, 8, 18, 21, 8, 2
Density6.77 g.cm-3 at 20°C
Atomic number59
Atomic Mass140.91 g.mol -1
Electronegativity according to Pauling1.13

Occurrence

Praseodymium is not a rare element. It is present in about 9.1mg/kg of the Earth’s crust. It is fourth most abundant element of the lanthanide’s series in the periodic table. It is present in phosphate, silicate and carbonate minerals like monazite (reddish brown phosphate mineral) and bastnaesite (family of three carbonate-fluoride mineral). Commercially, praseodymium is separated from lanthanides through ion exchange chromatography.

Physical Characteristics

Praseodymium is the third element of the lanthanide series. It is moderately soft and silver white metal. Its hardness is comparable with the silver and it is ductile in nature [2]. It exist in two allotropic forms α and β. Praseodymium has a strong paramagnetic nature (materials are attracted by external magnetic field).  It generally exist in trivalent state. Praseodymium chemical symbol is Pr. Its atomic number is 59. Atomic weight is 140.9. Melting point of Pr is about 931OC. The boiling point of Praseodymium is very high, i.e. 3520OC. Its density at room temperature is about 6.77 g/cm3.

Chemical Characteristics

Praseodymium is tarnished in air slowly and forming oxide layer like iron rust. Its centimeter sized sample of metal corrodes in a year. At 150OC it forms praseodymium (III, IV) oxide, and it may be reduced to praseodymium (III) oxide in the presence of hydrogen gas. Praseodymium (IV) oxide is obtained by reaction with pure oxygen at 400OC and 282 bar. It is electropositive and react slowly with cold water and quite quickly react with hot water to form praseodymium (III) hydroxide. Praseodymium metal react with all halogens to form trihalides. It dissolves readily in dilute sulfuric acid and form solution which contain chartreuse (color between yellow and green) Pr3+ ions. Praseodymium dissolve in water and the solution contains yellow Pr4+ ions.

Significance and Uses

  • Praseodymium is used in glass coloration (yellow and green color).
  • It is widely used in making high-power magnets with great strength and durability.
  • It is also used as an alloying agent in magnesium to make high strength metals which is used in aircraft engines.
  • It is used to color ceramics yellow.
  • Its salts are used to color glasses and enamels.
  • Praseodymium is used in the core of high intensity carbon arc lights which are used in film industry and food lightning.
  • It is also used for making specialized yellow glass goggles for glass blowers and welders.
  • It is also main component of didymium glass, which protects welder’s eyes from intense light.
  • Praseodymium is used as a doping agent in fiber optics as signal amplifier.
  • It is used to make 5% misch metal (mixed metal containing cerium, lanthanum and neodymium) to make flint for lighters.
  • It is also used as an alloying additions to ferrous and non- ferrous alloys.

Health Hazards

Praseodymium has low to moderate toxicity. Its soluble salts are mildly toxic in nature, but insoluble salts are non- toxic. In industrial working areas, it can cause lungs embolisms due to long term exposure. When it is accumulated in the body, it is dangerous for liver.

Isotopes of Praseodymium

Praseodymium has only one natural and stable isotope which is 141Pr. It has thirty two isotopes having mass number 121 to 154, all have half-lives under a day except 143Pr, which has a half-life of 13.6 days. Others isotopes have half-lives ranging from 10 millisecond to 13.57 days [3].

REFERENCES

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

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

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

 

Livermorium

Livermorium is a synthetic element that was discovered in 2000. It is highly radioactive and unstable element.

History and Discovery

According to the Mendeleev’s nomenclature of undiscovered elements, livermorium was named as eka-polonium or element-116. Livermorium was synthesized by team of scientist working in Dubna at Joint Institute for Nuclear Research (JINR) and Lawrence Livermore National Laboratory, Russia in 2000 [1]. They bombarded curium-248 (element 96) with calcium-48 (element 20) nuclei and obtained one atom of livermorium-293. Then in 2005, eight more atoms of livermorium were produced by repeating the make experiment. And during the experiment, another isotope, livermorium-292 was also discovered [2]. The discovery of livermorium has been verified by RIKEN in 2016 and GSI in 2012. Its name was confirmed as Livermorium by International Union of Pure and Applied Chemistry in 2012, after the Livermore city in California. Its symbol is Lv and was selected in 2012 in a ceremony held in Moscow.

Livermorium

Periodic Table ClassificationGroup 16
Period 7
State at 20CSolid (predicted)
Colorn/a
Electron Configuration[Rn] 5f14 6d10 7s2 7p4 (predicted)
Electron Number116
Proton Number116
Electron Shell2, 8, 18, 32, 32, 18, 6 (predicted)
Density12.90 g.cm-3 at 20°C (predicted)
Atomic number116
Atomic Mass293.00 g.mol -1 (most stable isotope)
Electronegativity according to Paulingn/a

Occurrence

Livermorium is an artificial element and does not exist in nature. It is extremely radioactive and thus is very unstable and only created in laboratory.

Physical Characteristics

Livermorium has been predicted to be a solid under normal conditions and a post transition metal. Only minute amount of livermorium have been produced so far and it is not enough to carry out a statistically significant analysis of its physical and chemical characteristics. However, it is predicted to be heaviest among the other chalcogens, including selenium, oxygen, polonium, tellurium and sulfur. Its density is predicted to be 12.9 g/cm3). Livermorium is present on the bottom of its group, below polonium. The melting and boiling points of livermorium is expected to be according to the trend in the group, and it will have lower boiling point and a higher melting point than polonium [3]. There are two allotropic forms of livermorium, alpha and beta.

Chemical Characteristics

The chemical characteristics of livermorium is not well studied yet. However, it is expected to resemble in characteristics with other members of chalcogen group. Livermorium belongs to the group 16 and period 7th of the periodic table. And is categorized as a 7p-block transactinide element.  Due to the orbit-spin interaction and relativistic effect, it is presumed that livermorium may vary in certain properties from the lighter members of its group. The most common oxidation state of livermorium is +2, while other members of the group have +6 and +4. The hydride of livermorium, livermorane (LvH2) is the heaviest chalcogen hydride.

Significance and Uses

  • Livermorium is used for research purposes.

Health Hazards

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

Isotopes of Livermorium

There are four isotopes of livermorium, and their mass numbers range between 290 and 293. They are unstable and unnatural. Livermorium-293 has a half-life of 60 millisecond and is considered as the most stable isotope of livermorium. Livermorium isotopes mostly decays through emission of alpha particles.

REFERENCES

[1]. Oganessian, Yu. Ts.; Utyonkov; Lobanov; Abdullin; Polyakov; Shirokovsky; Tsyganov; Gulbekian; Bogomolov; Gikal; Mezentsev; Iliev; Subbotin; Sukhov; Ivanov; Buklanov; Subotic; Itkis; Moody; Wild; Stoyer; Stoyer; Lougheed; Laue; Karelin; Tatarinov (2000). “Observation of the decay of 292116″. Physical Review C. 63: 011301. Bibcode:2001PhRvC..63a1301O.

 [2]. Oganessian, Yu. Ts.; Utyonkov, V.; Lobanov, Yu.; Abdullin, F.; Polyakov, A.; Shirokovsky, I.; Tsyganov, Yu.; Gulbekian, G.; Bogomolov, S.; Gikal, B. N.; et al. (2004). “Measurements of cross sections and decay properties of the isotopes of elements 112, 114, and 116 produced in the fusion reactions 233,238U, 242Pu, and 248Cm+48Ca” (PDF). Physical Review C. 70 (6): 064609. Bibcode:2004PhRvC..70f4609O. doi:10.1103/PhysRevC.70.064609.

[3]. Bonchev, Danail; Kamenska, Verginia (1981). “Predicting the Properties of the 113–120 Transactinide Elements”. Journal of Physical Chemistry. American Chemical Society. 85 (9): 1177–1186. doi:10.1021/j150609a021

Palladium

Palladium was discovered in 1802 and is member of the platinum group. It is widely used as a catalytic converter in automobile fuel industry.

History and Discovery

Palladium was discovered by William Hyde Wollaston in 1802. The element was named after the Greek goddess Athena, who slew a giant (according to one myth) named Pallas and wore his skin as her gown. The name palladion was given to the statue of Athena that was believed to have talismanic properties to keep the kingdom (Troy) safe. Pallas was also the name of an asteroid that was discovered in 1802, so Wollaston gave the newly discovered element the name palladium [1].

Palladium

Periodic Table ClassificationGroup 10
Period 5
State at 20CSolid
ColorSilvery white
Electron Configuration[Kr] 4d10
Electron Number46
Proton Number46
Electron Shell2, 8, 18, 18
Density12.02 g.cm-3 at 20°C
Atomic number46
Atomic Mass106.42 g.mol -1
Electronegativity according to Pauling2.20

Occurrence

Palladium is a rare metal. the minerals and ores of palladium are not very common in the earth’s crust. Palladium is found in the alloy form with gold and other metals of the platinum group. Commercially, palladium is produced by nickel-copper deposits that are found in Ontario and Siberia. Palladium is also present in form of rare minerals such as polarite and cooperite. The largest producers of palladium include South Africa, Canada and Russia, where large natural deposits of palladium are present [2].

Physical Characteristics

Palladium is silvery white transition metal. It has a shiny appearance and is soft and ductile in nature. Palladium belong to the group of exquisite metals (including rhodium, platinum, osmium, ruthenium, osmium and iridium) termed as the platinum group metals (PGMs). Palladium is a heavy metal and has a density of 12.03 g/cm3. In liquid form, palladium has a density of 10.38, which makes it a dense liquid. But palladium is less dense in both states as compared to other metals of PGM group. Palladium also have the lowest melting point among the elements of its group.

Chemical Characteristics

Palladium is not a reactive element. It is dissolved in concentrated sulfuric and nitric acid, but the reaction is quite slow. N powdered form, palladium dissolves in hydrochloric acid. Palladium readily dissolved in aqua regia at room temperature [3]. At standard temperature, palladium does not react with oxygen and that is why is resistant to tarnishing in air. At higher temperature, around 800C, palladium reacts with oxygen and forms a tarnished layer of palladium oxide on the surface. The most common oxidation states of palladium are 0 and +2. Palladium compounds resembles compounds of platinum. It reacts with halogens to form halides, and the common halide is palladium (II) chloride. Palladium chloride dissolves in a mixture of nitric acid and acetic acid to form palladium acetate.

Significance and Uses

  • A large amount of palladium is used in the catalytic conversion of harmful gases from automobile exhaust, carbon monoxide and hydrocarbon into less harmful gasses, such as nitrogen, carbon dioxide and water.
  • Palladium is widely used to make fuel cells, where it reacts with hydrogen and oxygen to produce electricity, along with heat and water.
  • Palladium is used in electronic industry and in making watches.
  • Palladium is used for various medical and dentistry purposes.
  • Palladium is used in making ornaments and jewellery.

Health Hazards

Palladium is a moderately toxic metal. and high dose of this element can be very dangerous, for animals and plants. Some experiments reflect the carcinogenic nature of palladium. However, little evidence is present regarding the toxic effects of palladium on humans.  In mice, however, the median lethal dose of soluble palladium has been estimated to be an intravenous dose of 5mg/kg and an oral does 200mg/ kg [4].

Isotopes of Palladium

There are seven isotopes of palladium, out of which only one is unstable. Palladium-107 is the most stable radio-iso tope and it have a half life of 6.5 million years. There are eighteen radioactive isotopes of palladium that are artificially produced and have atomic masses that range from 90.94 to 122.93. most of these isotopes have half-life of less than half an hour, expect palladium-101, palladium-109, and palladium-112.

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

[4]. Kielhorn, Janet; Melber, Christine; Keller, Detlef; Mangelsdorf, Inge (2002). “Palladium – A review of exposure and effects to human health”. International Journal of Hygiene and Environmental Health. 205 (6): 417–32. doi:10.1078/1438-4639-00180. PMID 12455264