Manganese is a chemical element with symbol Mn and atomic number 25. It is not found as a free element in nature. Manganese is a metal with important industrial metal alloy uses in stainless steels. Manganese is named for pyrolusite and other black minerals from the region of Magnesia in Greece, which gave its name to magnesium and the iron ore magnetite. By the mid-18th century, Swedish-German chemist Carl Wilhelm Scheele had used pyrolusite to produce chlorine. Scheele and others were aware that pyrolusite contained a new element, but they were unable to isolate it. Johan Gottlieb Gahn was the first to isolate an impure sample of manganese metal in 1774, which he did by reducing the dioxide with carbon. Manganese phosphating is used for corrosion prevention on steel. Ionized manganese is used industrially as pigments of various colors, which depend on the oxidation state of the ions; the permanganates of alkali and alkaline earth metals are powerful oxidizers. Manganese dioxide is used as the cathode material in alkaline batteries.
In biology, manganese ions function as cofactors for a large variety of enzymes with many functions. Manganese enzymes are essential in detoxification of superoxide free radicals in organisms that must deal with elemental oxygen. Manganese functions in the oxygen-evolving complex of photosynthetic plants. While the element is a required trace mineral for all known living organisms, it acts as a neurotoxin in larger amounts. Through inhalation, it can cause manganism, a condition in mammals leading to neurological damage, sometimes irreversible. Manganese is a silvery-gray metal, it is hard and brittle, difficult to fuse, but easy to oxidize. Manganese metal and its common ions are paramagnetic. Manganese tarnishes in air and oxidizes like iron in water containing dissolved oxygen. Occurring manganese is composed of one stable isotope, 55Mn. Eighteen radioisotopes have been isolated and described, ranging in atomic weight from 46 u to 65 u; the most stable are 53Mn with a half-life of 3.7 million years, 54Mn with a half-life of 312.3 days, 52Mn with a half-life of 5.591 days.
All of the remaining radioactive isotopes have half-lives of less than three hours, the majority of less than one minute. The primary decay mode before the most abundant stable isotope, 55Mn, is electron capture and the primary mode after is beta decay. Manganese has three meta states. Manganese is part of the iron group of elements, which are thought to be synthesized in large stars shortly before the supernova explosion. 53Mn decays to 53Cr with a half-life of 3.7 million years. Because of its short half-life, 53Mn is rare, produced by cosmic rays impact on iron. Manganese isotopic contents are combined with chromium isotopic contents and have found application in isotope geology and radiometric dating. Mn–Cr isotopic ratios reinforce the evidence from 26Al and 107Pd for the early history of the solar system. Variations in 53Cr/52Cr and Mn/Cr ratios from several meteorites suggest an initial 53Mn/55Mn ratio, which indicates that Mn–Cr isotopic composition must result from in situ decay of 53Mn in differentiated planetary bodies.
Hence, 53Mn provides additional evidence for nucleosynthetic processes before coalescence of the solar system. The most common oxidation states of manganese are +2, +3, +4, +6, +7, though all oxidation states from −3 to +7 have been observed. Mn2+ competes with Mg2+ in biological systems. Manganese compounds where manganese is in oxidation state +7, which are restricted to the unstable oxide Mn2O7, compounds of the intensely purple permanganate anion MnO4−, a few oxyhalides, are powerful oxidizing agents. Compounds with oxidation states +5 and +6 are strong oxidizing agents and are vulnerable to disproportionation; the most stable oxidation state for manganese is +2, which has a pale pink color, many manganese compounds are known, such as manganese sulfate and manganese chloride. This oxidation state is seen in the mineral rhodochrosite. Manganese most exists with a high spin, S = 5/2 ground state because of the high pairing energy for manganese. However, there are a few examples of S = 1/2 manganese.
There are no spin-allowed d–d transitions in manganese, explaining why manganese compounds are pale to colorless. The +3 oxidation state is known in compounds like manganese acetate, but these are quite powerful oxidizing agents and prone to disproportionation in solution, forming manganese and manganese. Solid compounds of manganese are characterized by its strong purple-red color and a preference for distorted octahedral coordination resulting from the Jahn-Teller effect; the oxidation state +5 can be produced by dissolving manganese dioxide in molten sodium nitrite. Manganate salts can be produced by dissolving Mn compounds, such as manganese dioxide, in molten alkali while exposed to air. Permanganate compounds are purple, can give glass a violet color. Potassium permanganate, sodium permanganate, barium permanganate are all potent oxidizers. Potassium permanganate called Condy's crystals, is a used laboratory reagent because of its oxidizing properties. Solutions of potassium permanganate were among the first stains and fixatives to be used in the preparation of biological cells and tissues for electron microscopy
In physics and chemistry, ionization energy or ionisation energy, denoted Ei, is the minimum amount of energy required to remove the most loosely bound electron, the valence electron, of an isolated neutral gaseous atom or molecule. It is quantitatively expressed as X + energy → X+ + e−where X is any atom or molecule capable of ionization, X+ is that atom or molecule with an electron removed, e− is the removed electron; this is an endothermic process. The closer the outermost electrons are to the nucleus of the atom, the higher the atom's or element's ionization energy; the sciences of physics and chemistry use different measures of ionization energy. In physics, the unit is the amount of energy required to remove a single electron from a single atom or molecule, expressed as electronvolts. In chemistry, the unit is the amount of energy required for all of the atoms in a mole of substance to lose one electron each: molar ionization energy or enthalpy, expressed as kilojoules per mole or kilocalories per mole.
Comparison of Ei of elements in the periodic table reveals two periodic trends: Ei increases as one moves from left to right within a given period. Ei decreases as one moves from top to bottom in a given group; the latter trend results from the outer electron shell being progressively farther from the nucleus, with the addition of one inner shell per row as one moves down the column. The nth ionization energy refers to the amount of energy required to remove an electron from the species with a charge of. For example, the first three ionization energies are defined as follows: 1st ionization energy X → X+ + e−2nd ionization energy X+ → X2+ + e−3rd ionization energy X2+ → X3+ + e−The term ionization potential is an older name for ionization energy, because the oldest method of measuring ionization energy was based on ionizing a sample and accelerating the electron removed using an electrostatic potential; however this term is now considered obsolete. Some factors affecting the ionization energy include: Nuclear charge: the greater the magnitude of nuclear charge the more the electrons are held by the nucleus and hence more will be ionization energy.
Number of electron shells: the greater the size of the atom less the electrons are held by the nucleus and ionization energy will be less. Effective nuclear charge: the greater the magnitude of electron shielding and penetration the less the electrons are held by the nucleus, the lower the Zeff of the electron, hence less will be the ionization energy. Type of orbital ionized: the atom having a more stable electronic configuration has less tendency to lose electrons and has high ionization energy. Occupancy of the orbital matters: if the orbital is half or filled it is harder to remove electrons Generally, the th ionization energy is larger than the nth ionization energy; when the next ionization energy involves removing an electron from the same electron shell, the increase in ionization energy is due to the increased net charge of the ion from which the electron is being removed. Electrons removed from more charged ions of a particular element experience greater forces of electrostatic attraction.
In addition, when the next ionization energy involves removing an electron from a lower electron shell, the decreased distance between the nucleus and the electron increases both the electrostatic force and the distance over which that force must be overcome to remove the electron. Both of these factors further increase the ionization energy; some values for elements of the third period are given in the following table: Large jumps in the successive molar ionization energies occur when passing noble gas configurations. For example, as can be seen in the table above, the first two molar ionization energies of magnesium are much smaller than the third, which requires stripping off a 2p electron from the neon configuration of Mg2+; that electron is much closer to the nucleus than the previous 3s electron. Ionization energy is a periodic trend within the periodic table organization. Moving left to right within a period, or upward within a group, the first ionization energy increases, with some exceptions such as aluminum and sulfur in the table above.
As the nuclear charge of the nucleus increases across the period, the atomic radius decreases and the electron cloud becomes closer towards the nucleus. Atomic ionization energy can be predicted by an analysis using electrostatic potential and the Bohr model of the atom, as follows. Consider an electron of charge -e and an atomic nucleus with charge +Ze, where Z is the number of protons in the nucleus. According to the Bohr model, if the electron were to approach and bond with the atom, it would come to rest at a certain radius a; the electrostatic potential V at distance a from the ionic nucleus, referenced to a point infinitely far away, is: V = Z e a Since the electron is negatively charged, it is drawn inwards by this positive electrostatic potential. The energy required for the electron to "climb out" and leave the atom is: E = e V = Z e 2 a This analysis is incomplete, as it leaves the distance a as an unknown variable, it can be made more rigorous by assigning to each electron of every chemical element a characteristic distance, ch
Neon is a chemical element with symbol Ne and atomic number 10. It is a noble gas. Neon is a colorless, inert monatomic gas under standard conditions, with about two-thirds the density of air, it was discovered in 1898 as one of the three residual rare inert elements remaining in dry air, after nitrogen, oxygen and carbon dioxide were removed. Neon was the second of these three rare gases to be discovered and was recognized as a new element from its bright red emission spectrum; the name neon is derived from the Greek νέον, neuter singular form of νέος, meaning new. Neon is chemically inert, no uncharged neon compounds are known; the compounds of neon known include ionic molecules, molecules held together by van der Waals forces and clathrates. During cosmic nucleogenesis of the elements, large amounts of neon are built up from the alpha-capture fusion process in stars. Although neon is a common element in the universe and solar system, it is rare on Earth, it composes about 18.2 ppm of air by a smaller fraction in Earth's crust.
The reason for neon's relative scarcity on Earth and the inner planets is that neon is volatile and forms no compounds to fix it to solids. As a result, it escaped from the planetesimals under the warmth of the newly ignited Sun in the early Solar System; the outer atmosphere of Jupiter is somewhat depleted of neon, although for a different reason. It is lighter than air, causing it to escape from Earth's atmosphere. Neon gives a distinct reddish-orange glow when used in low-voltage neon glow lamps, high-voltage discharge tubes and neon advertising signs; the red emission line from neon causes the well known red light of helium–neon lasers. Neon has few other commercial uses, it is commercially extracted by the fractional distillation of liquid air. Since air is the only source, it is more expensive than helium. Neon was discovered in 1898 by the British chemists Sir William Ramsay and Morris W. Travers in London. Neon was discovered when Ramsay chilled a sample of air until it became a liquid warmed the liquid and captured the gases as they boiled off.
The gases nitrogen and argon had been identified, but the remaining gases were isolated in their order of abundance, in a six-week period beginning at the end of May 1898. First to be identified was krypton; the next, after krypton had been removed, was a gas which gave a brilliant red light under spectroscopic discharge. This gas, identified in June, was named "neon", the Greek analogue of the Latin novum suggested by Ramsay's son; the characteristic brilliant red-orange color emitted by gaseous neon when excited electrically was noted immediately. Travers wrote: "the blaze of crimson light from the tube told its own story and was a sight to dwell upon and never forget."A second gas was reported along with neon, having the same density as argon but with a different spectrum – Ramsay and Travers named it metargon. However, subsequent spectroscopic analysis revealed it to be argon contaminated with carbon monoxide; the same team discovered xenon by the same process, in September 1898. Neon's scarcity precluded its prompt application for lighting along the lines of Moore tubes, which used nitrogen and which were commercialized in the early 1900s.
After 1902, Georges Claude's company Air Liquide produced industrial quantities of neon as a byproduct of his air-liquefaction business. In December 1910 Claude demonstrated modern neon lighting based on a sealed tube of neon. Claude tried to sell neon tubes for indoor domestic lighting, due to their intensity, but the market failed because homeowners objected to the color. In 1912, Claude's associate began selling neon discharge tubes as eye-catching advertising signs and was more successful. Neon tubes were introduced to the U. S. in 1923 with two large neon signs bought by a Los Angeles Packard car dealership. The glow and arresting red color made neon advertising different from the competition; the intense color and vibrancy of neon equated with American society at the time, suggesting a "century of progress" and transforming cities into sensational new environments filled with radiating advertisements and "electro-graphic architecture". Neon played a role in the basic understanding of the nature of atoms in 1913, when J. J. Thomson, as part of his exploration into the composition of canal rays, channeled streams of neon ions through a magnetic and an electric field and measured the deflection of the streams with a photographic plate.
Thomson observed two separate patches of light on the photographic plate, which suggested two different parabolas of deflection. Thomson concluded that some of the atoms in the neon gas were of higher mass than the rest. Though not understood at the time by Thomson, this was the first discovery of isotopes of stable atoms. Thomson's device was a crude version of the instrument. Neon is the second lightest inert gas. Neon has three stable isotopes: 21Ne and 22Ne. 21Ne and 22Ne are primordial and nucleogenic and their variations in natural abundance are well understood. In contrast, 20Ne is not known to be radiogenic; the causes of the variation of 20Ne in the Earth have thus been hotly debated. The princ
Argon is a chemical element with symbol Ar and atomic number 18. It is a noble gas. Argon is the third-most abundant gas in the Earth's atmosphere, at 0.934%. It is more than twice as abundant as water vapor, 23 times as abundant as carbon dioxide, more than 500 times as abundant as neon. Argon is the most abundant noble gas in Earth's crust, comprising 0.00015% of the crust. Nearly all of the argon in the Earth's atmosphere is radiogenic argon-40, derived from the decay of potassium-40 in the Earth's crust. In the universe, argon-36 is by far the most common argon isotope, as it is the most produced by stellar nucleosynthesis in supernovas; the name "argon" is derived from the Greek word ἀργόν, neuter singular form of ἀργός meaning "lazy" or "inactive", as a reference to the fact that the element undergoes no chemical reactions. The complete octet in the outer atomic shell makes argon stable and resistant to bonding with other elements, its triple point temperature of 83.8058 K is a defining fixed point in the International Temperature Scale of 1990.
Argon is produced industrially by the fractional distillation of liquid air. Argon is used as an inert shielding gas in welding and other high-temperature industrial processes where ordinarily unreactive substances become reactive. Argon is used in incandescent, fluorescent lighting, other gas-discharge tubes. Argon makes a distinctive blue-green gas laser. Argon is used in fluorescent glow starters. Argon has the same solubility in water as oxygen and is 2.5 times more soluble in water than nitrogen. Argon is colorless, odorless and nontoxic as a solid, liquid or gas. Argon is chemically inert under most conditions and forms no confirmed stable compounds at room temperature. Although argon is a noble gas, it can form some compounds under various extreme conditions. Argon fluorohydride, a compound of argon with fluorine and hydrogen, stable below 17 K, has been demonstrated. Although the neutral ground-state chemical compounds of argon are presently limited to HArF, argon can form clathrates with water when atoms of argon are trapped in a lattice of water molecules.
Ions, such as ArH+, excited-state complexes, such as ArF, have been demonstrated. Theoretical calculation predicts several more argon compounds that should be stable but have not yet been synthesized. Argon, is named in reference to its chemical inactivity; this chemical property of this first noble gas to be discovered impressed the namers. An unreactive gas was suspected to be a component of air by Henry Cavendish in 1785. Argon was first isolated from air in 1894 by Lord Rayleigh and Sir William Ramsay at University College London by removing oxygen, carbon dioxide and nitrogen from a sample of clean air, they had determined that nitrogen produced from chemical compounds was 0.5% lighter than nitrogen from the atmosphere. The difference was slight, they concluded. Argon was encountered in 1882 through independent research of H. F. Newall and W. N. Hartley; each observed new lines in the emission spectrum of air. Until 1957, the symbol for argon was "A", but now is "Ar". Argon constitutes 0.934% by volume and 1.288% by mass of the Earth's atmosphere, air is the primary industrial source of purified argon products.
Argon is isolated from air by fractionation, most by cryogenic fractional distillation, a process that produces purified nitrogen, neon and xenon. The Earth's crust and seawater contain 0.45 ppm of argon, respectively. The main isotopes of argon found on Earth are 40Ar, 36Ar, 38Ar. Occurring 40K, with a half-life of 1.25×109 years, decays to stable 40Ar by electron capture or positron emission, to stable 40Ca by beta decay. These properties and ratios are used to determine the age of rocks by K–Ar dating. In the Earth's atmosphere, 39Ar is made by cosmic ray activity by neutron capture of 40Ar followed by two-neutron emission. In the subsurface environment, it is produced through neutron capture by 39K, followed by proton emission. 37Ar is created from the neutron capture by 40Ca followed by an alpha particle emission as a result of subsurface nuclear explosions. It has a half-life of 35 days. Between locations in the Solar System, the isotopic composition of argon varies greatly. Where the major source of argon is the decay of 40K in rocks, 40Ar will be the dominant isotope, as it is on Earth.
Argon produced directly by stellar nucleosynthesis, is dominated by the alpha-process nuclide 36Ar. Correspondingly, solar argon contains 84.6% 36Ar, the ratio of the three isotopes 36Ar: 38Ar: 40Ar in the atmospheres of the outer planets is 8400: 1600: 1. This contrasts with the low abundance of primordial 36Ar in Earth's atmosphere, only 31.5 ppmv, comparable with that of neon on Earth and with interplanetary gasses, measured by probes. The atmospheres of Mars and Titan contain argon, predominantly as 40Ar, its content may be as high as 1.93%. The predominance of radiogenic 40Ar is the reason the standard atomic weight of terrestrial argon is greater than that of the next element, potassium, a fact that was
The boron group are the chemical elements in group 13 of the periodic table, comprising boron, gallium, indium and also the chemically uncharacterized nihonium. The elements in the boron group are characterized by having three electrons in their outer energy levels; these elements have been referred to as the triels. Boron is classified as a typical non-metal while the rest, with the possible exception of nihonium, are considered post-transition metals. Boron occurs sparsely because bombardment by the subatomic particles produced from natural radioactivity disrupts its nuclei. Aluminium occurs on earth, indeed is the third most abundant element in the Earth's crust. Gallium is found in the earth with an abundance of 13 ppm. Indium is the 61st most abundant element in the earth's crust, thallium is found in moderate amounts throughout the planet. Nihonium is never found in nature and therefore is termed a synthetic element. Several group 13 elements have biological roles in the ecosystem. Boron is essential for some plants.
Lack of boron can lead to stunted plant growth, while an excess can cause harm by inhibiting growth. Aluminium is considered safe. Indium and gallium can stimulate metabolism. Thallium is toxic, interfering with the function of numerous vital enzymes, has seen use as a pesticide. Like other groups, the members of this family show patterns in electron configuration in the outermost shells, resulting in trends in chemical behavior: The boron group is notable for trends in the electron configuration, as shown above, in some of its elements' characteristics. Boron differs from the other group members in its hardness and reluctance to participate in metallic bonding. An example of a trend in reactivity is boron's tendency to form reactive compounds with hydrogen. Most of the elements in the boron group show increasing reactivity as the elements get heavier in atomic mass and higher in atomic number. Boron, the first element in the group, is unreactive with many elements except at high temperatures, although it is capable of forming many compounds with hydrogen, sometimes called boranes.
The simplest borane is diborane, or B2H6. Another example is B10H14; the next group-13 elements and gallium, form fewer stable hydrides, although both AlH3 and GaH3 exist. Indium, the next element in the group, is not known to form many hydrides, except in complex compounds such as the phosphine complex H3InP3. No stable compound of thallium and hydrogen has been synthesized in any laboratory. All of the boron-group elements are known to form a trivalent oxide, with two atoms of the element bonded covalently with three atoms of oxygen; these elements show a trend of increasing pH. Boron oxide is acidic and gallium oxide are amphoteric, indium oxide is nearly amphoteric, thallium oxide is a Lewis base because it dissolves in acids to form salts; each of these compounds are stable, but thallium oxide decomposes at temperatures higher than 875 °C. The elements in group 13 are capable of forming stable compounds with the halogens with the formula MX3 Fluorine, the first halogen, is able to form stable compounds with every element, tested, the boron group is no exception.
It is hypothesized that nihonium could form a compound with fluorine, NhF3, before spontaneously decaying due to nihonium's radioactivity. Chlorine forms stable compounds with all of the elements in the boron group, including thallium, is hypothesized to react with nihonium. All of the elements will react with bromine under the right conditions, as with the other halogens but less vigorously than either chlorine or fluorine. Iodine will react with all natural elements in the periodic table except for the noble gases, is notable for its explosive reaction with aluminium to form 2AlI3. Astatine, the heaviest halogen, has only formed a few compounds, due to its radioactivity and short half-life, no reports of a compound with an At–B, –Al, –Ga, –In, –Tl, or –Nh bond have been seen, although scientists think that it should form salts with metals, it has been noticed that the elements in the boron group have similar physical properties, although most of boron's are exceptional. For example, all of the elements in the boron group, are soft.
Moreover, all of the other elements in group 13 are reactive at moderate temperatures, while boron's reactivity only becomes comparable at high temperatures. One characteristic that all do have in common is having three electrons in their valence shells. Boron, being a metalloid, is a thermal and electrical insulator at room temperature, but a good conductor of heat and electricity at high temperatures. Unlike boron, the metals in the group are good conductors under normal conditions; this is in accordance with the long-standing generalization that all metals conduct heat and electricity better than most non-metals. The inert s-pair effect is significant in the group-13 elements the heavier ones like thallium; this results in a variety of oxidation states. In the lighter elements, the +3 state is the most stable, but the +1 state becomes more prevalent with increasing atomic number, is the most stable for thallium. Boron is capable of forming compounds with lower oxidization states, of +1 or +2, aluminium can do the same.
Gallium can form compounds with the oxid
A permanganate is the general name for a chemical compound containing the manganate ion. Because manganese is in the +7 oxidation state, the permanganate ion is a strong oxidizing agent; the ion has tetrahedral geometry. Permanganate solutions are purple in color and are stable in neutral or alkaline media; the exact chemical reaction is dependent upon the organic contaminants present and the oxidant utilized. For example, trichloroethene is oxidized by sodium permanganate to form carbon dioxide, manganese dioxide, sodium ions, hydronium ions, chloride ions. In an acidic solution, permanganate is reduced to the pale pink +2 oxidation state of the manganese ion. 8 H+ + MnO−4 + 5 e− → Mn2+ + 4 H2OIn a basic solution, permanganate is reduced to the green +6 oxidation state of the manganate ion, MnO2−4. MnO−4 + e− → MnO2−4In a neutral medium, however, it gets reduced to the brown +4 oxidation state of manganese dioxide MnO2. 2 H2O + MnO−4 + 3 e− → MnO2 + 4 OH− Permanganates can be produced by oxidation of manganese compounds such as manganese chloride or manganese sulfate by strong oxidizing agents, for instance, sodium hypochlorite or lead dioxide: 2 MnCl2 + 5 NaClO + 6 NaOH → 2 NaMnO4 + 9 NaCl + 3 H2O 2 MnSO4 + 5 PbO2 + 3 H2SO4 → 2 HMnO4 + 5 PbSO4 + 2 H2OIt may be produced by the disproportionation of manganates, with manganese dioxide as a side-product: 3 Na2MnO4 + 2 H2O → 2 NaMnO4 + MnO2 + 4 NaOHThey are produced commercially by electrolysis or air oxidation of alkaline solutions of manganate salts.
Permanganates are salts of permanganic acid. They have a deep purple colour, due to a charge transfer transition. Permanganate is a strong oxidizer, similar to perchlorate, it is therefore in common use in qualitative analysis. According to theory, permanganate is strong enough to oxidize water, but this does not happen to any extent. Besides this, it is stable, it is a useful reagent, though with organic compounds, not selective. Potassium permanganate is used as a disinfectant. Manganates are not stable thermally. For instance, potassium permanganate decomposes at 230 °C to potassium manganate and manganese dioxide, releasing oxygen gas: 2 KMnO4 → K2MnO4 + MnO2 + O2A permanganate can oxidize an amine to a nitro compound, an alcohol to a ketone, an aldehyde to a carboxylic acid, a terminal alkene to a carboxylic acid, oxalic acid to carbon dioxide, an alkene to a diol; this list is not exhaustive. In alkene oxidations one intermediate is a cyclic Mn species: Ammonium permanganate, NH4MnO4 Calcium permanganate, Ca2 Potassium permanganate, KMnO4 Sodium permanganate, NaMnO4 Silver permanganate, AgMnO4 Perchlorate, a similar ion with a chlorine center Chromate, isoelectronic with permanganate Pertechnetate