Oleum, or fuming sulfuric acid, is a solution of various compositions of sulfur trioxide in sulfuric acid, or sometimes more to disulfuric acid. Oleum is identified by the CAS number 8014-95-7. Oleums can be described by the formula ySO3. H2O where y is the total molar sulfur trioxide content; the value of y can be varied. They can be described by the formula H2SO4.xSO3 where x is now defined as the molar free sulfur trioxide content. Oleum is assessed according to the free SO3 content by mass, it can be expressed as a percentage of sulfuric acid strength. For example, 10% oleum can be expressed as H2SO4.0.13611SO3, 1.0225SO3. H2O or 102.25% sulfuric acid. The conversion between % acid and % oleum is: % acid = 100 + 18/80 × % oleum A value for x of 1 gives the empirical formula H2S2O7 for disulfuric acid. Pure disulfuric acid is a solid at room temperature, melting at 36 °C and used either in the laboratory or industrial processes. Oleum is produced in the contact process, where sulfur is oxidized to sulfur trioxide, subsequently dissolved in concentrated sulfuric acid.
Sulfuric acid. The lead chamber process for sulfuric acid production was abandoned because it could not produce sulfur trioxide or concentrated sulfuric acid directly due to corrosion of the lead, absorption of NO2 gas; until this process was made obsolete by the contact process, oleum had to be obtained through indirect methods. The biggest production of oleum came from the distillation of iron sulfates at Nordhausen, from which the historical name Nordhausen sulfuric acid is derived. Oleum is an important intermediate in the manufacture of sulfuric acid due to its high enthalpy of hydration; when SO3 is added to water, rather than dissolving, it tends to form a fine mist of sulfuric acid, difficult to manage. However, SO3 added to concentrated sulfuric acid dissolves, forming oleum which can be diluted with water to produce additional concentrated sulfuric acid. Oleum is a useful form for transporting sulfuric acid compounds in rail tank cars, between oil refineries and industrial consumers.
Certain compositions of oleum are solid at room temperature, thus are safer to ship than as a liquid. Solid oleum can be converted into liquid at the destination by steam heating or dilution or concentration; this requires care to prevent overheating and evaporation of sulfur trioxide. To extract it from a tank car requires careful heating using steam conduits inside the tank car. Great care must be taken to avoid overheating, as this can increase the pressure in the tank car beyond the tank's safety valve limit. In addition, oleum is less corrosive to metals than sulfuric acid, because there is no free water to attack surfaces; because of that, sulfuric acid is sometimes concentrated to oleum for in-plant pipelines and diluted back to acid for use in industrial reactions. In Richmond, California in 1993 a significant release occurred due to overheating, causing a release of sulfur trioxide that absorbed moisture from the atmosphere, creating a mist of micrometre-sized sulfuric acid particles that formed an inhalation health hazard.
This mist spread over a wide area. Oleum is a harsh reagent, is corrosive. One important use of oleum as a reagent is the secondary nitration of nitrobenzene; the first nitration can occur with nitric acid in sulfuric acid, but this deactivates the ring towards further electrophilic substitution. A stronger reagent, oleum, is needed to introduce the second nitro group onto the aromatic ring. Oleum is used in the manufacture of many explosives with the notable exception of nitrocellulose; the chemical requirements for explosives manufacture require anhydrous mixtures containing nitric acid and sulfuric acid. Ordinary commercial grade nitric acid consists of the constant boiling azeotrope of nitric acid and water, contains 68% nitric acid. Mixtures of ordinary nitric acid in sulfuric acid therefore contain substantial amounts of water and are unsuitable for processes such as those that occur in the manufacture of trinitrotoluene; the synthesis of RDX and certain other explosives does not require oleum.
Anhydrous nitric acid, referred to as white fuming nitric acid, can be used to prepare water-free nitration mixtures, this method is used in laboratory scale operations where the cost of material is not of primary importance. Fuming nitric acid is hazardous to handle and transport, because it is corrosive and volatile. For industrial use, such strong nitration mixtures are prepared by mixing oleum with ordinary commercial nitric acid so that the free sulfur trioxide in the oleum consumes the water in the nitric acid. Like concentrated sulfuric acid, oleum is such a strong dehydrating agent that if poured onto powdered glucose, or any other sugar, it will draw the elements of water out of the sugar in an exothermic reaction, leaving nearly pure carbon as a solid; this carbon expands outward, hardening as a solid black substance with gas bubbles in it
Le Chatelier's principle
Le Chatelier's principle called Chatelier's principle or "The Equilibrium Law", can be used to predict the effect of a change in conditions on some chemical equilibria. The principle is named after Henry Louis Le Chatelier and sometimes Karl Ferdinand Braun who discovered it independently, it can be stated as: When any system at equilibrium for a long period of time is subjected to change in concentration, volume, or pressure, the system changes to a new equilibrium and this change counteracts the applied change. It is common to treat the principle as a more general observation, such as When a settled system is disturbed, it will adjust to diminish the change, made to it, or, “roughly stated”, Any change in status quo prompts an opposing reaction in the responding system or "The System always kicks back"; the principle has a variety of names, depending upon the discipline using it. In chemistry, the principle is used to manipulate the outcomes of reversible reactions to increase the yield of reactions.
In pharmacology, the binding of ligands to the receptor may shift the equilibrium according to Le Chatelier's principle, thereby explaining the diverse phenomena of receptor activation and desensitization. In economics, the principle has been generalized to help explain the price equilibrium of efficient economic systems. Phenomena in apparent contradiction to Le Chatelier's principle can arise in systems of simultaneous equilibrium: see the article on the theory of response reactions. Le Chatelier's principle describes the qualitative behavior of systems where there is an externally induced, instantaneous change in one parameter of a system; the duration of adjustment depends on the strength of the negative feedback to the initial shock. Where a shock induces positive feedback, the new equilibrium can be far from the old one, can take a long time to reach. In some dynamic systems, the end-state cannot be determined from the shock; the principle is used to describe closed negative-feedback systems, but applies, in general, to thermodynamically closed and isolated systems in nature, since the second law of thermodynamics ensures that the disequilibrium caused by an instantaneous shock must have a finite half-life.
The principle has analogs throughout the entire physical world. The principle while well rooted in chemical equilibrium and extended into economic theory, can be used in describing mechanical systems in that the system put under stress will respond in a way such as to reduce or minimize that stress. Moreover, the response will be via the mechanism that most relieves that stress. Shear pins and other such sacrificial devices are design elements that protect systems against stress applied in undesired manners to relieve it so as to prevent more extensive damage to the entire system, a practical engineering application of Le Chatelier's principle. Changing the concentration of a chemical will shift the equilibrium to the side that would reduce that change in concentration; the chemical system will attempt to oppose the change affected to the original state of equilibrium. In turn, the rate of reaction and yield of products will be altered corresponding to the impact on the system; this can be illustrated by the equilibrium of carbon monoxide and hydrogen gas, reacting to form methanol.
CO + 2 H2 ⇌ CH3OHSuppose. Using Le Chatelier's principle, we can predict that the amount of methanol will increase, decreasing the total change in CO. If we are to add a species to the overall reaction, the reaction will favor the side opposing the addition of the species; the subtraction of a species would cause the reaction to "fill the gap" and favor the side where the species was reduced. This observation is supported by the collision theory; as the concentration of CO is increased, the frequency of successful collisions of that reactant would increase allowing for an increase in forward reaction, generation of the product. If the desired product is not thermodynamically favored, the end-product can be obtained if it is continuously removed from the solution; the effect of a change in concentration is exploited synthetically for condensation reactions that are equilibrium processes. This can be achieved by physically sequestering water, by adding desiccants like anhydrous magnesium sulfate or molecular sieves, or by continuous removal of water by distillation facilitated by a Dean-Stark apparatus.
The effect of changing the temperature in the equilibrium can be made clear by 1) incorporating heat as either a reactant or a product, 2) assuming that an increase in temperature increases the heat content of a system. When the reaction is exothermic, heat is included as a product, when the reaction is endothermic, heat is included as a reactant. Hence, whether increasing or decreasing the temperature would favor the forward or the reverse reaction can be determined by applying the same principle as with concentration changes. Take, for example, the reversible reaction of nitrogen gas with hydrogen gas to form ammonia: N2 + 3 H2 ⇌ 2 NH3 ΔH = -92 kJ mol−1Because this reaction is exothermic, it produces heat: N2 + 3 H2 ⇌ 2 NH3 + heatIf the temperature was increased, the heat content of the system would increase, so the system would consume some of that heat b
Vanadium oxide is the inorganic compound with the formula V2O5. Known as vanadium pentoxide, it is a brown/yellow solid, although when freshly precipitated from aqueous solution, its colour is deep orange; because of its high oxidation state, it is an oxidizing agent. From the industrial perspective, it is the most important compound of vanadium, being the principal precursor to alloys of vanadium and is a used industrial catalyst; the mineral form of this compound, shcherbinaite, is rare always found among fumaroles. A mineral trihydrate, V2O5·3H2O, is known under the name of navajoite. Upon heating a mixture of vanadium oxide and vanadium oxide, comproportionation occurs to give vanadium oxide, as a deep-blue solid: V2O5 + V2O3 → 4 VO2The reduction can be effected by oxalic acid, carbon monoxide, sulfur dioxide. Further reduction using hydrogen or excess CO can lead to complex mixtures of oxides such as V4O7 and V5O9 before black V2O3 is reached. V2O5 is an amphoteric oxide. Unlike most metal oxides, it dissolves in water to give a pale yellow, acidic solution.
Thus V2O5 reacts with strong non-reducing acids to form solutions containing the pale yellow salts containing dioxovanadium centers: V2O5 + 2 HNO3 → 2 VO2 + H2OIt reacts with strong alkali to form polyoxovanadates, which have a complex structure that depends on pH. If excess aqueous sodium hydroxide is used, the product is a colourless salt, sodium orthovanadate, Na3VO4. If acid is added to a solution of Na3VO4, the colour deepens through orange to red before brown hydrated V2O5 precipitates around pH 2; these solutions contain the ions HVO42− and V2O74− between pH 9 and pH 13, but below pH 9 more exotic species such as V4O124− and HV10O285− predominate. Upon treatment with thionyl chloride, it converts to the volatile liquid vanadium oxychloride, VOCl3: V2O5 + 3 SOCl2 → 2 VOCl3 + 3 SO2 Hydrochloric acid and hydrobromic acid are oxidised to the corresponding halogen, e.g. V2O5 + 6HCl + 7H2O → 22+ + 4Cl− + Cl2Vanadates or vanadyl compounds in acid solution are reduced by zinc amalgam through the colourful pathway: The ions are all hydrated to varying degrees.
Technical grade V2O5 is produced as a black powder used for the production of vanadium metal and ferrovanadium. A vanadium ore or vanadium-rich residue is treated with sodium carbonate and an ammonium salt to produce sodium metavanadate, NaVO3; this material is acidified to pH 2–3 using H2SO4 to yield a precipitate of "red cake". The red cake is melted at 690 °C to produce the crude V2O5. Vanadium oxide is produced when vanadium metal is heated with excess oxygen, but this product is contaminated with other, lower oxides. A more satisfactory laboratory preparation involves the decomposition of ammonium metavanadate at 500-550 °C: 2 NH4VO3 → V2O5 + 2 NH3 + H2O In terms of quantity, the dominant use for vanadium oxide is in the production of ferrovanadium; the oxide is heated with scrap iron and ferrosilicon, with lime added to form a calcium silicate slag. Aluminium may be used, producing the iron-vanadium alloy along with alumina as a by-product. Another important use of vanadium oxide is in the manufacture of sulfuric acid, an important industrial chemical with an annual worldwide production of 165 million metric tons in 2001, with an approximate value of US$8 billion.
Vanadium oxide serves the crucial purpose of catalysing the mildly exothermic oxidation of sulfur dioxide to sulfur trioxide by air in the contact process: 2 SO2 + O2 ⇌ 2 SO3The discovery of this simple reaction, for which V2O5 is the most effective catalyst, allowed sulfuric acid to become the cheap commodity chemical it is today. The reaction is performed between 400 and 620 °C. Since it is known that V2O5 can be reduced to VO2 by SO2, one catalytic cycle is as follows: SO2 + V2O5 → SO3 + 2VO2followed by 2VO2 +½O2 → V2O5It is used as catalyst in the selective catalytic reduction of NOx emissions in some power plants. Due to its effectiveness in converting sulfur dioxide into sulfur trioxide, thereby sulfuric acid, special care must be taken with the operating temperatures and placement of a power plant's SCR unit when firing sulfur-containing fuels. Maleic anhydride is produced by the V2O5-catalysed oxidation of butane with air: C4H10 + 4 O2 → C2H22O + 8 H2OMaleic anhydride is used for the production of polyester resins and alkyd resins.
Phthalic anhydride is produced by V2O5-catalysed oxidation of ortho-xylene or naphthalene at 350–400 °C. The equation is for the xylene oxidation: C6H42 + 3 O2 → C6H42O + 3 H2OPhthalic anhydride is a precursor to plasticisers, used for conferring pliability to polymers. A variety of other industrial compounds are produced including adipic acid, acrylic acid, oxalic acid, anthraquinone. Due to its high coefficient of thermal resistance, vanadium oxide finds use as a detector material in bolometers and microbolometer arrays for thermal imaging, it finds application as an ethanol sensor in ppm levels. Vanadium redox batteries are a type of flow battery used for energy storage, including large power facilities such as wind farms. Vanadium oxide exhibits modest acute toxicity to humans, with an LD50 of about 470 mg/kg; the greater hazard is with inhalation of the dust, where the LD50 ranges from 4–11 mg/kg for a 14-day exposure. Vanadate, formed by hydrolysis of V2O5 at high pH, appears to inhibit enzymes that process phosphate.
However the mode of action remains elusive. "Vanadium Pentoxide", Cobalt in Hard Metal
Sulfuric acid known as vitriol, is a mineral acid composed of the elements sulfur and hydrogen, with molecular formula H2SO4. It is a colorless and syrupy liquid, soluble in water, in a reaction, exothermic, its corrosiveness can be ascribed to its strong acidic nature, and, if at a high concentration, its dehydrating and oxidizing properties. It is hygroscopic absorbing water vapor from the air. Upon contact, sulfuric acid can cause severe chemical burns and secondary thermal burns. Sulfuric acid is a important commodity chemical, a nation's sulfuric acid production is a good indicator of its industrial strength, it is produced with different methods, such as contact process, wet sulfuric acid process, lead chamber process and some other methods. Sulfuric acid is a key substance in the chemical industry, it is most used in fertilizer manufacture, but is important in mineral processing, oil refining, wastewater processing, chemical synthesis. It has a wide range of end applications including in domestic acidic drain cleaners, as an electrolyte in lead-acid batteries, in various cleaning agents.
Although nearly 100% sulfuric acid can be made, the subsequent loss of SO3 at the boiling point brings the concentration to 98.3% acid. The 98.3% grade is more stable in storage, is the usual form of what is described as "concentrated sulfuric acid". Other concentrations are used for different purposes; some common concentrations are: "Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process, chamber acid being the acid produced in the lead chamber itself and tower acid being the acid recovered from the bottom of the Glover tower. They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10M" sulfuric acid is prepared by adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C or higher. Sulfuric acid reacts with its anhydride, SO3, to form H2S2O7, called pyrosulfuric acid, fuming sulfuric acid, Disulfuric acid or oleum or, less Nordhausen acid.
Concentrations of oleum are either expressed in terms of % SO3 or as % H2SO4. Pure H2S2O7 is a solid with melting point of 36 °C. Pure sulfuric acid has a vapor pressure of <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C, 98% sulfuric acid has a <1 mmHg vapor pressure at 40 °C. Pure sulfuric acid is a viscous clear liquid, like oil, this explains the old name of the acid. Commercial sulfuric acid is sold in several different purity grades. Technical grade H2SO4 is impure and colored, but is suitable for making fertilizer. Pure grades, such as United States Pharmacopeia grade, are used for making pharmaceuticals and dyestuffs. Analytical grades are available. Nine hydrates are known, but four of them were confirmed to be tetrahydrate and octahydrate. Anhydrous H2SO4 is a polar liquid, having a dielectric constant of around 100, it has a high electrical conductivity, caused by dissociation through protonating itself, a process known as autoprotolysis. 2 H2SO4 ⇌ H3SO+4 + HSO−4The equilibrium constant for the autoprotolysis is Kap = = 2.7×10−4The comparable equilibrium constant for water, Kw is 10−14, a factor of 1010 smaller.
In spite of the viscosity of the acid, the effective conductivities of the H3SO+4 and HSO−4 ions are high due to an intramolecular proton-switch mechanism, making sulfuric acid a good conductor of electricity. It is an excellent solvent for many reactions; because the hydration reaction of sulfuric acid is exothermic, dilution should always be performed by adding the acid to the water rather than the water to the acid. Because the reaction is in an equilibrium that favors the rapid protonation of water, addition of acid to the water ensures that the acid is the limiting reagent; this reaction is best thought of as the formation of hydronium ions: H2SO4 + H2O → H3O+ + HSO−4 Ka1 = 2.4×106 HSO−4 + H2O → H3O+ + SO2−4 Ka2 = 1.0×10−2 HSO−4 is the bisulfate anion and SO2−4 is the sulfate anion. Ka1 and Ka2 are the acid dissociation constants; because the hydration of sulfuric acid is thermodynamically favorable and the affinity of it for water is sufficiently strong, sulfuric acid is an excellent dehydrating agent.
Concentrated sulfuric acid has a powerful dehydrating property, removing water from other chemical compounds including sugar and other carbohydrates and producing carbon and steam. In the laboratory, this is demonstrated by mixing table sugar into sulfuric acid; the sugar changes from white to dark brown and to black as carbon is formed. A rigid column of black, porous carbon will emerge as well; the carbon will smell of caramel due to the heat generated. C 12 H 22 O 11 ⏞ sucrose → H 2 SO 4 12 C + 11 H 2
An exothermic reaction is a chemical reaction that releases energy through light or heat. It is the opposite of an endothermic reaction. Expressed in a chemical equation: reactants → products + energy. Exothermic Reaction means "thermic" means heat. So the reaction in which there is release of heat with or without light is called exothermic reaction. An exothermic reaction is a chemical reaction, it gives net energy to its surroundings. That is, the energy needed to initiate the reaction is less than the energy released; when the medium in which the reaction is taking place collects heat, the reaction is exothermic. When using a calorimeter, the total amount of heat that flows into the calorimeter is the negative of the net change in energy of the system; the absolute amount of energy in a chemical system is difficult to calculate. The enthalpy change, ΔH, of a chemical reaction is much easier to work with; the enthalpy change equals the change in internal energy of the system plus the work needed to change the volume of the system against constant ambient pressure.
A bomb calorimeter is suitable for measuring the energy change, ΔH, of a combustion reaction. Measured and calculated ΔH values are related to bond energies by: ΔH = − In an exothermic reaction, by definition, the enthalpy change has a negative value: ΔH < 0since a larger value is subtracted from a smaller value. For example, when hydrogen burns: 2H2 + O2 → 2H2O ΔH = −483.6 kJ/mol of O2 In an adiabatic system, the temperature raise due to enthalpy change can be expressed as −ΔH298.15 K = ∫T1T0Cp, pdT + ∫T0298 KdTwhere ΔH298.15 K is the standard enthalpy of reaction at 298 K, T0 and T1 are the initial and final temperature of the system and Cp,p and Cp,r are the heat capacities of the product and reactant, respectively. Assuming the heat capacity of the system remains as a constant value Cp,p=Cp,r=Cp, the change of temperature ΔT=T1−T0 can be expressed as −ΔH298.15 K = ∫T0+ΔTT0Cp, pdT = ΔTCp, pThe most available hand warmers make use of the oxidation of iron to achieve an exothermic reaction: 4Fe + 3O2 → 2Fe2O3 .
Combustion reactions of fuels or a substance e.g. Burning of natural gas: CH 4 + 2 O 2 ⟶ CO 2 + 2 H 2 O C 6 H 12 O 6 + 6 O 2 ⟶ 6 CO 2 + 6 H 2 O Neutralization The thermite reaction Reactions taking place in a self-heating can based on lime aluminium Many corrosion reactions such as oxidation of metals Most polymerization reactions The Haber process of ammonia production Respiration Decomposition of vegetable matter into compost Solution of sulfuric acid into water Dehydration of sugars upon contact with sulfuric acid Detonation of nitroglycerin Nuclear fission of uranium-235 The concept and its opposite number endothermic relate to the enthalpy change in any process, not just chemical reactions. In endergonic reactions and exergonic reactions it is the sign of the Gibbs free energy that determines the equilibrium point, not enthalpy; the related concepts endergonic and exergonic apply to all physical processes. The conceptually related endotherm and ectotherm are concepts in animal physiology.
In quantum numbers, when any excited energy level goes down to its original level for example: when n=4 fall to n=2, energy is released so, it is exothermic. Where an exothermic reaction causes heating of the reaction vessel, not controlled, the rate of reaction can increase, in turn causing heat to be evolved more quickly; this positive feedback situation is known as thermal runaway. An explosion can result from the problem. Heat production or absorption in either a physical process or chemical reaction is measured using calorimetry. One common laboratory instrument is the reaction calorimeter, where the heat flow into or from the reaction vessel is monitored; the technique can be used to follow chemical reactions as well as physical processes such as crystallization and dissolution. Energy released is measured in Joule per mole; the reaction has a negative ΔH value due to heat loss. E.g.: -123 J/mol Chemical thermodynamics Differential scanning calorimetry Endergonic Exergonic Endergonic reaction Exergonic reaction Exothermic process Endothermic reaction Endotherm
Platinum is a chemical element with symbol Pt and atomic number 78. It is a dense, ductile unreactive, silverish-white transition metal, its name is derived from the Spanish term platino, meaning "little silver". Platinum is a member of the platinum group of elements and group 10 of the periodic table of elements, it has six occurring isotopes. It is one of the rarer elements in Earth's crust, with an average abundance of 5 μg/kg, it occurs in some nickel and copper ores along with some native deposits in South Africa, which accounts for 80% of the world production. Because of its scarcity in Earth's crust, only a few hundred tonnes are produced annually, given its important uses, it is valuable and is a major precious metal commodity. Platinum is one of the least reactive metals, it has remarkable resistance to corrosion at high temperatures, is therefore considered a noble metal. Platinum is found chemically uncombined as native platinum; because it occurs in the alluvial sands of various rivers, it was first used by pre-Columbian South American natives to produce artifacts.
It was referenced in European writings as early as 16th century, but it was not until Antonio de Ulloa published a report on a new metal of Colombian origin in 1748 that it began to be investigated by scientists. Platinum is used in catalytic converters, laboratory equipment, electrical contacts and electrodes, platinum resistance thermometers, dentistry equipment, jewelry. Being a heavy metal, it leads to health problems upon exposure to its salts. Compounds containing platinum, such as cisplatin and carboplatin, are applied in chemotherapy against certain types of cancer; as of 2018, the value of platinum is $833.00 per ounce. Pure platinum is a lustrous and malleable, silver-white metal. Platinum is more ductile than gold, silver or copper, thus being the most ductile of pure metals, but it is less malleable than gold; the metal has excellent resistance to corrosion, is stable at high temperatures and has stable electrical properties. Platinum does oxidize, forming PtO2, at 500 °C, it reacts vigorously with fluorine at 500 °C to form platinum tetrafluoride.
It is attacked by chlorine, bromine and sulfur. Platinum is insoluble in hydrochloric and nitric acid, but dissolves in hot aqua regia, to form chloroplatinic acid, H2PtCl6, its physical characteristics and chemical stability make it useful for industrial applications. Its resistance to wear and tarnish is well suited to use in fine jewellery; the most common oxidation states of platinum are +2 and +4. The +1 and +3 oxidation states are less common, are stabilized by metal bonding in bimetallic species; as is expected, tetracoordinate platinum compounds tend to adopt 16-electron square planar geometries. Although elemental platinum is unreactive, it dissolves in hot aqua regia to give aqueous chloroplatinic acid: Pt + 4 HNO3 + 6 HCl → H2PtCl6 + 4 NO2 + 4 H2OAs a soft acid, platinum has a great affinity for sulfur, such as on dimethyl sulfoxide. In 2007, Gerhard Ertl won the Nobel Prize in Chemistry for determining the detailed molecular mechanisms of the catalytic oxidation of carbon monoxide over platinum.
Platinum has six occurring isotopes: 190Pt, 192Pt, 194Pt, 195Pt, 196Pt, 198Pt. The most abundant of these is 195 Pt, it is the only stable isotope with a non-zero spin. 190Pt is the least abundant at only 0.01%. Of the occurring isotopes, only 190Pt is unstable, though it decays with a half-life of 6.5×1011 years, causing an activity of 15 Bq/kg of natural platinum. 198 Pt can undergo alpha decay. Platinum has 31 synthetic isotopes ranging in atomic mass from 166 to 204, making the total number of known isotopes 39; the least stable of these is 166Pt, with a half-life of 300 µs, whereas the most stable is 193Pt with a half-life of 50 years. Most platinum isotopes decay by some combination of beta alpha decay. 188Pt, 191Pt, 193Pt decay by electron capture. 190Pt and 198Pt are predicted to have energetically favorable double beta decay paths. Platinum is an rare metal, occurring at a concentration of only 0.005 ppm in Earth's crust. It is sometimes mistaken for silver. Platinum is found chemically uncombined as native platinum and as alloy with the other platinum-group metals and iron mostly.
Most the native platinum is found in secondary deposits in alluvial deposits. The alluvial deposits used by pre-Columbian people in the Chocó Department, Colombia are still a source for platinum-group metals. Another large alluvial deposit is in the Ural Mountains, it is still mined. In nickel and copper deposits, platinum-group metals occur as sulfides, tellurides and arsenides, as end alloys with nickel or copper. Platinum arsenide, sperrylite, is a major source of platinum associated with nickel ores in the Sudbury Basin deposit in Ontario, Canada. At Platinum, about 17,000 kg was mined between 1927 and 1975; the mine ceased operations in 1990. The rare sulfide minera
Arsenic is a chemical element with symbol As and atomic number 33. Arsenic occurs in many minerals in combination with sulfur and metals, but as a pure elemental crystal. Arsenic is a metalloid, it has various allotropes, but only the gray form, which has a metallic appearance, is important to industry. The primary use of arsenic is in alloys of lead. Arsenic is a common n-type dopant in semiconductor electronic devices, the optoelectronic compound gallium arsenide is the second most used semiconductor after doped silicon. Arsenic and its compounds the trioxide, are used in the production of pesticides, treated wood products and insecticides; these applications are declining due to the toxicity of its compounds. A few species of bacteria are able to use arsenic compounds as respiratory metabolites. Trace quantities of arsenic are an essential dietary element in rats, goats and other species. A role in human metabolism is not known. However, arsenic poisoning occurs in multicellular life. Arsenic contamination of groundwater is a problem.
The United States' Environmental Protection Agency states that all forms of arsenic are a serious risk to human health. The United States' Agency for Toxic Substances and Disease Registry ranked arsenic as number 1 in its 2001 Priority List of Hazardous Substances at Superfund sites. Arsenic is classified as a Group-A carcinogen; the three most common arsenic allotropes are gray and black arsenic, with gray being the most common. Gray arsenic adopts a double-layered structure consisting of many interlocked, six-membered rings; because of weak bonding between the layers, gray arsenic is brittle and has a low Mohs hardness of 3.5. Nearest and next-nearest neighbors form a distorted octahedral complex, with the three atoms in the same double-layer being closer than the three atoms in the next; this close packing leads to a high density of 5.73 g/cm3. Gray arsenic becomes a semiconductor with a bandgap of 1.2 -- 1.4 eV if amorphized. Gray arsenic is the most stable form. Yellow arsenic is soft and waxy, somewhat similar to tetraphosphorus.
Both have four atoms arranged in a tetrahedral structure in which each atom is bound to each of the other three atoms by a single bond. This unstable allotrope, being molecular, is the most volatile, least dense, most toxic. Solid yellow arsenic is produced by rapid cooling of arsenic vapor, As4, it is transformed into gray arsenic by light. The yellow form has a density of 1.97 g/cm3. Black arsenic is similar in structure to black phosphorus. Black arsenic can be formed by cooling vapor at around 100–220 °C, it is brittle. It is a poor electrical conductor. Arsenic occurs in nature as a monoisotopic element, composed of 75As; as of 2003, at least 33 radioisotopes have been synthesized, ranging in atomic mass from 60 to 92. The most stable of these is 73As with a half-life of 80.30 days. All other isotopes have half-lives of under one day, with the exception of 71As, 72As, 74As, 76As, 77As. Isotopes that are lighter than the stable 75As tend to decay by β+ decay, those that are heavier tend to decay by β− decay, with some exceptions.
At least 10 nuclear isomers have been described, ranging in atomic mass from 66 to 84. The most stable of arsenic's isomers is 68mAs with a half-life of 111 seconds. Arsenic has a similar electronegativity and ionization energies to its lighter congener phosphorus and as such forms covalent molecules with most of the nonmetals. Though stable in dry air, arsenic forms a golden-bronze tarnish upon exposure to humidity which becomes a black surface layer; when heated in air, arsenic oxidizes to arsenic trioxide. This odor can be detected on striking arsenide minerals such as arsenopyrite with a hammer, it burns in oxygen to form arsenic trioxide and arsenic pentoxide, which have the same structure as the more well-known phosphorus compounds, in fluorine to give arsenic pentafluoride. Arsenic sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state at 887 K; the triple point is 3.63 MPa and 1,090 K. Arsenic makes arsenic acid with concentrated nitric acid, arsenous acid with dilute nitric acid, arsenic trioxide with concentrated sulfuric acid.
Arsenic reacts with metals to form arsenides, though these are not ionic compounds containing the As3− ion as the formation of such an anion would be endothermic and the group 1 arsenides have properties of intermetallic compounds. Like germanium and bromine, which like arsenic succeed the 3d transition series, arsenic is much less stable in the group oxidation state of +5 than its vertical neighbors phosphorus and antimony, hence arsenic pentoxide and arsenic acid are potent oxidizers. Compounds of arsenic resemble in some respects those of phosphorus which occupies the same group of the periodic table; the most common oxidation states for arsenic are: −3 in the arsenides, which are alloy-like intermetallic compounds, +3 in the arsenites, +5 in the arsenates and most organoarsenic compounds. Arsenic bonds to itself as seen in the square As3−4 ions in the mineral skutterudite. In the +3 oxidation state, arsenic is pyramidal owing to the i