In chemistry, bases are substances that, in aqueous solution, release hydroxide ions, are slippery to the touch, can taste bitter if an alkali, change the color of indicators, react with acids to form salts, promote certain chemical reactions, accept protons from any proton donor or contain or displaceable OH− ions. Examples of bases are the hydroxides of the alkaline earth metals; these particular substances produce hydroxide ions in aqueous solutions, are thus classified as Arrhenius bases. For a substance to be classified as an Arrhenius base, it must produce hydroxide ions in an aqueous solution. Arrhenius believed; this makes the Arrhenius model limited, as it cannot explain the basic properties of aqueous solutions of ammonia or its organic derivatives. There are bases that do not contain a hydroxide ion but react with water, resulting in an increase in the concentration of the hydroxide ion. An example of this is the reaction between water to produce ammonium and hydroxide. In this reaction ammonia is the base.
Ammonia and other bases similar to it have the ability to form a bond with a proton due to the unshared pair of electrons that they possess. In the more general Brønsted–Lowry acid–base theory, a base is a substance that can accept hydrogen cations —otherwise known as protons. In the Lewis model, a base is an electron pair donor. In water, by altering the autoionization equilibrium, bases yield solutions in which the hydrogen ion activity is lower than it is in pure water, i.e. the water has a pH higher than 7.0 at standard conditions. A soluble base is called an alkali if it releases OH − ions quantitatively. However, it is important to realize. Metal oxides and alkoxides are basic, conjugate bases of weak acids are weak bases. Bases can be thought of as the chemical opposite of acids. However, some strong acids are able to act as bases. Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium concentration in water, whereas bases reduce this concentration.
A reaction between an acid and a base is called neutralization. In a neutralization reaction, an aqueous solution of a base reacts with an aqueous solution of an acid to produce a solution of water and salt in which the salt separates into its component ions. If the aqueous solution is saturated with a given salt solute, any additional such salt precipitates out of the solution; the notion of a base as a concept in chemistry was first introduced by the French chemist Guillaume François Rouelle in 1754. He noted that acids, which at that time were volatile liquids, turned into solid salts only when combined with specific substances. Rouelle considered that such a substance serves as a "base" for the salt, giving the salt a "concrete or solid form". General properties of bases include: Concentrated or strong bases are caustic on organic matter and react violently with acidic substances. Aqueous solutions or molten bases dissociate in ions and conduct electricity. Reactions with indicators: bases turn red litmus paper blue, phenolphthalein pink, keep bromothymol blue in its natural colour of blue, turn methyl orange yellow.
The pH of a basic solution at standard conditions is greater than seven. Bases are bitter in taste; the following reaction represents the general reaction between a base and water to produce a conjugate acid and a conjugate base: B + H2O ⇌ BH+ + OH−The equilibrium constant, Kb, for this reaction can be found using the following general equation: Kb = /In this equation, both the base and the strong base compete with one another for the proton. As a result, bases that react with water have small equilibrium constant values; the base is weaker. Bases react with acids to neutralize each other at a fast rate both in alcohol; when dissolved in water, the strong base sodium hydroxide ionizes into hydroxide and sodium ions: NaOH → Na+ + OH−and in water the acid hydrogen chloride forms hydronium and chloride ions: HCl + H2O → H3O+ + Cl−When the two solutions are mixed, the H3O+ and OH− ions combine to form water molecules: H3O+ + OH− → 2 H2OIf equal quantities of NaOH and HCl are dissolved, the base and the acid neutralize leaving only NaCl table salt, in solution.
Weak bases, such as baking soda or egg white, should be used to neutralize any acid spills. Neutralizing acid spills with strong bases, such as sodium hydroxide or potassium hydroxide, can cause a violent exothermic reaction, the base itself can cause just as much damage as the original acid spill. Bases are compounds that can neutralize an amount of acids. Both sodium carbonate and ammonia are bases, although neither of these substances contains OH− groups. Both compounds accept H+ when dissolved in protic solvents such as water: Na2CO3 + H2O → 2 Na+ + HCO3− + OH− NH3 + H2O → NH4+ + OH−From this, a pH, or acidity, can be calculated for aqueous solutions of bases. Bases directly act as electron-pair donors themselves: CO32− + H+ → HCO3− NH3 + H+ → NH4+A base is defined as a molecule that has the ability to accept an electron pair bond by entering another atom's valence shell through its possession of one electron pair. There are a limited number of elements that have atoms with the ability to provide a molecule with basic properties
Palladium is a chemical element with symbol Pd and atomic number 46. It is a lustrous silvery-white metal discovered in 1803 by William Hyde Wollaston, he named it after the asteroid Pallas, itself named after the epithet of the Greek goddess Athena, acquired by her when she slew Pallas. Palladium, rhodium, ruthenium and osmium form a group of elements referred to as the platinum group metals; these have similar chemical properties, but palladium has the lowest melting point and is the least dense of them. More than half the supply of palladium and its congener platinum is used in catalytic converters, which convert as much as 90% of the harmful gases in automobile exhaust into less noxious substances. Palladium is used in electronics, medicine, hydrogen purification, chemical applications, groundwater treatment, jewelry. Palladium is a key component of fuel cells, which react hydrogen with oxygen to produce electricity and water. Ore deposits of palladium and other PGMs are rare; the most extensive deposits have been found in the norite belt of the Bushveld Igneous Complex covering the Transvaal Basin in South Africa.
Recycling is a source from scrapped catalytic converters. The numerous applications and limited supply sources result in considerable investment interest. Palladium belongs to group 10 in the periodic table, but the configuration in the outermost electrons are in accordance with Hund's rule. Electrons in the s-shell migrate to fill the d orbitals. Palladium is a soft silver-white metal, it has the lowest melting point of the platinum group metals. It is soft and ductile when annealed and is increased in strength and hardness when cold-worked. Palladium dissolves in concentrated nitric acid, in hot, concentrated sulfuric acid, when finely ground, in hydrochloric acid, it dissolves at room temperature in aqua regia. Palladium does not react with oxygen at standard temperature. Palladium heated to 800 °C will produce a layer of palladium oxide, it tarnishes in a moist atmosphere containing sulfur. Palladium films with defects produced by alpha particle bombardment at low temperature exhibit superconductivity having Tc=3.2 K.
Occurring palladium is composed of seven isotopes, six of which are stable. The most stable radioisotopes are 107Pd with a half-life of 6.5 million years, 103Pd with 17 days, 100Pd with 3.63 days. Eighteen other radioisotopes have been characterized with atomic weights ranging from 90.94948 u to 122.93426 u. These have half-lives of less than thirty minutes, except 101Pd, 109Pd, 112Pd. For isotopes with atomic mass unit values less than that of the most abundant stable isotope, 106Pd, the primary decay mode is electron capture with the primary decay product being rhodium; the primary mode of decay for those isotopes of Pd with atomic mass greater than 106 is beta decay with the primary product of this decay being silver. Radiogenic 107Ag is a decay product of 107Pd and was first discovered in 1978 in the Santa Clara meteorite of 1976; the discoverers suggest that the coalescence and differentiation of iron-cored small planets may have occurred 10 million years after a nucleosynthetic event. 107Pd versus Ag correlations observed in bodies, which have been melted since accretion of the solar system, must reflect the presence of short-lived nuclides in the early solar system.
Palladium compounds exist in the 0 and +2 oxidation state. Other less common states are recognized; the compounds of palladium are more similar to those of platinum than those of any other element. Palladium chloride is the principal starting material for other palladium compounds, it arises by the reaction of palladium with chlorine. It is used to prepare heterogeneous palladium catalysts such as palladium on barium sulfate, palladium on carbon, palladium chloride on carbon. Solutions of PdCl2 in nitric acid react with acetic acid to give palladium acetate a versatile reagent. PdCl2 reacts with ligands to give square planar complexes of the type PdCl2L2. One example of such complexes is the benzonitrile derivative PdX22. PdCl2 + 2 L → PdCl2L2. Palladium forms a range of zerovalent complexes with the formula PdL4, PdL3 and PdL2. For example, reduction of a mixture of PdCl22 and PPh3 gives tetrakispalladium: 2 PdCl22 + 4 PPh3 + 5 N2H4 → 2 Pd4 + N2 + 4 N2H5+Cl−Another major palladium complex, trisdipalladium, is prepared by reducing sodium tetrachloropalladate in the presence of dibenzylideneacetone.
Palladium, as well as palladium, are catalysts in coupling reactions, as has been recognized by the 2010 Nobel Prize in Chemistry to Richard F. Heck, Ei-ichi Negishi, Akira Suzuki; such reactions are practiced for the synthesis of fine chemicals. Prominent coupling reactions include the Heck, Sonogashira coupling, Stille reactions, the Kumada coupling. Palladium acetate, tetrakispalladium (Pd4, trisdipalladium serve either as catalysts or precatalysts. Although Pd compounds are comparatively rare, one example is sodium hexachloropalladate
Calcium is a chemical element with symbol Ca and atomic number 20. As an alkaline earth metal, calcium is a reactive metal that forms a dark oxide-nitride layer when exposed to air, its physical and chemical properties are most similar to its heavier homologues strontium and barium. It is the fifth most abundant element in Earth's crust and the third most abundant metal, after iron and aluminium; the most common calcium compound on Earth is calcium carbonate, found in limestone and the fossilised remnants of early sea life. The name derives from Latin calx "lime", obtained from heating limestone; some calcium compounds were known to the ancients, though their chemistry was unknown until the seventeenth century. Pure calcium was isolated in 1808 via electrolysis of its oxide by Humphry Davy, who named the element. Calcium compounds are used in many industries: in foods and pharmaceuticals for calcium supplementation, in the paper industry as bleaches, as components in cement and electrical insulators, in the manufacture of soaps.
On the other hand, the metal in pure form has few applications due to its high reactivity. Calcium is the fifth-most abundant element in the human body; as electrolytes, calcium ions play a vital role in the physiological and biochemical processes of organisms and cells: in signal transduction pathways where they act as a second messenger. Calcium ions outside cells are important for maintaining the potential difference across excitable cell membranes as well as proper bone formation. Calcium is a ductile silvery metal whose properties are similar to the heavier elements in its group, strontium and radium. A calcium atom has twenty electrons, arranged in the electron configuration 4s2. Like the other elements placed in group 2 of the periodic table, calcium has two valence electrons in the outermost s-orbital, which are easily lost in chemical reactions to form a dipositive ion with the stable electron configuration of a noble gas, in this case argon. Hence, calcium is always divalent in its compounds, which are ionic.
Hypothetical univalent salts of calcium would be stable with respect to their elements, but not to disproportionation to the divalent salts and calcium metal, because the enthalpy of formation of MX2 is much higher than those of the hypothetical MX. This occurs because of the much greater lattice energy afforded by the more charged Ca2+ cation compared to the hypothetical Ca+ cation. Calcium, strontium and radium are always considered to be alkaline earth metals. Beryllium and magnesium are different from the other members of the group in their physical and chemical behaviour: they behave more like aluminium and zinc and have some of the weaker metallic character of the post-transition metals, why the traditional definition of the term "alkaline earth metal" excludes them; this classification is obsolete in English-language sources, but is still used in other countries such as Japan. As a result, comparisons with strontium and barium are more germane to calcium chemistry than comparisons with magnesium.
Calcium metal melts at 842 °C and boils at 1494 °C. It crystallises in the face-centered cubic arrangement like strontium, its density of 1.55 g/cm3 is the lowest in its group. Calcium can be cut with a knife with effort. While calcium is a poorer conductor of electricity than copper or aluminium by volume, it is a better conductor by mass than both due to its low density. While calcium is infeasible as a conductor for most terrestrial applications as it reacts with atmospheric oxygen, its use as such in space has been considered; the chemistry of calcium is that of a typical heavy alkaline earth metal. For example, calcium spontaneously reacts with water more than magnesium and less than strontium to produce calcium hydroxide and hydrogen gas, it reacts with the oxygen and nitrogen in the air to form a mixture of calcium oxide and calcium nitride. When finely divided, it spontaneously burns in air to produce the nitride. In bulk, calcium is less reactive: it forms a hydration coating in moist air, but below 30% relative humidity it may be stored indefinitely at room temperature.
Besides the simple oxide CaO, the peroxide CaO2 can be made by direct oxidation of calcium metal under a high pressure of oxygen, there is some evidence for a yellow superoxide Ca2. Calcium hydroxide, Ca2, is a strong base, though it is not as strong as the hydroxides of strontium, barium or the alkali metals. All four dihalides of calcium are known. Calcium carbonate and calcium sulfate are abundant minerals. Like strontium and barium, as well as the alkali metals and the divalent lanthanides europium and ytterbium, calcium metal dissolves directly in liquid ammonia to give a dark blue solution. Due to the large size of the Ca2+ ion, high coordination numbers are common, up to 24 in some intermetallic compounds such as CaZn13. Calcium is complexed by oxygen chelates such as EDTA and polyphosphates, which are useful in an
Rhodium is a chemical element with symbol Rh and atomic number 45. It is a rare, silvery-white, corrosion-resistant, chemically inert transition metal, it is a member of the platinum group. It has only one occurring isotope, 103Rh. Occurring rhodium is found as the free metal, alloyed with similar metals, as a chemical compound in minerals such as bowieite and rhodplumsite, it is one of most valuable precious metals. Rhodium is found in platinum or nickel ores together with the other members of the platinum group metals, it was discovered in 1803 by William Hyde Wollaston in one such ore, named for the rose color of one of its chlorine compounds, produced after it reacted with the powerful acid mixture aqua regia. The element's major use is as one of the catalysts in the three-way catalytic converters in automobiles; because rhodium metal is inert against corrosion and most aggressive chemicals, because of its rarity, rhodium is alloyed with platinum or palladium and applied in high-temperature and corrosion-resistive coatings.
White gold is plated with a thin rhodium layer to improve its appearance while sterling silver is rhodium-plated for tarnish resistance. Rhodium detectors are used in nuclear reactors to measure the neutron flux level. Rhodium was discovered in 1803 by William Hyde Wollaston, soon after his discovery of palladium, he used crude platinum ore obtained from South America. His procedure involved dissolving the ore in aqua regia and neutralizing the acid with sodium hydroxide, he precipitated the platinum as ammonium chloroplatinate by adding ammonium chloride. Most other metals like copper, lead and rhodium were precipitated with zinc. Diluted nitric acid dissolved all but rhodium. Of these, palladium dissolved in aqua regia but rhodium did not, the rhodium was precipitated by the addition of sodium chloride as Na3·nH2O. After being washed with ethanol, the rose-red precipitate was reacted with zinc, which displaced the rhodium in the ionic compound and thereby released the rhodium as free metal. After the discovery, the rare element had only minor applications.
The first major application was electroplating for decorative uses and as corrosion-resistant coating. The introduction of the three-way catalytic converter by Volvo in 1976 increased the demand for rhodium; the previous catalytic converters used platinum or palladium, while the three-way catalytic converter used rhodium to reduce the amount of NOx in the exhaust. Rhodium is a hard, durable metal that has a high reflectance. Rhodium metal does not form an oxide when heated. Oxygen is absorbed from the atmosphere only at the melting point of rhodium, but is released on solidification. Rhodium has both lower density than platinum, it is not attacked by most acids: it is insoluble in nitric acid and dissolves in aqua regia. Rhodium belongs to group 9 of the periodic table, but the configuration of electrons in the outermost shells is atypical for the group; this anomaly is observed in the neighboring elements, niobium and palladium. The common oxidation state of rhodium is +3, but oxidation states from 0 to +6 are observed.
Unlike ruthenium and osmium, rhodium forms no volatile oxygen compounds. The known stable oxides include Rh2O3, RhO2, RhO2·xH2O, Na2RhO3, Sr3LiRhO6 and Sr3NaRhO6. Halogen compounds are known in nearly the full range of possible oxidation states. Rhodium chloride, rhodium fluoride, rhodium fluoride and rhodium fluoride are examples; the lower oxidation states are stable only in the presence of ligands. The best-known rhodium-halogen compound is the Wilkinson's catalyst chlorotrisrhodium; this catalyst is used in the hydrogenation of alkenes. Occurring rhodium is composed of only one isotope, 103Rh; the most stable radioisotopes are 101Rh with a half-life of 3.3 years, 102Rh with a half-life of 207 days, 102mRh with a half-life of 2.9 years, 99Rh with a half-life of 16.1 days. Twenty other radioisotopes have been characterized with atomic weights ranging from 92.926 u to 116.925 u. Most of these have half-lives shorter except 100Rh and 105Rh. Rhodium has numerous meta states, the most stable being 102mRh with a half-life of about 2.9 years and 101mRh with a half-life of 4.34 days.
In isotopes weighing less than 103, the primary decay mode is electron capture and the primary decay product is ruthenium. In isotopes greater than 103, the primary decay mode is beta emission and the primary product is palladium. Rhodium is one of the rarest elements in the Earth's crust, comprising an estimated 0.0002 parts per million. Its rarity affects its use in commercial applications; the industrial extraction of rhodium is complex because the ores are mixed with other metals such as palladium, silver and gold and there are few rhodium-bearing minerals. It is found in platinum ores and extracted as a white inert metal, difficult to fuse. Principal sources are located in South Africa. Although the quantity at Sudbury is small, the large amount of processed nickel ore makes rhodium recovery cost-effective; the main exporter of rhodium is South
A metal is a material that, when freshly prepared, polished, or fractured, shows a lustrous appearance, conducts electricity and heat well. Metals are malleable or ductile. A metal may be an alloy such as stainless steel. In physics, a metal is regarded as any substance capable of conducting electricity at a temperature of absolute zero. Many elements and compounds that are not classified as metals become metallic under high pressures. For example, the nonmetal iodine becomes a metal at a pressure of between 40 and 170 thousand times atmospheric pressure; some materials regarded as metals can become nonmetals. Sodium, for example, becomes a nonmetal at pressure of just under two million times atmospheric pressure. In chemistry, two elements that would otherwise qualify as brittle metals—arsenic and antimony—are instead recognised as metalloids, on account of their predominately non-metallic chemistry. Around 95 of the 118 elements in the periodic table are metals; the number is inexact as the boundaries between metals and metalloids fluctuate due to a lack of universally accepted definitions of the categories involved.
In astrophysics the term "metal" is cast more to refer to all chemical elements in a star that are heavier than the lightest two and helium, not just traditional metals. A star fuses lighter atoms hydrogen and helium, into heavier atoms over its lifetime. Used in that sense, the metallicity of an astronomical object is the proportion of its matter made up of the heavier chemical elements. Metals are present in many aspects of modern life; the strength and resilience of some metals has led to their frequent use in, for example, high-rise building and bridge construction, as well as most vehicles, many home appliances, tools and railroad tracks. Precious metals were used as coinage, but in the modern era, coinage metals have extended to at least 23 of the chemical elements; the history of metals is thought to begin with the use of copper about 11,000 years ago. Gold, iron and brass were in use before the first known appearance of bronze in the 5th millennium BCE. Subsequent developments include the production of early forms of steel.
Metals are lustrous, at least when freshly prepared, polished, or fractured. Sheets of metal thicker than a few micrometres appear opaque; the solid or liquid state of metals originates in the capacity of the metal atoms involved to lose their outer shell electrons. Broadly, the forces holding an individual atom’s outer shell electrons in place are weaker than the attractive forces on the same electrons arising from interactions between the atoms in the solid or liquid metal; the electrons involved become delocalised and the atomic structure of a metal can be visualised as a collection of atoms embedded in a cloud of mobile electrons. This type of interaction is called a metallic bond; the strength of metallic bonds for different elemental metals reaches a maximum around the center of the transition metal series, as these elements have large numbers of delocalized electrons. Although most elemental metals have higher densities than most nonmetals, there is a wide variation in their densities, lithium being the least dense and osmium the most dense.
Magnesium and titanium are light metals of significant commercial importance. Their respective densities of 1.7, 2.7 and 4.5 g/cm3 can be compared to those of the older structural metals, like iron at 7.9 and copper at 8.9 g/cm3. An iron ball would thus weigh about as much as three aluminium balls. Metals are malleable and ductile, deforming under stress without cleaving; the nondirectional nature of metallic bonding is thought to contribute to the ductility of most metallic solids. In contrast, in an ionic compound like table salt, when the planes of an ionic bond slide past one another, the resultant change in location shifts ions of the same charge into close proximity, resulting in the cleavage of the crystal; such a shift is not observed in a covalently bonded crystal, such as a diamond, where fracture and crystal fragmentation occurs. Reversible elastic deformation in metals can be described by Hooke's Law for restoring forces, where the stress is linearly proportional to the strain. Heat or forces larger than a metal's elastic limit may cause a permanent deformation, known as plastic deformation or plasticity.
An applied force may be a compressive force, or a shear, bending or torsion force. A temperature change may affect the movement or displacement of structural defects in the metal such as grain boundaries, point vacancies and screw dislocations, stacking faults and twins in both crystalline and non-crystalline metals. Internal slip and metal fatigue may ensue; the atoms of metallic substances are arranged in one of three common crystal structures, namely body-centered cubic, face-centered cubic, hexagonal close-packed. In bcc, each atom is positioned at the center of a cube of eight others. In fcc and hcp, each atom is surrounded by twelve others; some metals adopt different structures depending on the temperature. The
Fluorine is a chemical element with symbol F and atomic number 9. It is the lightest halogen and exists as a toxic pale yellow diatomic gas at standard conditions; as the most electronegative element, it is reactive, as it reacts with all other elements, except for helium and neon. Among the elements, fluorine ranks 24th in universal 13th in terrestrial abundance. Fluorite, the primary mineral source of fluorine which gave the element its name, was first described in 1529. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, several early experimenters died or sustained injuries from their attempts. Only in 1886 did French chemist Henri Moissan isolate elemental fluorine using low-temperature electrolysis, a process still employed for modern production. Industrial production of fluorine gas for uranium enrichment, its largest application, began during the Manhattan Project in World War II. Owing to the expense of refining pure fluorine, most commercial applications use fluorine compounds, with about half of mined fluorite used in steelmaking.
The rest of the fluorite is converted into corrosive hydrogen fluoride en route to various organic fluorides, or into cryolite, which plays a key role in aluminium refining. Molecules containing a Carbon–fluorine bond have high chemical and thermal stability. Pharmaceuticals such as atorvastatin and fluoxetine contain C-F bonds, the fluoride ion inhibits dental cavities, so finds use in toothpaste and water fluoridation. Global fluorochemical sales amount to more than US$15 billion a year. Fluorocarbon gases are greenhouse gases with global-warming potentials 100 to 20,000 times that of carbon dioxide. Organofluorine compounds persist in the environment due to the strength of the carbon–fluorine bond. Fluorine has no known metabolic role in mammals. Fluorine atoms have nine electrons, one fewer than neon, electron configuration 1s22s22p5: two electrons in a filled inner shell and seven in an outer shell requiring one more to be filled; the outer electrons are ineffective at nuclear shielding, experience a high effective nuclear charge of 9 − 2 = 7.
Fluorine's first ionization energy is third-highest among all elements, behind helium and neon, which complicates the removal of electrons from neutral fluorine atoms. It has a high electron affinity, second only to chlorine, tends to capture an electron to become isoelectronic with the noble gas neon. Fluorine atoms have a small covalent radius of around 60 picometers, similar to those of its period neighbors oxygen and neon; the bond energy of difluorine is much lower than that of either Cl2 or Br2 and similar to the cleaved peroxide bond. Conversely, bonds to other atoms are strong because of fluorine's high electronegativity. Unreactive substances like powdered steel, glass fragments, asbestos fibers react with cold fluorine gas. Reactions of elemental fluorine with metals require varying conditions. Alkali metals cause; some solid nonmetals react vigorously in liquid air temperature fluorine. Hydrogen sulfide and sulfur dioxide combine with fluorine, the latter sometimes explosively. Hydrogen, like some of the alkali metals, reacts explosively with fluorine.
Carbon, as lamp black, reacts at room temperature to yield fluoromethane. Graphite combines with fluorine above 400 °C to produce non-stoichiometric carbon monofluoride. Carbon dioxide and carbon monoxide react at or just above room temperature, whereas paraffins and other organic chemicals generate strong reactions: fully substituted haloalkanes such as carbon tetrachloride incombustible, may explode. Although nitrogen trifluoride is stable, nitrogen requires an electric discharge at elevated temperatures for reaction with fluorine to occur, due to the strong triple bond in elemental nitrogen. Oxygen does not combine with fluorine under ambient conditions, but can be made to react using electric discharge at low temperatures and pressures. Heavier halogens react with fluorine as does the noble gas radon. At room temperature, fluorine is a gas of diatomic molecules, pale yellow, it has a characteristic halogen-like biting odor detectable at 20 ppb. Fluorine condenses into a bright yellow liquid at −188 °C, a transition temperature similar to those of oxygen and nitrogen.
Fluorine has two solid forms, α- and β-fluorine. The latter crystallizes at −220 °C and is transparent and sof
Gold is a chemical element with symbol Au and atomic number 79, making it one of the higher atomic number elements that occur naturally. In its purest form, it is a bright reddish yellow, soft and ductile metal. Chemically, gold is a group 11 element, it is solid under standard conditions. Gold occurs in free elemental form, as nuggets or grains, in rocks, in veins, in alluvial deposits, it occurs in a solid solution series with the native element silver and naturally alloyed with copper and palladium. Less it occurs in minerals as gold compounds with tellurium. Gold is resistant to most acids, though it does dissolve in aqua regia, a mixture of nitric acid and hydrochloric acid, which forms a soluble tetrachloroaurate anion. Gold is insoluble in nitric acid, which dissolves silver and base metals, a property that has long been used to refine gold and to confirm the presence of gold in metallic objects, giving rise to the term acid test. Gold dissolves in alkaline solutions of cyanide, which are used in mining and electroplating.
Gold dissolves in mercury, forming amalgam alloys. A rare element, gold is a precious metal, used for coinage and other arts throughout recorded history. In the past, a gold standard was implemented as a monetary policy, but gold coins ceased to be minted as a circulating currency in the 1930s, the world gold standard was abandoned for a fiat currency system after 1971. A total of 186,700 tonnes of gold exists above ground, as of 2015; the world consumption of new gold produced is about 50% in jewelry, 40% in investments, 10% in industry. Gold's high malleability, resistance to corrosion and most other chemical reactions, conductivity of electricity have led to its continued use in corrosion resistant electrical connectors in all types of computerized devices. Gold is used in infrared shielding, colored-glass production, gold leafing, tooth restoration. Certain gold salts are still used as anti-inflammatories in medicine; as of 2017, the world's largest gold producer by far was China with 440 tonnes per year.
Gold is the most malleable of all metals. It can be drawn into a monoatomic wire, stretched about twice before it breaks; such nanowires distort via formation and migration of dislocations and crystal twins without noticeable hardening. A single gram of gold can be beaten into a sheet of 1 square meter, an avoirdupois ounce into 300 square feet. Gold leaf can be beaten thin enough to become semi-transparent; the transmitted light appears greenish blue, because gold reflects yellow and red. Such semi-transparent sheets strongly reflect infrared light, making them useful as infrared shields in visors of heat-resistant suits, in sun-visors for spacesuits. Gold is a good conductor of electricity. Gold has a density of 19.3 g/cm3 identical to that of tungsten at 19.25 g/cm3. By comparison, the density of lead is 11.34 g/cm3, that of the densest element, osmium, is 22.588±0.015 g/cm3. Whereas most metals are gray or silvery white, gold is reddish-yellow; this color is determined by the frequency of plasma oscillations among the metal's valence electrons, in the ultraviolet range for most metals but in the visible range for gold due to relativistic effects affecting the orbitals around gold atoms.
Similar effects impart a golden hue to metallic caesium. Common colored gold alloys include the distinctive eighteen-karat rose gold created by the addition of copper. Alloys containing palladium or nickel are important in commercial jewelry as these produce white gold alloys. Fourteen-karat gold-copper alloy is nearly identical in color to certain bronze alloys, both may be used to produce police and other badges. White gold alloys can be made with nickel. Fourteen- and eighteen-karat gold alloys with silver alone appear greenish-yellow and are referred to as green gold. Blue gold can be made by alloying with iron, purple gold can be made by alloying with aluminium. Less addition of manganese, aluminium and other elements can produce more unusual colors of gold for various applications. Colloidal gold, used by electron-microscopists, is red. Gold has only one stable isotope, 197Au, its only occurring isotope, so gold is both a mononuclidic and monoisotopic element. Thirty-six radioisotopes have been synthesized, ranging in atomic mass from 169 to 205.
The most stable of these is 195Au with a half-life of 186.1 days. The least stable is 171Au. Most of gold's radioisotopes with atomic masses below 197 decay by some combination of proton emission, α decay, β+ decay; the exceptions are 195Au, which decays by electron capture, 196Au, which decays most by electron capture with a minor β− decay path. All of gold's radioisotopes with atomic masses above 197 decay by β− decay. At least 32 nuclear isomers have been characterized, ranging in atomic mass from 170 to 200. Within that range, only 178Au, 180Au, 181Au, 182Au, 188Au do not have isomers. Gold's most stable isomer is 198m2Au with a half-life of 2.27 days. Gold's least stable isomer is 177m2Au with a half-life of only 7 ns. 184m1Au has three decay paths: β+ decay, isomeric