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Gödel's completeness theorem for first-order logic: every consistent set of first-order sentences has a completion. That is, every consistent set of first-order sentences can be extended to a maximal consistent set.
The compactness theorem: If is a set of first-order (or alternatively, zero-order) sentences such that every finite subset of has a model, then has a model. | Axiom of choice | Wikipedia | 82 | 840 | https://en.wikipedia.org/wiki/Axiom%20of%20choice | Mathematics | Discrete mathematics | null |
Possibly equivalent implications of AC
There are several historically important set-theoretic statements implied by AC whose equivalence to AC is open. Zermelo cited the partition principle, which was formulated before AC itself, as a justification for believing AC. In 1906, Russell declared PP to be equivalent, but whether the partition principle implies AC is the oldest open problem in set theory, and the equivalences of the other statements are similarly hard old open problems. In every known model of ZF where choice fails, these statements fail too, but it is unknown whether they can hold without choice.
Set theory
Partition principle: if there is a surjection from A to B, there is an injection from B to A. Equivalently, every partition P of a set S is less than or equal to S in size.
Converse Schröder–Bernstein theorem: if two sets have surjections to each other, they are equinumerous.
Weak partition principle: if there is an injection and a surjection from A to B, then A and B are equinumerous. Equivalently, a partition of a set S cannot be strictly larger than S. If WPP holds, this already implies the existence of a non-measurable set. Each of the previous three statements is implied by the preceding one, but it is unknown if any of these implications can be reversed.
There is no infinite decreasing sequence of cardinals. The equivalence was conjectured by Schoenflies in 1905.
Abstract algebra
Hahn embedding theorem: Every ordered abelian group G order-embeds as a subgroup of the additive group endowed with a lexicographical order, where Ω is the set of Archimedean equivalence classes of G. This equivalence was conjectured by Hahn in 1907.
Stronger forms of the negation of AC
If we abbreviate by BP the claim that every set of real numbers has the property of Baire, then BP is stronger than ¬AC, which asserts the nonexistence of any choice function on perhaps only a single set of nonempty sets. Strengthened negations may be compatible with weakened forms of AC. For example, ZF + DC + BP is consistent, if ZF is. | Axiom of choice | Wikipedia | 451 | 840 | https://en.wikipedia.org/wiki/Axiom%20of%20choice | Mathematics | Discrete mathematics | null |
It is also consistent with ZF + DC that every set of reals is Lebesgue measurable, but this consistency result, due to Robert M. Solovay, cannot be proved in ZFC itself, but requires a mild large cardinal assumption (the existence of an inaccessible cardinal). The much stronger axiom of determinacy, or AD, implies that every set of reals is Lebesgue measurable, has the property of Baire, and has the perfect set property (all three of these results are refuted by AC itself). ZF + DC + AD is consistent provided that a sufficiently strong large cardinal axiom is consistent (the existence of infinitely many Woodin cardinals).
Quine's system of axiomatic set theory, New Foundations (NF), takes its name from the title ("New Foundations for Mathematical Logic") of the 1937 article that introduced it. In the NF axiomatic system, the axiom of choice can be disproved.
Statements implying the negation of AC
There are models of Zermelo-Fraenkel set theory in which the axiom of choice is false. We shall abbreviate "Zermelo-Fraenkel set theory plus the negation of the axiom of choice" by ZF¬C. For certain models of ZF¬C, it is possible to validate the negation of some standard ZFC theorems. As any model of ZF¬C is also a model of ZF, it is the case that for each of the following statements, there exists a model of ZF in which that statement is true. | Axiom of choice | Wikipedia | 337 | 840 | https://en.wikipedia.org/wiki/Axiom%20of%20choice | Mathematics | Discrete mathematics | null |
The negation of the weak partition principle: There is a set that can be partitioned into strictly more equivalence classes than the original set has elements, and a function whose domain is strictly smaller than its range. In fact, this is the case in all known models.
There is a function f from the real numbers to the real numbers such that f is not continuous at a, but f is sequentially continuous at a, i.e., for any sequence {xn} converging to a, limn f(xn)=f(a).
There is an infinite set of real numbers without a countably infinite subset.
The real numbers are a countable union of countable sets. This does not imply that the real numbers are countable: As pointed out above, to show that a countable union of countable sets is itself countable requires the Axiom of countable choice.
There is a field with no algebraic closure.
In all models of ZF¬C there is a vector space with no basis.
There is a vector space with two bases of different cardinalities.
There is a free complete Boolean algebra on countably many generators.
There is a set that cannot be linearly ordered.
There exists a model of ZF¬C in which every set in Rn is measurable. Thus it is possible to exclude counterintuitive results like the Banach–Tarski paradox which are provable in ZFC. Furthermore, this is possible whilst assuming the Axiom of dependent choice, which is weaker than AC but sufficient to develop most of real analysis.
In all models of ZF¬C, the generalized continuum hypothesis does not hold.
For proofs, see .
Additionally, by imposing definability conditions on sets (in the sense of descriptive set theory) one can often prove restricted versions of the axiom of choice from axioms incompatible with general choice. This appears, for example, in the Moschovakis coding lemma.
Axiom of choice in type theory | Axiom of choice | Wikipedia | 408 | 840 | https://en.wikipedia.org/wiki/Axiom%20of%20choice | Mathematics | Discrete mathematics | null |
In type theory, a different kind of statement is known as the axiom of choice. This form begins with two types, σ and τ, and a relation R between objects of type σ and objects of type τ. The axiom of choice states that if for each x of type σ there exists a y of type τ such that R(x,y), then there is a function f from objects of type σ to objects of type τ such that R(x,f(x)) holds for all x of type σ:
Unlike in set theory, the axiom of choice in type theory is typically stated as an axiom scheme, in which R varies over all formulas or over all formulas of a particular logical form. | Axiom of choice | Wikipedia | 148 | 840 | https://en.wikipedia.org/wiki/Axiom%20of%20choice | Mathematics | Discrete mathematics | null |
An aircraft (: aircraft) is a vehicle that is able to fly by gaining support from the air. It counters the force of gravity by using either static lift or the dynamic lift of an airfoil, or, in a few cases, direct downward thrust from its engines. Common examples of aircraft include airplanes, rotorcraft, helicopters, airships (including blimps), gliders, paramotors, and hot air balloons. Part 1 (Definitions and Abbreviations) of Subchapter A of Chapter I of Title 14 of the U. S. Code of Federal Regulations states that aircraft "means a device that is used or intended to be used for flight in the air."
The human activity that surrounds aircraft is called aviation. The science of aviation, including designing and building aircraft, is called aeronautics. Crewed aircraft are flown by an onboard pilot, whereas unmanned aerial vehicles may be remotely controlled or self-controlled by onboard computers. Aircraft may be classified by different criteria, such as lift type, aircraft propulsion (if any), usage and others.
History
Flying model craft and stories of manned flight go back many centuries; however, the first manned ascent — and safe descent — in modern times took place by larger hot-air balloons developed in the 18th century. Each of the two World Wars led to great technical advances. Consequently, the history of aircraft can be divided into five eras:
Pioneers of flight, from the earliest experiments to 1914
First World War, 1914 to 1918
Aviation in the interwar period, 1918 to 1939
Second World War, 1939 to 1945
Postwar era, also called the Jet Age, 1945 to the present day
Methods of lift
Lighter-than-air
Lighter-than-air aircraft or aerostats use buoyancy to float in the air in much the same way that ships float on the water. They are characterized by one or more large cells or canopies, filled with a lifting gas such as helium, hydrogen or hot air, which is less dense than the surrounding air. When the weight of the lifting gas is added to the weight of the aircraft itself, it is same or less than the mass of the air that the craft displaces.
Small hot-air balloons, called sky lanterns, were first invented in ancient China prior to the 3rd century BC and used primarily in cultural celebrations, and were only the second type of aircraft to fly, the first being kites, which were also first invented in ancient China over two thousand years ago (see Han Dynasty). | Aircraft | Wikipedia | 507 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
A balloon was originally any aerostat, while the term airship was used for large, powered aircraft designs — usually fixed-wing. In 1919, Frederick Handley Page was reported as referring to "ships of the air," with smaller passenger types as "Air yachts." In the 1930s, large intercontinental flying boats were also sometimes referred to as "ships of the air" or "flying-ships". — though none had yet been built. The advent of powered balloons, called dirigible balloons, and later of rigid hulls allowing a great increase in size, began to change the way these words were used. Huge powered aerostats, characterized by a rigid outer framework and separate aerodynamic skin surrounding the gas bags, were produced, the Zeppelins being the largest and most famous. There were still no fixed-wing aircraft or non-rigid balloons large enough to be called airships, so "airship" came to be synonymous with these aircraft. Then several accidents, such as the Hindenburg disaster in 1937, led to the demise of these airships. Nowadays a "balloon" is an unpowered aerostat and an "airship" is a powered one.
A powered, steerable aerostat is called a dirigible. Sometimes this term is applied only to non-rigid balloons, and sometimes dirigible balloon is regarded as the definition of an airship (which may then be rigid or non-rigid). Non-rigid dirigibles are characterized by a moderately aerodynamic gasbag with stabilizing fins at the back. These soon became known as blimps. During World War II, this shape was widely adopted for tethered balloons; in windy weather, this both reduces the strain on the tether and stabilizes the balloon. The nickname blimp was adopted along with the shape. In modern times, any small dirigible or airship is called a blimp, though a blimp may be unpowered as well as powered.
Heavier-than-air
Heavier-than-air aircraft or aerodynes are denser than air and thus must find some way to obtain enough lift that can overcome the aircraft's weight. There are two ways to produce dynamic upthrust — aerodynamic lift by having air flowing past an aerofoil (such dynamic interaction of aerofoils with air is the origin of the term "aerodyne"), or powered lift in the form of reactional lift from downward engine thrust. | Aircraft | Wikipedia | 502 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
Aerodynamic lift involving wings is the most common, and can be achieved via two methods. Fixed-wing aircraft (airplanes and gliders) achieve airflow past the wings by having the entire aircraft moving forward through the air, while rotorcraft (helicopters and autogyros) do so by having mobile, elongated wings spinning rapidly around a mast in an assembly known as the rotor. As aerofoils, there must be air flowing over the wing to create pressure difference between above and below, thus generating upward lift over the entire wetted area of the wing. A flexible wing is a wing made of fabric or thin sheet material, often stretched over a rigid frame, similar to the flight membranes on many flying and gliding animals. A kite is tethered to the ground and relies on the speed of the wind over its wings, which may be flexible or rigid, fixed, or rotary.
With powered lift, the aircraft directs its engine thrust vertically downward. V/STOL aircraft, such as the Harrier jump jet and Lockheed Martin F-35B take off and land vertically using powered lift and transfer to aerodynamic lift in steady flight.
A pure rocket is not usually regarded as an aerodyne because its flight does not depend on interaction with the air at all (and thus can even fly in the vacuum of outer space); however, many aerodynamic lift vehicles have been powered or assisted by rocket motors. Rocket-powered missiles that obtain aerodynamic lift at very high speed due to airflow over their bodies are a marginal case.
Fixed-wing
The forerunner of the fixed-wing aircraft is the kite. Whereas a fixed-wing aircraft relies on its forward speed to create airflow over the wings, a kite is tethered to the ground and relies on the wind blowing over its wings to provide lift. Kites were the first kind of aircraft to fly and were invented in China around 500 BC. Much aerodynamic research was done with kites before test aircraft, wind tunnels, and computer modelling programs became available.
The first heavier-than-air craft capable of controlled free-flight were gliders. A glider designed by George Cayley carried out the first true manned, controlled flight in 1853. The first powered and controllable fixed-wing aircraft (the airplane or aeroplane) was invented by Wilbur and Orville Wright. | Aircraft | Wikipedia | 467 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
Besides the method of propulsion (if any), fixed-wing aircraft are in general characterized by their wing configuration. The most important wing characteristics are:
Number of wings – Monoplane, biplane, triplane, or multiplane.
Wing support – Braced or cantilever, rigid or flexible.
Wing planform – including aspect ratio, angle of sweep, and any variations along the span (including the important class of delta wings).
Location of the horizontal stabilizer, if any.
Dihedral angle – positive, zero, or negative (anhedral).
A variable geometry aircraft can change its wing configuration during flight.
A flying wing has no fuselage, though it may have small blisters or pods. The opposite of this is a lifting body, which has no wings, though it may have small stabilizing and control surfaces.
Wing-in-ground-effect vehicles are generally not considered aircraft. They "fly" efficiently close to the surface of the ground or water, like conventional aircraft during takeoff. An example is the Russian ekranoplan nicknamed the "Caspian Sea Monster". Man-powered aircraft also rely on ground effect to remain airborne with minimal pilot power, but this is only because they are so underpowered—in fact, the airframe is capable of flying higher.
Rotorcraft
Rotorcraft, or rotary-wing aircraft, use a spinning rotor with aerofoil cross-section blades (a rotary wing) to provide lift. Types include helicopters, autogyros, and various hybrids such as gyrodynes and compound rotorcraft.
Helicopters have a rotor turned by an engine-driven shaft. The rotor pushes air downward to create lift. By tilting the rotor forward, the downward flow is tilted backward, producing thrust for forward flight. Some helicopters have more than one rotor and a few have rotors turned by gas jets at the tips. Some have a tail rotor to counteract the rotation of the main rotor, and to aid directional control.
Autogyros have unpowered rotors, with a separate power plant to provide thrust. The rotor is tilted backward. As the autogyro moves forward, air blows upward across the rotor, making it spin. This spinning increases the speed of airflow over the rotor, to provide lift. Rotor kites are unpowered autogyros, which are towed to give them forward speed or tethered to a static anchor in high-wind for kited flight. | Aircraft | Wikipedia | 490 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
Compound rotorcraft have wings that provide some or all of the lift in forward flight. They are nowadays classified as powered lift types and not as rotorcraft. Tiltrotor aircraft (such as the Bell Boeing V-22 Osprey), tiltwing, tail-sitter, and coleopter aircraft have their rotors/propellers horizontal for vertical flight and vertical for forward flight.
Other methods of lift
A lifting body is an aircraft body shaped to produce lift. If there are any wings, they are too small to provide significant lift and are used only for stability and control. Lifting bodies are not efficient: they suffer from high drag, and must also travel at high speed to generate enough lift to fly. Many of the research prototypes, such as the Martin Marietta X-24, which led up to the Space Shuttle, were lifting bodies, though the Space Shuttle is not, and some supersonic missiles obtain lift from the airflow over a tubular body.
Powered lift types rely on engine-derived lift for vertical takeoff and landing (VTOL). Most types transition to fixed-wing lift for horizontal flight. Classes of powered lift types include VTOL jet aircraft (such as the Harrier jump jet) and tiltrotors, such as the Bell Boeing V-22 Osprey, among others. A few experimental designs rely entirely on engine thrust to provide lift throughout the whole flight, including personal fan-lift hover platforms and jetpacks. VTOL research designs include the Rolls-Royce Thrust Measuring Rig.
Some rotor wings employ horizontal-axis wings, in which airflow across a spinning rotor generates lift. The Flettner airplane uses a rotating cylinder, obtaining lift from the Magnus effect. The FanWing uses a cross-flow fan, while the mechanically more complex cyclogyro comprises multiple wings which rotate together around a central axis.
The ornithopter obtains thrust by flapping its wings.
Size and speed extremes
Size
The smallest aircraft are toys/recreational items, and nano aircraft.
The largest aircraft by dimensions and volume (as of 2016) is the long British Airlander 10, a hybrid blimp, with helicopter and fixed-wing features, and reportedly capable of speeds up to , and an airborne endurance of two weeks with a payload of up to . | Aircraft | Wikipedia | 463 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
The largest aircraft by weight and largest regular fixed-wing aircraft ever built, , was the Antonov An-225 Mriya. That Soviet-built (Ukrainian SSR) six-engine transport of the 1980s was long, with an wingspan. It holds the world payload record, after transporting of goods, and has flown loads commercially. With a maximum loaded weight of , it was also the heaviest aircraft built to date. It could cruise at . The aircraft was destroyed during the Russo-Ukrainian War.
The largest military airplanes are the Ukrainian Antonov An-124 Ruslan (world's second-largest airplane, also used as a civilian transport), and American Lockheed C-5 Galaxy transport, weighing, loaded, over . The 8-engine, piston/propeller Hughes H-4 Hercules "Spruce Goose" — an American World War II wooden flying boat transport with a greater wingspan (94m/260 ft) than any current aircraft and a tail height equal to the tallest (Airbus A380-800 at 24.1m/78 ft) — flew only one short hop in the late 1940s and never flew out of ground effect.
The largest civilian airplanes, apart from the above-noted An-225 and An-124, are the Airbus Beluga cargo transport derivative of the Airbus A300 jet airliner, the Boeing Dreamlifter cargo transport derivative of the Boeing 747 jet airliner/transport (the 747-200B was, at its creation in the 1960s, the heaviest aircraft ever built, with a maximum weight of over ), and the double-decker Airbus A380 "super-jumbo" jet airliner (the world's largest passenger airliner).
Speeds
The fastest fixed-wing aircraft and fastest glider, is the Space Shuttle, which re-entered the atmosphere at nearly Mach 25 or
The fastest recorded powered aircraft flight and fastest recorded aircraft flight of an air-breathing powered aircraft was of the NASA X-43A Pegasus, a scramjet-powered, hypersonic, lifting body experimental research aircraft, at Mach 9.68 or on 16 November 2004.
Prior to the X-43A, the fastest recorded powered airplane flight, and still the record for the fastest manned powered airplane, was the North American X-15, rocket-powered airplane at Mach 6.7 or 7,274 km/h (4,520 mph) on 3 October 1967. | Aircraft | Wikipedia | 484 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
The fastest manned, air-breathing powered airplane is the Lockheed SR-71 Blackbird, a U.S. reconnaissance jet fixed-wing aircraft, having reached on 28 July 1976.
Propulsion
Unpowered aircraft
Gliders are heavier-than-air aircraft that do not employ propulsion once airborne. Take-off may be by launching forward and downward from a high location, or by pulling into the air on a tow-line, either by a ground-based winch or vehicle, or by a powered "tug" aircraft. For a glider to maintain its forward air speed and lift, it must descend in relation to the air (but not necessarily in relation to the ground). Many gliders can "soar", i.e., gain height from updrafts such as thermal currents. The first practical, controllable example was designed and built by the British scientist and pioneer George Cayley, whom many recognise as the first aeronautical engineer. Common examples of gliders are sailplanes, hang gliders and paragliders.
Balloons drift with the wind, though normally the pilot can control the altitude, either by heating the air or by releasing ballast, giving some directional control (since the wind direction changes with altitude). A wing-shaped hybrid balloon can glide directionally when rising or falling; but a spherically shaped balloon does not have such directional control.
Kites are aircraft that are tethered to the ground or other object (fixed or mobile) that maintains tension in the tether or kite line; they rely on virtual or real wind blowing over and under them to generate lift and drag. Kytoons are balloon-kite hybrids that are shaped and tethered to obtain kiting deflections, and can be lighter-than-air, neutrally buoyant, or heavier-than-air.
Powered aircraft
Powered aircraft have one or more onboard sources of mechanical power, typically aircraft engines although rubber and manpower have also been used. Most aircraft engines are either lightweight reciprocating engines or gas turbines. Engine fuel is stored in tanks, usually in the wings but larger aircraft also have additional fuel tanks in the fuselage.
Propeller aircraft
Propeller aircraft use one or more propellers (airscrews) to create thrust in a forward direction. The propeller is usually mounted in front of the power source in tractor configuration but can be mounted behind in pusher configuration. Variations of propeller layout include contra-rotating propellers and ducted fans. | Aircraft | Wikipedia | 496 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
Many kinds of power plant have been used to drive propellers. Early airships used man power or steam engines. The more practical internal combustion piston engine was used for virtually all fixed-wing aircraft until World War II and is still used in many smaller aircraft. Some types use turbine engines to drive a propeller in the form of a turboprop or propfan. Human-powered flight has been achieved, but has not become a practical means of transport. Unmanned aircraft and models have also used power sources such as electric motors and rubber bands.
Jet aircraft
Jet aircraft use airbreathing jet engines, which take in air, burn fuel with it in a combustion chamber, and accelerate the exhaust rearwards to provide thrust.
Different jet engine configurations include the turbojet and turbofan, sometimes with the addition of an afterburner. Those with no rotating turbomachinery include the pulsejet and ramjet. These mechanically simple engines produce no thrust when stationary, so the aircraft must be launched to flying speed using a catapult, like the V-1 flying bomb, or a rocket, for example. Other engine types include the motorjet and the dual-cycle Pratt & Whitney J58.
Compared to engines using propellers, jet engines can provide much higher thrust, higher speeds and, above about , greater efficiency. They are also much more fuel-efficient than rockets. As a consequence nearly all large, high-speed or high-altitude aircraft use jet engines.
Rotorcraft
Some rotorcraft, such as helicopters, have a powered rotary wing or rotor, where the rotor disc can be angled slightly forward so that a proportion of its lift is directed forwards. The rotor may, like a propeller, be powered by a variety of methods such as a piston engine or turbine. Experiments have also used jet nozzles at the rotor blade tips.
Other types of powered aircraft
Rocket-powered aircraft have occasionally been experimented with, and the Messerschmitt Me 163 Komet fighter even saw action in the Second World War. Since then, they have been restricted to research aircraft, such as the North American X-15, which traveled up into space where air-breathing engines cannot work (rockets carry their own oxidant). Rockets have more often been used as a supplement to the main power plant, typically for the rocket-assisted take off of heavily loaded aircraft, but also to provide high-speed dash capability in some hybrid designs such as the Saunders-Roe SR.53.
The ornithopter obtains thrust by flapping its wings. | Aircraft | Wikipedia | 512 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
Design and construction
Aircraft are designed according to many factors such as customer and manufacturer demand, safety protocols and physical and economic constraints. For many types of aircraft the design process is regulated by national airworthiness authorities.
The key parts of an aircraft are generally divided into three categories:
The structure ("airframe") comprises the main load-bearing elements and associated equipment, as well as flight controls.
The propulsion system ("powerplant") (if it is powered) comprises the power source and associated equipment, as described above.
The avionics comprise the electrical and electronic control, navigation and communication systems.
Structure
The approach to structural design varies widely between different types of aircraft. Some, such as paragliders, comprise only flexible materials that act in tension and rely on aerodynamic pressure to hold their shape. A balloon similarly relies on internal gas pressure, but may have a rigid basket or gondola slung below it to carry its payload. Early aircraft, including airships, often employed flexible doped aircraft fabric covering to give a reasonably smooth aeroshell stretched over a rigid frame. Later aircraft employed semi-monocoque techniques, where the skin of the aircraft is stiff enough to share much of the flight loads. In a true monocoque design there is no internal structure left.
The key structural parts of an aircraft depend on what type it is.
Aerostats
Lighter-than-air types are characterised by one or more gasbags, typically with a supporting structure of flexible cables or a rigid framework called its hull. Other elements such as engines or a gondola may also be attached to the supporting structure.
Aerodynes
Heavier-than-air types are characterised by one or more wings and a central fuselage. The fuselage typically also carries a tail or empennage for stability and control, and an undercarriage for takeoff and landing. Engines may be located on the fuselage or wings. On a fixed-wing aircraft the wings are rigidly attached to the fuselage, while on a rotorcraft the wings are attached to a rotating vertical shaft. Smaller designs sometimes use flexible materials for part or all of the structure, held in place either by a rigid frame or by air pressure. The fixed parts of the structure comprise the airframe.
Power | Aircraft | Wikipedia | 452 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
The source of motive power for an aircraft is normally called the powerplant, and includes engine or motor, propeller or rotor, (if any), jet nozzles and thrust reversers (if any), and accessories essential to the functioning of the engine or motor (e.g.: starter, ignition system, intake system, exhaust system, fuel system, lubrication system, engine cooling system, and engine controls).
Powered aircraft are typically powered by internal combustion engines (piston or turbine) burning fossil fuels—typically gasoline (avgas) or jet fuel. A very few are powered by rocket power, ramjet propulsion, or by electric motors, or by internal combustion engines of other types, or using other fuels. A very few have been powered, for short flights, by human muscle energy (e.g.: Gossamer Condor).
Avionics
The avionics comprise any electronic aircraft flight control systems and related equipment, including electronic cockpit instrumentation, navigation, radar, monitoring, and communications systems.
Flight characteristics
Flight envelope
The flight envelope of an aircraft refers to its approved design capabilities in terms of airspeed, load factor and altitude. The term can also refer to other assessments of aircraft performance such as maneuverability. When an aircraft is abused, for instance by diving it at too-high a speed, it is said to be flown outside the envelope, something considered foolhardy since it has been taken beyond the design limits which have been established by the manufacturer. Going beyond the envelope may have a known outcome such as flutter or entry to a non-recoverable spin (possible reasons for the boundary).
Range
The range is the distance an aircraft can fly between takeoff and landing, as limited by the time it can remain airborne.
For a powered aircraft the time limit is determined by the fuel load and rate of consumption.
For an unpowered aircraft, the maximum flight time is limited by factors such as weather conditions and pilot endurance. Many aircraft types are restricted to daylight hours, while balloons are limited by their supply of lifting gas. The range can be seen as the average ground speed multiplied by the maximum time in the air.
The Airbus A350-900ULR is among the longest range airliners.
Flight dynamics | Aircraft | Wikipedia | 452 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
Flight dynamics is the science of air vehicle orientation and control in three dimensions. The three critical flight dynamics parameters are the angles of rotation around three axes which pass through the vehicle's center of gravity, known as pitch, roll, and yaw.
Roll is a rotation about the longitudinal axis (equivalent to the rolling or heeling of a ship) giving an up-down movement of the wing tips measured by the roll or bank angle.
Pitch is a rotation about the sideways horizontal axis giving an up-down movement of the aircraft nose measured by the angle of attack.
Yaw is a rotation about the vertical axis giving a side-to-side movement of the nose known as sideslip.
Flight dynamics is concerned with the stability and control of an aircraft's rotation about each of these axes.
Stability
An aircraft that is unstable tends to diverge from its intended flight path and so is difficult to fly. A very stable aircraft tends to stay on its flight path and is difficult to maneuver. Therefore, it is important for any design to achieve the desired degree of stability. Since the widespread use of digital computers, it is increasingly common for designs to be inherently unstable and rely on computerised control systems to provide artificial stability.
A fixed wing is typically unstable in pitch, roll, and yaw. Pitch and yaw stabilities of conventional fixed wing designs require horizontal and vertical stabilisers, which act similarly to the feathers on an arrow. These stabilizing surfaces allow equilibrium of aerodynamic forces and to stabilise the flight dynamics of pitch and yaw. They are usually mounted on the tail section (empennage), although in the canard layout, the main aft wing replaces the canard foreplane as pitch stabilizer. Tandem wing and tailless aircraft rely on the same general rule to achieve stability, the aft surface being the stabilising one.
A rotary wing is typically unstable in yaw, requiring a vertical stabiliser.
A balloon is typically very stable in pitch and roll due to the way the payload is slung underneath the center of lift.
Control
Flight control surfaces enable the pilot to control an aircraft's flight attitude and are usually part of the wing or mounted on, or integral with, the associated stabilizing surface. Their development was a critical advance in the history of aircraft, which had until that point been uncontrollable in flight. | Aircraft | Wikipedia | 480 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
Aerospace engineers develop control systems for a vehicle's orientation (attitude) about its center of mass. The control systems include actuators, which exert forces in various directions, and generate rotational forces or moments about the aerodynamic center of the aircraft, and thus rotate the aircraft in pitch, roll, or yaw. For example, a pitching moment is a vertical force applied at a distance forward or aft from the aerodynamic center of the aircraft, causing the aircraft to pitch up or down. Control systems are also sometimes used to increase or decrease drag, for example to slow the aircraft to a safe speed for landing.
The two main aerodynamic forces acting on any aircraft are lift supporting it in the air and drag opposing its motion. Control surfaces or other techniques may also be used to affect these forces directly, without inducing any rotation.
Environmental impact
Aircraft permit long distance, high speed travel and may be a more fuel efficient mode of transportation in some circumstances. Aircraft have environmental and climate impacts beyond fuel efficiency considerations, however. They are also relatively noisy compared to other forms of travel and high altitude aircraft generate contrails, which experimental evidence suggests may alter weather patterns.
Uses for aircraft
Aircraft are produced in several different types optimized for various uses; military aircraft, which includes not just combat types but many types of supporting aircraft, and civil aircraft, which include all non-military types, experimental and model.
Military
A military aircraft is any aircraft that is operated by a legal or insurrectionary armed service of any type. Military aircraft can be either combat or non-combat:
Combat aircraft are aircraft designed to destroy enemy equipment using its own armament. Combat aircraft divide broadly into fighters and bombers, with several in-between types, such as fighter-bombers and attack aircraft, including attack helicopters.
Non-combat aircraft are not designed for combat as their primary function, but may carry weapons for self-defense. Non-combat roles include search and rescue, reconnaissance, observation, transport, training, and aerial refueling. These aircraft are often variants of civil aircraft.
Most military aircraft are powered heavier-than-air types. Other types, such as gliders and balloons, have also been used as military aircraft; for example, balloons were used for observation during the American Civil War and World War I, and military gliders were used during World War II to land troops.
Civil
Civil aircraft divide into commercial and general types, however there are some overlaps. | Aircraft | Wikipedia | 487 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
Commercial aircraft include types designed for scheduled and charter airline flights, carrying passengers, mail and other cargo. The larger passenger-carrying types are the airliners, the largest of which are wide-body aircraft. Some of the smaller types are also used in general aviation, and some of the larger types are used as VIP aircraft.
General aviation is a catch-all covering other kinds of private (where the pilot is not paid for time or expenses) and commercial use, and involving a wide range of aircraft types such as business jets (bizjets), trainers, homebuilt, gliders, warbirds and hot air balloons to name a few. The vast majority of aircraft today are general aviation types.
Experimental
An experimental aircraft is one that has not been fully proven in flight, or that carries a Special Airworthiness Certificate, called an Experimental Certificate in United States parlance. This often implies that the aircraft is testing new aerospace technologies, though the term also refers to amateur-built and kit-built aircraft, many of which are based on proven designs.
Model
A model aircraft is a small unmanned type made to fly for fun, for static display, for aerodynamic research or for other purposes. A scale model is a replica of some larger design. | Aircraft | Wikipedia | 253 | 849 | https://en.wikipedia.org/wiki/Aircraft | Technology | Transportation | null |
Argon is a chemical element; it has symbol Ar and atomic number 18. It is in group 18 of the periodic table and is a noble gas. Argon is the third most abundant gas in Earth's atmosphere, at 0.934% (9340 ppmv). It is more than twice as abundant as water vapor (which averages about 4000 ppmv, but varies greatly), 23 times as abundant as carbon dioxide (400 ppmv), and more than 500 times as abundant as neon (18 ppmv). Argon is the most abundant noble gas in Earth's crust, comprising 0.00015% of the crust.
Nearly all argon in Earth's atmosphere is radiogenic argon-40, derived from the decay of potassium-40 in Earth's crust. In the universe, argon-36 is by far the most common argon isotope, as it is the most easily produced by stellar nucleosynthesis in supernovas.
The name "argon" is derived from the Greek word , neuter singular form of meaning 'lazy' or 'inactive', as a reference to the fact that the element undergoes almost no chemical reactions. The complete octet (eight electrons) 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 extracted industrially by the fractional distillation of liquid air. It is mostly used as an inert shielding gas in welding and other high-temperature industrial processes where ordinarily unreactive substances become reactive; for example, an argon atmosphere is used in graphite electric furnaces to prevent the graphite from burning. It is also used in incandescent and fluorescent lighting, and other gas-discharge tubes. It makes a distinctive blue-green gas laser. It is also used in fluorescent glow starters.
Characteristics
Argon has approximately the same solubility in water as oxygen and is 2.5 times more soluble in water than nitrogen. Argon is colorless, odorless, nonflammable and nontoxic as a solid, liquid or gas. Argon is chemically inert under most conditions and forms no confirmed stable compounds at room temperature. | Argon | Wikipedia | 476 | 896 | https://en.wikipedia.org/wiki/Argon | Physical sciences | Chemical elements_2 | null |
Although argon is a noble gas, it can form some compounds under various extreme conditions. Argon fluorohydride (HArF), a compound of argon with fluorine and hydrogen that is stable below , 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 , and 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.
History
Argon (Greek , neuter singular form of meaning "lazy" or "inactive") 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, water, and nitrogen from a sample of clean air. They first accomplished this by replicating an experiment of Henry Cavendish's. They trapped a mixture of atmospheric air with additional oxygen in a test-tube (A) upside-down over a large quantity of dilute alkali solution (B), which in Cavendish's original experiment was potassium hydroxide, and conveyed a current through wires insulated by U-shaped glass tubes (CC) which sealed around the platinum wire electrodes, leaving the ends of the wires (DD) exposed to the gas and insulated from the alkali solution. The arc was powered by a battery of five Grove cells and a Ruhmkorff coil of medium size. The alkali absorbed the oxides of nitrogen produced by the arc and also carbon dioxide. They operated the arc until no more reduction of volume of the gas could be seen for at least an hour or two and the spectral lines of nitrogen disappeared when the gas was examined. The remaining oxygen was reacted with alkaline pyrogallate to leave behind an apparently non-reactive gas which they called argon. | Argon | Wikipedia | 451 | 896 | https://en.wikipedia.org/wiki/Argon | Physical sciences | Chemical elements_2 | null |
Before isolating the gas, they had determined that nitrogen produced from chemical compounds was 0.5% lighter than nitrogen from the atmosphere. The difference was slight, but it was important enough to attract their attention for many months. They concluded that there was another gas in the air mixed in with the nitrogen. Argon was also encountered in 1882 through independent research of H. F. Newall and W. N. Hartley. Each observed new lines in the emission spectrum of air that did not match known elements.
Prior to 1957, the symbol for argon was "A". This was changed to Ar after the International Union of Pure and Applied Chemistry published the work Nomenclature of Inorganic Chemistry in 1957.
Occurrence
Argon constitutes 0.934% by volume and 1.288% by mass of Earth's atmosphere. Air is the primary industrial source of purified argon products. Argon is isolated from air by fractionation, most commonly by cryogenic fractional distillation, a process that also produces purified nitrogen, oxygen, neon, krypton and xenon. Earth's crust and seawater contain 1.2 ppm and 0.45 ppm of argon, respectively.
Isotopes
The main isotopes of argon found on Earth are (99.6%), (0.34%), and (0.06%). Naturally occurring , with a half-life of 1.25 years, decays to stable (11.2%) by electron capture or positron emission, and also to stable (88.8%) by beta decay. These properties and ratios are used to determine the age of rocks by K–Ar dating.
In Earth's atmosphere, is made by cosmic ray activity, primarily by neutron capture of followed by two-neutron emission. In the subsurface environment, it is also produced through neutron capture by , followed by proton emission. is created from the neutron capture by followed by an alpha particle emission as a result of subsurface nuclear explosions. It has a half-life of 35 days. | Argon | Wikipedia | 423 | 896 | https://en.wikipedia.org/wiki/Argon | Physical sciences | Chemical elements_2 | null |
Between locations in the Solar System, the isotopic composition of argon varies greatly. Where the major source of argon is the decay of in rocks, will be the dominant isotope, as it is on Earth. Argon produced directly by stellar nucleosynthesis is dominated by the alpha-process nuclide . Correspondingly, solar argon contains 84.6% (according to solar wind measurements), and 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 in Earth's atmosphere, which is only 31.5 ppmv (= 9340 ppmv × 0.337%), comparable with that of neon (18.18 ppmv) on Earth and with interplanetary gasses, measured by probes.
The atmospheres of Mars, Mercury and Titan (the largest moon of Saturn) contain argon, predominantly as .
The predominance of radiogenic is the reason the standard atomic weight of terrestrial argon is greater than that of the next element, potassium, a fact that was puzzling when argon was discovered. Mendeleev positioned the elements on his periodic table in order of atomic weight, but the inertness of argon suggested a placement before the reactive alkali metal. Henry Moseley later solved this problem by showing that the periodic table is actually arranged in order of atomic number (see History of the periodic table).
Compounds | Argon | Wikipedia | 316 | 896 | https://en.wikipedia.org/wiki/Argon | Physical sciences | Chemical elements_2 | null |
Argon's complete octet of electrons indicates full s and p subshells. This full valence shell makes argon very stable and extremely resistant to bonding with other elements. Before 1962, argon and the other noble gases were considered to be chemically inert and unable to form compounds; however, compounds of the heavier noble gases have since been synthesized. The first argon compound with tungsten pentacarbonyl, W(CO)5Ar, was isolated in 1975. However, it was not widely recognised at that time. In August 2000, another argon compound, argon fluorohydride (HArF), was formed by researchers at the University of Helsinki, by shining ultraviolet light onto frozen argon containing a small amount of hydrogen fluoride with caesium iodide. This discovery caused the recognition that argon could form weakly bound compounds, even though it was not the first. It is stable up to 17 kelvins (−256 °C). The metastable dication, which is valence-isoelectronic with carbonyl fluoride and phosgene, was observed in 2010. Argon-36, in the form of argon hydride (argonium) ions, has been detected in interstellar medium associated with the Crab Nebula supernova; this was the first noble-gas molecule detected in outer space.
Solid argon hydride (Ar(H2)2) has the same crystal structure as the MgZn2 Laves phase. It forms at pressures between 4.3 and 220 GPa, though Raman measurements suggest that the H2 molecules in Ar(H2)2 dissociate above 175 GPa.
Production
Argon is extracted industrially by the fractional distillation of liquid air in a cryogenic air separation unit; a process that separates liquid nitrogen, which boils at 77.3 K, from argon, which boils at 87.3 K, and liquid oxygen, which boils at 90.2 K. About 700,000 tonnes of argon are produced worldwide every year.
Applications
Argon has several desirable properties:
Argon is a chemically inert gas.
Argon is the cheapest alternative when nitrogen is not sufficiently inert.
Argon has low thermal conductivity.
Argon has electronic properties (ionization and/or the emission spectrum) desirable for some applications. | Argon | Wikipedia | 497 | 896 | https://en.wikipedia.org/wiki/Argon | Physical sciences | Chemical elements_2 | null |
Other noble gases would be equally suitable for most of these applications, but argon is by far the cheapest. It is inexpensive, since it occurs naturally in air and is readily obtained as a byproduct of cryogenic air separation in the production of liquid oxygen and liquid nitrogen: the primary constituents of air are used on a large industrial scale. The other noble gases (except helium) are produced this way as well, but argon is the most plentiful by far. The bulk of its applications arise simply because it is inert and relatively cheap.
Industrial processes
Argon is used in some high-temperature industrial processes where ordinarily non-reactive substances become reactive. For example, an argon atmosphere is used in graphite electric furnaces to prevent the graphite from burning.
For some of these processes, the presence of nitrogen or oxygen gases might cause defects within the material. Argon is used in some types of arc welding such as gas metal arc welding and gas tungsten arc welding, as well as in the processing of titanium and other reactive elements. An argon atmosphere is also used for growing crystals of silicon and germanium.
Argon is used in the poultry industry to asphyxiate birds, either for mass culling following disease outbreaks, or as a means of slaughter more humane than electric stunning. Argon is denser than air and displaces oxygen close to the ground during inert gas asphyxiation. Its non-reactive nature makes it suitable in a food product, and since it replaces oxygen within the dead bird, argon also enhances shelf life.
Argon is sometimes used for extinguishing fires where valuable equipment may be damaged by water or foam. | Argon | Wikipedia | 349 | 896 | https://en.wikipedia.org/wiki/Argon | Physical sciences | Chemical elements_2 | null |
Scientific research
Liquid argon is used as the target for neutrino experiments and direct dark matter searches. The interaction between the hypothetical WIMPs and an argon nucleus produces scintillation light that is detected by photomultiplier tubes. Two-phase detectors containing argon gas are used to detect the ionized electrons produced during the WIMP–nucleus scattering. As with most other liquefied noble gases, argon has a high scintillation light yield (about 51 photons/keV), is transparent to its own scintillation light, and is relatively easy to purify. Compared to xenon, argon is cheaper and has a distinct scintillation time profile, which allows the separation of electronic recoils from nuclear recoils. On the other hand, its intrinsic beta-ray background is larger due to contamination, unless one uses argon from underground sources, which has much less contamination. Most of the argon in Earth's atmosphere was produced by electron capture of long-lived ( + e− → + ν) present in natural potassium within Earth. The activity in the atmosphere is maintained by cosmogenic production through the knockout reaction (n,2n) and similar reactions. The half-life of is only 269 years. As a result, the underground Ar, shielded by rock and water, has much less contamination. Dark-matter detectors currently operating with liquid argon include DarkSide, WArP, ArDM, microCLEAN and DEAP. Neutrino experiments include ICARUS and MicroBooNE, both of which use high-purity liquid argon in a time projection chamber for fine grained three-dimensional imaging of neutrino interactions.
At Linköping University, Sweden, the inert gas is being utilized in a vacuum chamber in which plasma is introduced to ionize metallic films. This process results in a film usable for manufacturing computer processors. The new process would eliminate the need for chemical baths and use of expensive, dangerous and rare materials.
Preservative
Argon is used to displace oxygen- and moisture-containing air in packaging material to extend the shelf-lives of the contents (argon has the European food additive code E938). Aerial oxidation, hydrolysis, and other chemical reactions that degrade the products are retarded or prevented entirely. High-purity chemicals and pharmaceuticals are sometimes packed and sealed in argon. | Argon | Wikipedia | 488 | 896 | https://en.wikipedia.org/wiki/Argon | Physical sciences | Chemical elements_2 | null |
In winemaking, argon is used in a variety of activities to provide a barrier against oxygen at the liquid surface, which can spoil wine by fueling both microbial metabolism (as with acetic acid bacteria) and standard redox chemistry.
Argon is sometimes used as the propellant in aerosol cans.
Argon is also used as a preservative for such products as varnish, polyurethane, and paint, by displacing air to prepare a container for storage.
Since 2002, the American National Archives stores important national documents such as the Declaration of Independence and the Constitution within argon-filled cases to inhibit their degradation. Argon is preferable to the helium that had been used in the preceding five decades, because helium gas escapes through the intermolecular pores in most containers and must be regularly replaced.
Laboratory equipment
Argon may be used as the inert gas within Schlenk lines and gloveboxes. Argon is preferred to less expensive nitrogen in cases where nitrogen may react with the reagents or apparatus.
Argon may be used as the carrier gas in gas chromatography and in electrospray ionization mass spectrometry; it is the gas of choice for the plasma used in ICP spectroscopy. Argon is preferred for the sputter coating of specimens for scanning electron microscopy. Argon gas is also commonly used for sputter deposition of thin films as in microelectronics and for wafer cleaning in microfabrication.
Medical use
Cryosurgery procedures such as cryoablation use liquid argon to destroy tissue such as cancer cells. It is used in a procedure called "argon-enhanced coagulation", a form of argon plasma beam electrosurgery. The procedure carries a risk of producing gas embolism and has resulted in the death of at least one patient.
Blue argon lasers are used in surgery to weld arteries, destroy tumors, and correct eye defects.
Argon has also been used experimentally to replace nitrogen in the breathing or decompression mix known as Argox, to speed the elimination of dissolved nitrogen from the blood.
Lighting | Argon | Wikipedia | 444 | 896 | https://en.wikipedia.org/wiki/Argon | Physical sciences | Chemical elements_2 | null |
Incandescent lights are filled with argon, to preserve the filaments at high temperature from oxidation. It is used for the specific way it ionizes and emits light, such as in plasma globes and calorimetry in experimental particle physics. Gas-discharge lamps filled with pure argon provide lilac/violet light; with argon and some mercury, blue light. Argon is also used for blue and green argon-ion lasers.
Miscellaneous uses
Argon is used for thermal insulation in energy-efficient windows. Argon is also used in technical scuba diving to inflate a dry suit because it is inert and has low thermal conductivity.
Argon is used as a propellant in the development of the Variable Specific Impulse Magnetoplasma Rocket (VASIMR). Compressed argon gas is allowed to expand, to cool the seeker heads of some versions of the AIM-9 Sidewinder missile and other missiles that use cooled thermal seeker heads. The gas is stored at high pressure.
Argon-39, with a half-life of 269 years, has been used for a number of applications, primarily ice core and ground water dating. Also, potassium–argon dating and related argon-argon dating are used to date sedimentary, metamorphic, and igneous rocks.
Argon has been used by athletes as a doping agent to simulate hypoxic conditions. In 2014, the World Anti-Doping Agency (WADA) added argon and xenon to the list of prohibited substances and methods, although at this time there is no reliable test for abuse.
Safety
Although argon is non-toxic, it is 38% more dense than air and therefore considered a dangerous asphyxiant in closed areas. It is difficult to detect because it is colorless, odorless, and tasteless. A 1994 incident, in which a man was asphyxiated after entering an argon-filled section of oil pipe under construction in Alaska, highlights the dangers of argon tank leakage in confined spaces and emphasizes the need for proper use, storage and handling. | Argon | Wikipedia | 429 | 896 | https://en.wikipedia.org/wiki/Argon | Physical sciences | Chemical elements_2 | null |
Arsenic is a chemical element with the symbol As and the atomic number 33. It is a metalloid and one of the pnictogens, and therefore shares many properties with its group 15 neighbors phosphorus and antimony. Arsenic is a notoriously toxic heavy metal. It occurs naturally in many minerals, usually in combination with sulfur and metals, but also as a pure elemental crystal. It has various allotropes, but only the grey form, which has a metallic appearance, is important to industry.
The primary use of arsenic is in alloys of lead (for example, in car batteries and ammunition). Arsenic is also a common n-type dopant in semiconductor electronic devices, and a component of the III–V compound semiconductor gallium arsenide. Arsenic and its compounds, especially the trioxide, are used in the production of pesticides, treated wood products, herbicides, and insecticides. These applications are declining with the increasing recognition of the toxicity of arsenic and its compounds.
Arsenic has been known since ancient times to be poisonous to humans. However, a few species of bacteria are able to use arsenic compounds as respiratory metabolites. Trace quantities of arsenic have been proposed to be an essential dietary element in rats, hamsters, goats, and chickens. Research has not been conducted to determine whether small amounts of arsenic may play a role in human metabolism. However, arsenic poisoning occurs in multicellular life if quantities are larger than needed. Arsenic contamination of groundwater is a problem that affects millions of people across the world.
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 number 1 in its 2001 prioritized list of hazardous substances at Superfund sites. Arsenic is classified as a Group-A carcinogen.
Characteristics
Physical characteristics | Arsenic | Wikipedia | 376 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
The three most common arsenic allotropes are grey, yellow, and black arsenic, with grey being the most common. Grey arsenic (α-As, space group Rm No. 166) adopts a double-layered structure consisting of many interlocked, ruffled, six-membered rings. Because of weak bonding between the layers, grey arsenic is brittle and has a relatively 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 slightly closer than the three atoms in the next. This relatively close packing leads to a high density of 5.73 g/cm3. Grey arsenic is a semimetal, but becomes a semiconductor with a bandgap of 1.2–1.4 eV if amorphized. Grey arsenic is also the most stable form.
Yellow arsenic is soft and waxy, and 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, and most toxic. Solid yellow arsenic is produced by rapid cooling of arsenic vapor, . It is rapidly transformed into grey 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 also be formed by cooling vapor at around 100–220 °C and by crystallization of amorphous arsenic in the presence of mercury vapors. It is glassy and brittle. Black arsenic is also a poor electrical conductor.
Arsenic sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state at . The triple point is at 3.63 MPa and .
Isotopes | Arsenic | Wikipedia | 386 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Arsenic occurs in nature as one stable isotope, 75As, and is therefore called a monoisotopic element. As of 2024, at least 32 radioisotopes have also been synthesized, ranging in atomic mass from 64 to 95. 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 (t1/2=65.30 hours), 72As (t1/2=26.0 hours), 74As (t1/2=17.77 days), 76As (t1/2=26.26 hours), and 77As (t1/2=38.83 hours). Isotopes that are lighter than the stable 75As tend to decay by β+ decay, and 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.
Chemistry | Arsenic | Wikipedia | 232 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Arsenic has a similar electronegativity and ionization energies to its lighter pnictogen congener phosphorus and therefore readily forms covalent molecules with most of the nonmetals. Though stable in dry air, arsenic forms a golden-bronze tarnish upon exposure to humidity which eventually becomes a black surface layer. When heated in air, arsenic oxidizes to arsenic trioxide; the fumes from this reaction have an odor resembling garlic. 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, and in fluorine to give arsenic pentafluoride. Arsenic makes arsenic acid with concentrated nitric acid, arsenous acid with dilute nitric acid, and arsenic trioxide with concentrated sulfuric acid; however, it does not react with water, alkalis, or non-oxidising acids. 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 highly endothermic and even the group 1 arsenides have properties of intermetallic compounds. Like germanium, selenium, and bromine, which like arsenic succeed the 3d transition series, arsenic is much less stable in the +5 oxidation state than its vertical neighbors phosphorus and antimony, and hence arsenic pentoxide and arsenic acid are potent oxidizers.
Compounds
Compounds of arsenic resemble, in some respects, those of phosphorus, which occupies the same group (column) 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, and +5 in the arsenates and most organoarsenic compounds. Arsenic also bonds readily to itself as seen in the square ions in the mineral skutterudite. In the +3 oxidation state, arsenic is typically pyramidal owing to the influence of the lone pair of electrons.
Inorganic compounds | Arsenic | Wikipedia | 439 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
One of the simplest arsenic compounds is the trihydride, the highly toxic, flammable, pyrophoric arsine (AsH3). This compound is generally regarded as stable, since at room temperature it decomposes only slowly. At temperatures of 250–300 °C decomposition to arsenic and hydrogen is rapid. Several factors, such as humidity, presence of light and certain catalysts (namely aluminium) facilitate the rate of decomposition. It oxidises readily in air to form arsenic trioxide and water, and analogous reactions take place with sulfur and selenium instead of oxygen.
Arsenic forms colorless, odorless, crystalline oxides As2O3 ("white arsenic") and As2O5 which are hygroscopic and readily soluble in water to form acidic solutions. Arsenic(V) acid is a weak acid and its salts, known as arsenates, are a major source of arsenic contamination of groundwater in regions with high levels of naturally-occurring arsenic minerals. Synthetic arsenates include Scheele's Green (cupric hydrogen arsenate, acidic copper arsenate), calcium arsenate, and lead hydrogen arsenate. These three have been used as agricultural insecticides and poisons.
The protonation steps between the arsenate and arsenic acid are similar to those between phosphate and phosphoric acid. Unlike phosphorous acid, arsenous acid is genuinely tribasic, with the formula As(OH)3.
A broad variety of sulfur compounds of arsenic are known. Orpiment (As2S3) and realgar (As4S4) are somewhat abundant and were formerly used as painting pigments. In As4S10, arsenic has a formal oxidation state of +2 in As4S4 which features As-As bonds so that the total covalency of As is still 3. Both orpiment and realgar, as well as As4S3, have selenium analogs; the analogous As2Te3 is known as the mineral kalgoorlieite, and the anion As2Te− is known as a ligand in cobalt complexes. | Arsenic | Wikipedia | 444 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
All trihalides of arsenic(III) are well known except the astatide, which is unknown. Arsenic pentafluoride (AsF5) is the only important pentahalide, reflecting the lower stability of the +5 oxidation state; even so, it is a very strong fluorinating and oxidizing agent. (The pentachloride is stable only below −50 °C, at which temperature it decomposes to the trichloride, releasing chlorine gas.)
Alloys
Arsenic is used as the group 5 element in the III-V semiconductors gallium arsenide, indium arsenide, and aluminium arsenide. The valence electron count of GaAs is the same as a pair of Si atoms, but the band structure is completely different which results in distinct bulk properties. Other arsenic alloys include the II-V semiconductor cadmium arsenide.
Organoarsenic compounds
A large variety of organoarsenic compounds are known. Several were developed as chemical warfare agents during World War I, including vesicants such as lewisite and vomiting agents such as adamsite. Cacodylic acid, which is of historic and practical interest, arises from the methylation of arsenic trioxide, a reaction that has no analogy in phosphorus chemistry. Cacodyl was the first organometallic compound known (even though arsenic is not a true metal) and was named from the Greek κακωδία "stink" for its offensive, garlic-like odor; it is very toxic.
Occurrence and production
Arsenic is the 53rd most abundant element in the Earth's crust, comprising about 1.5 parts per million (0.00015%). Typical background concentrations of arsenic do not exceed 3 ng/m3 in the atmosphere; 100 mg/kg in soil; 400 μg/kg in vegetation; 10 μg/L in freshwater and 1.5 μg/L in seawater. Arsenic is the 22nd most abundant element in seawater and ranks 41st in abundance in the universe.
Minerals with the formula MAsS and MAs2 (M = Fe, Ni, Co) are the dominant commercial sources of arsenic, together with realgar (an arsenic sulfide mineral) and native (elemental) arsenic. An illustrative mineral is arsenopyrite (FeAsS), which is structurally related to iron pyrite. Many minor As-containing minerals are known. Arsenic also occurs in various organic forms in the environment. | Arsenic | Wikipedia | 509 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
In 2014, China was the top producer of white arsenic with almost 70% world share, followed by Morocco, Russia, and Belgium, according to the British Geological Survey and the United States Geological Survey. Most arsenic refinement operations in the US and Europe have closed over environmental concerns. Arsenic is found in the smelter dust from copper, gold, and lead smelters, and is recovered primarily from copper refinement dust.
On roasting arsenopyrite in air, arsenic sublimes as arsenic(III) oxide leaving iron oxides, while roasting without air results in the production of gray arsenic. Further purification from sulfur and other chalcogens is achieved by sublimation in vacuum, in a hydrogen atmosphere, or by distillation from molten lead-arsenic mixture.
History
The word arsenic has its origin in the Syriac word zarnika, from Arabic al-zarnīḵ 'the orpiment', based on Persian zar ("gold") from the word zarnikh, meaning "yellow" (literally "gold-colored") and hence "(yellow) orpiment". It was adopted into Greek (using folk etymology) as arsenikon () – a neuter form of the Greek adjective arsenikos (), meaning "male", "virile".
Latin-speakers adopted the Greek term as , which in French ultimately became , whence the English word "arsenic".
Arsenic sulfides (orpiment, realgar) and oxides have been known and used since ancient times. Zosimos () describes roasting sandarach (realgar) to obtain cloud of arsenic (arsenic trioxide), which he then reduces to gray arsenic. As the symptoms of arsenic poisoning are not very specific, the substance was frequently used for murder until the advent in the 1830s of the Marsh test, a sensitive chemical test for its presence. (Another less sensitive but more general test is the Reinsch test.) Owing to its use by the ruling class to murder one another and its potency and discreetness, arsenic has been called the "poison of kings" and the "king of poisons". Arsenic became known as "the inheritance powder" due to its use in killing family members in the Renaissance era. | Arsenic | Wikipedia | 465 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
During the Bronze Age, arsenic was melted with copper to make arsenical bronze.
Jabir ibn Hayyan described the isolation of arsenic before 815 AD. Albertus Magnus (Albert the Great, 1193–1280) later isolated the element from a compound in 1250, by heating soap together with arsenic trisulfide. In 1649, Johann Schröder published two ways of preparing arsenic. Crystals of elemental (native) arsenic are found in nature, although rarely.
Cadet's fuming liquid (impure cacodyl), often claimed as the first synthetic organometallic compound, was synthesized in 1760 by Louis Claude Cadet de Gassicourt through the reaction of potassium acetate with arsenic trioxide.
In the Victorian era, women would eat "arsenic" ("white arsenic" or arsenic trioxide) mixed with vinegar and chalk to improve the complexion of their faces, making their skin paler (to show they did not work in the fields). The accidental use of arsenic in the adulteration of foodstuffs led to the Bradford sweet poisoning in 1858, which resulted in 21 deaths. From the late-18th century wallpaper production began to use dyes made from arsenic,
which was thought to increase the pigment's brightness. One account of the illness and 1821 death of Napoleon I implicates arsenic poisoning involving wallpaper.
Two arsenic pigments have been widely used since their discovery – Paris Green in 1814 and Scheele's Green in 1775. After the toxicity of arsenic became widely known, these chemicals were used less often as pigments and more often as insecticides. In the 1860s, an arsenic byproduct of dye production, London Purple, was widely used. This was a solid mixture of arsenic trioxide, aniline, lime, and ferrous oxide, insoluble in water and very toxic by inhalation or ingestion But it was later replaced with Paris Green, another arsenic-based dye. With better understanding of the toxicology mechanism, two other compounds were used starting in the 1890s. Arsenite of lime and arsenate of lead were used widely as insecticides until the discovery of DDT in 1942. | Arsenic | Wikipedia | 445 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
In small doses, soluble arsenic compounds act as stimulants, and were once popular as medicine by people in the mid-18th to 19th centuries; this use was especially prevalent for sport animals such as race horses or work dogs and continued into the 20th century.
A 2006 study of the remains of the Australian racehorse Phar Lap determined that its 1932 death was caused by a massive overdose of arsenic. Sydney veterinarian Percy Sykes stated, "In those days, arsenic was quite a common tonic, usually given in the form of a solution (Fowler's Solution) ... It was so common that I'd reckon 90 per cent of the horses had arsenic in their system."
Applications
Agricultural
The toxicity of arsenic to insects, bacteria, and fungi led to its use as a wood preservative. In the 1930s, a process of treating wood with chromated copper arsenate (also known as CCA or Tanalith) was invented, and for decades, this treatment was the most extensive industrial use of arsenic. An increased appreciation of the toxicity of arsenic led to a ban of CCA in consumer products in 2004, initiated by the European Union and United States. However, CCA remains in heavy use in other countries (such as on Malaysian rubber plantations).
Arsenic was also used in various agricultural insecticides and poisons. For example, lead hydrogen arsenate was a common insecticide on fruit trees, but contact with the compound sometimes resulted in brain damage among those working the sprayers. In the second half of the 20th century, monosodium methyl arsenate (MSMA) and disodium methyl arsenate (DSMA) – less toxic organic forms of arsenic – replaced lead arsenate in agriculture. These organic arsenicals were in turn phased out in the United States by 2013 in all agricultural activities except cotton farming.
The biogeochemistry of arsenic is complex and includes various adsorption and desorption processes. The toxicity of arsenic is connected to its solubility and is affected by pH. Arsenite () is more soluble than arsenate () and is more toxic; however, at a lower pH, arsenate becomes more mobile and toxic. It was found that addition of sulfur, phosphorus, and iron oxides to high-arsenite soils greatly reduces arsenic phytotoxicity. | Arsenic | Wikipedia | 482 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Arsenic is used as a feed additive in poultry and swine production, in particular it was used in the U.S. until 2015 to increase weight gain, improve feed efficiency, and prevent disease. An example is roxarsone, which had been used as a broiler starter by about 70% of U.S. broiler growers. In 2011, Alpharma, a subsidiary of Pfizer Inc., which produces roxarsone, voluntarily suspended sales of the drug in response to studies showing elevated levels of inorganic arsenic, a carcinogen, in treated chickens. A successor to Alpharma, Zoetis, continued to sell nitarsone until 2015, primarily for use in turkeys.
Medical use
During the 17th, 18th, and 19th centuries, a number of arsenic compounds were used as medicines, including arsphenamine (by Paul Ehrlich) and arsenic trioxide (by Thomas Fowler), for treating diseases such as cancer or psoriasis. Arsphenamine, as well as neosalvarsan, was indicated for syphilis, but has been superseded by modern antibiotics. However, arsenicals such as melarsoprol are still used for the treatment of trypanosomiasis in spite of their severe toxicity, since the disease is almost uniformly fatal if untreated. In 2000 the US Food and Drug Administration approved arsenic trioxide for the treatment of patients with acute promyelocytic leukemia that is resistant to all-trans retinoic acid.
A 2008 paper reports success in locating tumors using arsenic-74 (a positron emitter). This isotope produces clearer PET scan images than the previous radioactive agent, iodine-124, because the body tends to transport iodine to the thyroid gland producing signal noise. Nanoparticles of arsenic have shown ability to kill cancer cells with lesser cytotoxicity than other arsenic formulations.
Alloys | Arsenic | Wikipedia | 390 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
The main use of arsenic is in alloying with lead. Lead components in car batteries are strengthened by the presence of a very small percentage of arsenic. Dezincification of brass (a copper-zinc alloy) is greatly reduced by the addition of arsenic. "Phosphorus Deoxidized Arsenical Copper" with an arsenic content of 0.3% has an increased corrosion stability in certain environments. Gallium arsenide is an important semiconductor material, used in integrated circuits. Circuits made from GaAs are much faster (but also much more expensive) than those made from silicon. Unlike silicon, GaAs has a direct bandgap, and can be used in laser diodes and LEDs to convert electrical energy directly into light.
Military
After World War I, the United States built a stockpile of 20,000 tons of weaponized lewisite (ClCH=CHAsCl2), an organoarsenic vesicant (blister agent) and lung irritant. The stockpile was neutralized with bleach and dumped into the Gulf of Mexico in the 1950s. During the Vietnam War, the United States used Agent Blue, a mixture of sodium cacodylate and its acid form, as one of the rainbow herbicides to deprive North Vietnamese soldiers of foliage cover and rice.
Other uses
Copper acetoarsenite was used as a green pigment known under many names, including Paris Green and Emerald Green. It caused numerous arsenic poisonings. Scheele's Green, a copper arsenate, was used in the 19th century as a coloring agent in sweets.
Arsenic is used in bronzing.
As much as 2% of produced arsenic is used in lead alloys for lead shot and bullets.
Arsenic is added in small quantities to alpha-brass to make it dezincification-resistant. This grade of brass is used in plumbing fittings and other wet environments.
Arsenic is also used for taxonomic sample preservation. It was also used in embalming fluids historically.
Arsenic was used in the taxidermy process up until the 1980s.
Arsenic was used as an opacifier in ceramics, creating white glazes.
Until recently, arsenic was used in optical glass. Modern glass manufacturers have ceased using both arsenic and lead.
Biological role | Arsenic | Wikipedia | 462 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Bacteria
Some species of bacteria obtain their energy in the absence of oxygen by oxidizing various fuels while reducing arsenate to arsenite. Under oxidative environmental conditions some bacteria use arsenite as fuel, which they oxidize to arsenate. The enzymes involved are known as arsenate reductases (Arr).
In 2008, bacteria were discovered that employ a version of photosynthesis in the absence of oxygen with arsenites as electron donors, producing arsenates (just as ordinary photosynthesis uses water as electron donor, producing molecular oxygen). Researchers conjecture that, over the course of history, these photosynthesizing organisms produced the arsenates that allowed the arsenate-reducing bacteria to thrive. One strain, PHS-1, has been isolated and is related to the gammaproteobacterium Ectothiorhodospira shaposhnikovii. The mechanism is unknown, but an encoded Arr enzyme may function in reverse to its known homologues.
In 2011, it was postulated that the Halomonadaceae strain GFAJ-1 could be grown in the absence of phosphorus if that element were substituted with arsenic, exploiting the fact that the arsenate and phosphate anions are similar structurally. The study was widely criticised and subsequently refuted by independent researcher groups.
Potential role in higher animals
Arsenic may be an essential trace mineral in birds, involved in the synthesis of methionine metabolites. However, the role of arsenic in bird nutrition is disputed, as other authors state that arsenic is toxic in small amounts
Some evidence indicates that arsenic is an essential trace mineral in mammals.
Heredity
Arsenic has been linked to epigenetic changes, heritable changes in gene expression that occur without changes in DNA sequence. These include DNA methylation, histone modification, and RNA interference. Toxic levels of arsenic cause significant DNA hypermethylation of tumor suppressor genes p16 and p53, thus increasing risk of carcinogenesis. These epigenetic events have been studied in vitro using human kidney cells and in vivo using rat liver cells and peripheral blood leukocytes in humans. Inductively coupled plasma mass spectrometry (ICP-MS) is used to detect precise levels of intracellular arsenic and other arsenic bases involved in epigenetic modification of DNA. Studies investigating arsenic as an epigenetic factor can be used to develop precise biomarkers of exposure and susceptibility. | Arsenic | Wikipedia | 509 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
The Chinese brake fern (Pteris vittata) hyperaccumulates arsenic from the soil into its leaves and has a proposed use in phytoremediation.
Biomethylation
Inorganic arsenic and its compounds, upon entering the food chain, are progressively metabolized through a process of methylation. For example, the mold Scopulariopsis brevicaulis produces trimethylarsine if inorganic arsenic is present. The organic compound arsenobetaine is found in some marine foods such as fish and algae, and also in mushrooms in larger concentrations. The average person's intake is about 10–50 μg/day. Values about 1000 μg are not unusual following consumption of fish or mushrooms, but there is little danger in eating fish because this arsenic compound is nearly non-toxic.
Environmental issues
Exposure
Naturally occurring sources of human exposure include volcanic ash, weathering of minerals and ores, and mineralized groundwater. Arsenic is also found in food, water, soil, and air. Arsenic is absorbed by all plants, but is more concentrated in leafy vegetables, rice, apple and grape juice, and seafood. An additional route of exposure is inhalation of atmospheric gases and dusts.
During the Victorian era, arsenic was widely used in home decor, especially wallpapers. In Europe, an analysis based on 20,000 soil samples across all 28 countries show that 98% of sampled soils have concentrations less than 20 mg kg-1. In addition, the As hotspots are related to frequent fertilization and close distance to mining activities.
Occurrence in drinking water | Arsenic | Wikipedia | 324 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Extensive arsenic contamination of groundwater has led to widespread arsenic poisoning in Bangladesh and neighboring countries. It is estimated that approximately 57 million people in the Bengal basin are drinking groundwater with arsenic concentrations elevated above the World Health Organization's standard of 10 parts per billion (ppb). However, a study of cancer rates in Taiwan suggested that significant increases in cancer mortality appear only at levels above 150 ppb. The arsenic in the groundwater is of natural origin, and is released from the sediment into the groundwater, caused by the anoxic conditions of the subsurface. This groundwater was used after local and western NGOs and the Bangladeshi government undertook a massive shallow tube well drinking-water program in the late twentieth century. This program was designed to prevent drinking of bacteria-contaminated surface waters, but failed to test for arsenic in the groundwater. Many other countries and districts in Southeast Asia, such as Vietnam and Cambodia, have geological environments that produce groundwater with a high arsenic content. Arsenicosis was reported in Nakhon Si Thammarat, Thailand, in 1987, and the Chao Phraya River probably contains high levels of naturally occurring dissolved arsenic without being a public health problem because much of the public uses bottled water. In Pakistan, more than 60 million people are exposed to arsenic polluted drinking water indicated by a 2017 report in Science. Podgorski's team investigated more than 1200 samples and more than 66% exceeded the WHO minimum contamination level.
Since the 1980s, residents of the Ba Men region of Inner Mongolia, China have been chronically exposed to arsenic through drinking water from contaminated wells. A 2009 research study observed an elevated presence of skin lesions among residents with well water arsenic concentrations between 5 and 10 μg/L, suggesting that arsenic induced toxicity may occur at relatively low concentrations with chronic exposure. Overall, 20 of China's 34 provinces have high arsenic concentrations in the groundwater supply, potentially exposing 19 million people to hazardous drinking water.
A study by IIT Kharagpur found high levels of Arsenic in groundwater of 20% of India's land, exposing more than 250 million people. States such as Punjab, Bihar, West Bengal, Assam, Haryana, Uttar Pradesh, and Gujarat have highest land area exposed to arsenic. | Arsenic | Wikipedia | 443 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
In the United States, arsenic is most commonly found in the ground waters of the southwest. Parts of New England, Michigan, Wisconsin, Minnesota and the Dakotas are also known to have significant concentrations of arsenic in ground water. Increased levels of skin cancer have been associated with arsenic exposure in Wisconsin, even at levels below the 10 ppb drinking water standard. According to a recent film funded by the US Superfund, millions of private wells have unknown arsenic levels, and in some areas of the US, more than 20% of the wells may contain levels that exceed established limits.
Low-level exposure to arsenic at concentrations of 100 ppb (i.e., above the 10 ppb drinking water standard) compromises the initial immune response to H1N1 or swine flu infection according to NIEHS-supported scientists. The study, conducted in laboratory mice, suggests that people exposed to arsenic in their drinking water may be at increased risk for more serious illness or death from the virus.
Some Canadians are drinking water that contains inorganic arsenic. Private-dug–well waters are most at risk for containing inorganic arsenic. Preliminary well water analysis typically does not test for arsenic. Researchers at the Geological Survey of Canada have modeled relative variation in natural arsenic hazard potential for the province of New Brunswick. This study has important implications for potable water and health concerns relating to inorganic arsenic.
Epidemiological evidence from Chile shows a dose-dependent connection between chronic arsenic exposure and various forms of cancer, in particular when other risk factors, such as cigarette smoking, are present. These effects have been demonstrated at contaminations less than 50 ppb. Arsenic is itself a constituent of tobacco smoke.
Analyzing multiple epidemiological studies on inorganic arsenic exposure suggests a small but measurable increase in risk for bladder cancer at 10 ppb. According to Peter Ravenscroft of the Department of Geography at the University of Cambridge, roughly 80 million people worldwide consume between 10 and 50 ppb arsenic in their drinking water. If they all consumed exactly 10 ppb arsenic in their drinking water, the previously cited multiple epidemiological study analysis would predict an additional 2,000 cases of bladder cancer alone. This represents a clear underestimate of the overall impact, since it does not include lung or skin cancer, and explicitly underestimates the exposure. Those exposed to levels of arsenic above the current WHO standard should weigh the costs and benefits of arsenic remediation. | Arsenic | Wikipedia | 491 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Early (1973) evaluations of the processes for removing dissolved arsenic from drinking water demonstrated the efficacy of co-precipitation with either iron or aluminium oxides. In particular, iron as a coagulant was found to remove arsenic with an efficacy exceeding 90%. Several adsorptive media systems have been approved for use at point-of-service in a study funded by the United States Environmental Protection Agency (US EPA) and the National Science Foundation (NSF). A team of European and Indian scientists and engineers have set up six arsenic treatment plants in West Bengal based on in-situ remediation method (SAR Technology). This technology does not use any chemicals and arsenic is left in an insoluble form (+5 state) in the subterranean zone by recharging aerated water into the aquifer and developing an oxidation zone that supports arsenic oxidizing micro-organisms. This process does not produce any waste stream or sludge and is relatively cheap.
Another effective and inexpensive method to avoid arsenic contamination is to sink wells 500 feet or deeper to reach purer waters. A recent 2011 study funded by the US National Institute of Environmental Health Sciences' Superfund Research Program shows that deep sediments can remove arsenic and take it out of circulation. In this process, called adsorption, arsenic sticks to the surfaces of deep sediment particles and is naturally removed from the ground water.
Magnetic separations of arsenic at very low magnetic field gradients with high-surface-area and monodisperse magnetite (Fe3O4) nanocrystals have been demonstrated in point-of-use water purification. Using the high specific surface area of Fe3O4 nanocrystals, the mass of waste associated with arsenic removal from water has been dramatically reduced.
Epidemiological studies have suggested a correlation between chronic consumption of drinking water contaminated with arsenic and the incidence of all leading causes of mortality. The literature indicates that arsenic exposure is causative in the pathogenesis of diabetes.
Chaff-based filters have recently been shown to reduce the arsenic content of water to 3 μg/L. This may find applications in areas where the potable water is extracted from underground aquifers.
San Pedro de Atacama
For several centuries, the people of San Pedro de Atacama in Chile have been drinking water that is contaminated with arsenic, and some evidence suggests they have developed some immunity. | Arsenic | Wikipedia | 487 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Hazard maps for contaminated groundwater
Around one-third of the world's population drinks water from groundwater resources. Of this, about 10 percent, approximately 300 million people, obtains water from groundwater resources that are contaminated with unhealthy levels of arsenic or fluoride. These trace elements derive mainly from minerals and ions in the ground.
Redox transformation of arsenic in natural waters
Arsenic is unique among the trace metalloids and oxyanion-forming trace metals (e.g. As, Se, Sb, Mo, V, Cr, U, Re). It is sensitive to mobilization at pH values typical of natural waters (pH 6.5–8.5) under both oxidizing and reducing conditions. Arsenic can occur in the environment in several oxidation states (−3, 0, +3 and +5), but in natural waters it is mostly found in inorganic forms as oxyanions of trivalent arsenite [As(III)] or pentavalent arsenate [As(V)]. Organic forms of arsenic are produced by biological activity, mostly in surface waters, but are rarely quantitatively important. Organic arsenic compounds may, however, occur where waters are significantly impacted by industrial pollution.
Arsenic may be solubilized by various processes. When pH is high, arsenic may be released from surface binding sites that lose their positive charge. When water level drops and sulfide minerals are exposed to air, arsenic trapped in sulfide minerals can be released into water. When organic carbon is present in water, bacteria are fed by directly reducing As(V) to As(III) or by reducing the element at the binding site, releasing inorganic arsenic.
The aquatic transformations of arsenic are affected by pH, reduction-oxidation potential, organic matter concentration and the concentrations and forms of other elements, especially iron and manganese. The main factors are pH and the redox potential. Generally, the main forms of arsenic under oxic conditions are , , , and at pH 2, 2–7, 7–11 and 11, respectively. Under reducing conditions, is predominant at pH 2–9. | Arsenic | Wikipedia | 429 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Oxidation and reduction affects the migration of arsenic in subsurface environments. Arsenite is the most stable soluble form of arsenic in reducing environments and arsenate, which is less mobile than arsenite, is dominant in oxidizing environments at neutral pH. Therefore, arsenic may be more mobile under reducing conditions. The reducing environment is also rich in organic matter which may enhance the solubility of arsenic compounds. As a result, the adsorption of arsenic is reduced and dissolved arsenic accumulates in groundwater. That is why the arsenic content is higher in reducing environments than in oxidizing environments.
The presence of sulfur is another factor that affects the transformation of arsenic in natural water. Arsenic can precipitate when metal sulfides form. In this way, arsenic is removed from the water and its mobility decreases. When oxygen is present, bacteria oxidize reduced sulfur to generate energy, potentially releasing bound arsenic.
Redox reactions involving Fe also appear to be essential factors in the fate of arsenic in aquatic systems. The reduction of iron oxyhydroxides plays a key role in the release of arsenic to water. So arsenic can be enriched in water with elevated Fe concentrations. Under oxidizing conditions, arsenic can be mobilized from pyrite or iron oxides especially at elevated pH. Under reducing conditions, arsenic can be mobilized by reductive desorption or dissolution when associated with iron oxides. The reductive desorption occurs under two circumstances. One is when arsenate is reduced to arsenite which adsorbs to iron oxides less strongly. The other results from a change in the charge on the mineral surface which leads to the desorption of bound arsenic. | Arsenic | Wikipedia | 347 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Some species of bacteria catalyze redox transformations of arsenic. Dissimilatory arsenate-respiring prokaryotes (DARP) speed up the reduction of As(V) to As(III). DARP use As(V) as the electron acceptor of anaerobic respiration and obtain energy to survive. Other organic and inorganic substances can be oxidized in this process. Chemoautotrophic arsenite oxidizers (CAO) and heterotrophic arsenite oxidizers (HAO) convert As(III) into As(V). CAO combine the oxidation of As(III) with the reduction of oxygen or nitrate. They use obtained energy to fix produce organic carbon from CO2. HAO cannot obtain energy from As(III) oxidation. This process may be an arsenic detoxification mechanism for the bacteria.
Equilibrium thermodynamic calculations predict that As(V) concentrations should be greater than As(III) concentrations in all but strongly reducing conditions, i.e. where sulfate reduction is occurring. However, abiotic redox reactions of arsenic are slow. Oxidation of As(III) by dissolved O2 is a particularly slow reaction. For example, Johnson and Pilson (1975) gave half-lives for the oxygenation of As(III) in seawater ranging from several months to a year. In other studies, As(V)/As(III) ratios were stable over periods of days or weeks during water sampling when no particular care was taken to prevent oxidation, again suggesting relatively slow oxidation rates. Cherry found from experimental studies that the As(V)/As(III) ratios were stable in anoxic solutions for up to 3 weeks but that gradual changes occurred over longer timescales. Sterile water samples have been observed to be less susceptible to speciation changes than non-sterile samples. Oremland found that the reduction of As(V) to As(III) in Mono Lake was rapidly catalyzed by bacteria with rate constants ranging from 0.02 to 0.3-day−1. | Arsenic | Wikipedia | 432 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Wood preservation in the US
As of 2002, US-based industries consumed 19,600 metric tons of arsenic. Ninety percent of this was used for treatment of wood with chromated copper arsenate (CCA). In 2007, 50% of the 5,280 metric tons of consumption was still used for this purpose. In the United States, the voluntary phasing-out of arsenic in production of consumer products and residential and general consumer construction products began on 31 December 2003, and alternative chemicals are now used, such as Alkaline Copper Quaternary, borates, copper azole, cyproconazole, and propiconazole.
Although discontinued, this application is also one of the most concerning to the general public. The vast majority of older pressure-treated wood was treated with CCA. CCA lumber is still in widespread use in many countries, and was heavily used during the latter half of the 20th century as a structural and outdoor building material. Although the use of CCA lumber was banned in many areas after studies showed that arsenic could leach out of the wood into the surrounding soil (from playground equipment, for instance), a risk is also presented by the burning of older CCA timber. The direct or indirect ingestion of wood ash from burnt CCA lumber has caused fatalities in animals and serious poisonings in humans; the lethal human dose is approximately 20 grams of ash. Scrap CCA lumber from construction and demolition sites may be inadvertently used in commercial and domestic fires. Protocols for safe disposal of CCA lumber are not consistent throughout the world. Widespread landfill disposal of such timber raises some concern, but other studies have shown no arsenic contamination in the groundwater.
Mapping of industrial releases in the US
One tool that maps the location (and other information) of arsenic releases in the United States is TOXMAP. TOXMAP is a Geographic Information System (GIS) from the Division of Specialized Information Services of the United States National Library of Medicine (NLM) funded by the US Federal Government. With marked-up maps of the United States, TOXMAP enables users to visually explore data from the United States Environmental Protection Agency's (EPA) Toxics Release Inventory and Superfund Basic Research Programs. TOXMAP's chemical and environmental health information is taken from NLM's Toxicology Data Network (TOXNET), PubMed, and from other authoritative sources.
Bioremediation | Arsenic | Wikipedia | 497 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Physical, chemical, and biological methods have been used to remediate arsenic contaminated water. Bioremediation is said to be cost-effective and environmentally friendly. Bioremediation of ground water contaminated with arsenic aims to convert arsenite, the toxic form of arsenic to humans, to arsenate. Arsenate (+5 oxidation state) is the dominant form of arsenic in surface water, while arsenite (+3 oxidation state) is the dominant form in hypoxic to anoxic environments. Arsenite is more soluble and mobile than arsenate. Many species of bacteria can transform arsenite to arsenate in anoxic conditions by using arsenite as an electron donor. This is a useful method in ground water remediation. Another bioremediation strategy is to use plants that accumulate arsenic in their tissues via phytoremediation but the disposal of contaminated plant material needs to be considered.
Bioremediation requires careful evaluation and design in accordance with existing conditions. Some sites may require the addition of an electron acceptor while others require microbe supplementation (bioaugmentation). Regardless of the method used, only constant monitoring can prevent future contamination.
Arsenic removal
Coagulation and flocculation are closely related processes common in arsenate removal from water. Due to the net negative charge carried by arsenate ions, they settle slowly or not at all due to charge repulsion. In coagulation, a positively charged coagulent such as iron and aluminum (commonly used salts: FeCl3, Fe2(SO4)3, Al2(SO4)3) neutralize the negatively charged arsenate, enable it to settle. Flocculation follows where a flocculant bridges smaller particles and allows the aggregate to precipitate out from water. However, such methods may not be efficient on arsenite as As(III) exists in uncharged arsenious acid, H3AsO3, at near-neutral pH.
The major drawbacks of coagulation and flocculation are the costly disposal of arsenate-concentrated sludge, and possible secondary contamination of environment. Moreover, coagulents such as iron may produce ion contamination that exceeds safety levels.
Toxicity and precautions | Arsenic | Wikipedia | 470 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Arsenic and many of its compounds are especially potent poisons (e.g. arsine). Small amount of arsenic can be detected by pharmacopoial methods which includes reduction of arsenic to arsenious with help of zinc and can be confirmed with mercuric chloride paper.
Classification
Elemental arsenic and arsenic sulfate and trioxide compounds are classified as "toxic" and "dangerous for the environment" in the European Union under directive 67/548/EEC.
The International Agency for Research on Cancer (IARC) recognizes arsenic and inorganic arsenic compounds as group 1 carcinogens, and the EU lists arsenic trioxide, arsenic pentoxide, and arsenate salts as category 1 carcinogens.
Arsenic is known to cause arsenicosis when present in drinking water, "the most common species being arsenate [; As(V)] and arsenite [; As(III)]".
Legal limits, food, and drink
In the United States since 2006, the maximum concentration in drinking water allowed by the Environmental Protection Agency (EPA) is 10 ppb and the FDA set the same standard in 2005 for bottled water. The Department of Environmental Protection for New Jersey set a drinking water limit of 5 ppb in 2006. The IDLH (immediately dangerous to life and health) value for arsenic metal and inorganic arsenic compounds is 5 mg/m3 (5 ppb). The Occupational Safety and Health Administration has set the permissible exposure limit (PEL) to a time-weighted average (TWA) of 0.01 mg/m3 (0.01 ppb), and the National Institute for Occupational Safety and Health (NIOSH) has set the recommended exposure limit (REL) to a 15-minute constant exposure of 0.002 mg/m3 (0.002 ppb). The PEL for organic arsenic compounds is a TWA of 0.5 mg/m3. (0.5 ppb). | Arsenic | Wikipedia | 399 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
In 2008, based on its ongoing testing of a wide variety of American foods for toxic chemicals, the U.S. Food and Drug Administration set the "level of concern" for inorganic arsenic in apple and pear juices at 23 ppb, based on non-carcinogenic effects, and began blocking importation of products in excess of this level; it also required recalls for non-conforming domestic products. In 2011, the national Dr. Oz television show broadcast a program highlighting tests performed by an independent lab hired by the producers. Though the methodology was disputed (it did not distinguish between organic and inorganic arsenic) the tests showed levels of arsenic up to 36 ppb. In response, the FDA tested the worst brand from the Dr. Oz show and found much lower levels. Ongoing testing found 95% of the apple juice samples were below the level of concern. Later testing by Consumer Reports showed inorganic arsenic at levels slightly above 10 ppb, and the organization urged parents to reduce consumption. In July 2013, on consideration of consumption by children, chronic exposure, and carcinogenic effect, the FDA established an "action level" of 10 ppb for apple juice, the same as the drinking water standard.
Concern about arsenic in rice in Bangladesh was raised in 2002, but at the time only Australia had a legal limit for food (one milligram per kilogram, or 1000 ppb). Concern was raised about people who were eating U.S. rice exceeding WHO standards for personal arsenic intake in 2005. In 2011, the People's Republic of China set a food standard of 150 ppb for arsenic.
In the United States in 2012, testing by separate groups of researchers at the Children's Environmental Health and Disease Prevention Research Center at Dartmouth College (early in the year, focusing on urinary levels in children) and Consumer Reports (in November) found levels of arsenic in rice that resulted in calls for the FDA to set limits. The FDA released some testing results in September 2012, and as of July 2013, is still collecting data in support of a new potential regulation. It has not recommended any changes in consumer behavior. | Arsenic | Wikipedia | 431 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Consumer Reports recommended:
That the EPA and FDA eliminate arsenic-containing fertilizer, drugs, and pesticides in food production;
That the FDA establish a legal limit for food;
That industry change production practices to lower arsenic levels, especially in food for children; and
That consumers test home water supplies, eat a varied diet, and cook rice with excess water, then draining it off (reducing inorganic arsenic by about one third along with a slight reduction in vitamin content).
Evidence-based public health advocates also recommend that, given the lack of regulation or labeling for arsenic in the U.S., children should eat no more than 1.5 servings per week of rice and should not drink rice milk as part of their daily diet before age 5. They also offer recommendations for adults and infants on how to limit arsenic exposure from rice, drinking water, and fruit juice.
A 2014 World Health Organization advisory conference was scheduled to consider limits of 200–300 ppb for rice.
Reducing arsenic content in rice
In 2020, scientists assessed multiple preparation procedures of rice for their capacity to reduce arsenic content and preserve nutrients, recommending a procedure involving parboiling and water-absorption.
Occupational exposure limits
Ecotoxicity
Arsenic is bioaccumulative in many organisms, marine species in particular, but it does not appear to biomagnify significantly in food webs. In polluted areas, plant growth may be affected by root uptake of arsenate, which is a phosphate analog and therefore readily transported in plant tissues and cells. In polluted areas, uptake of the more toxic arsenite ion (found more particularly in reducing conditions) is likely in poorly-drained soils.
Toxicity in animals
Biological mechanism
Arsenic's toxicity comes from the affinity of arsenic(III) oxides for thiols. Thiols, in the form of cysteine residues and cofactors such as lipoic acid and coenzyme A, are situated at the active sites of many important enzymes. | Arsenic | Wikipedia | 402 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Arsenic disrupts ATP production through several mechanisms. At the level of the citric acid cycle, arsenic inhibits lipoic acid, which is a cofactor for pyruvate dehydrogenase. By competing with phosphate, arsenate uncouples oxidative phosphorylation, thus inhibiting energy-linked reduction of NAD+, mitochondrial respiration and ATP synthesis. Hydrogen peroxide production is also increased, which, it is speculated, has potential to form reactive oxygen species and oxidative stress. These metabolic interferences lead to death from multi-system organ failure. The organ failure is presumed to be from necrotic cell death, not apoptosis, since energy reserves have been too depleted for apoptosis to occur.
Exposure risks and remediation
Occupational exposure and arsenic poisoning may occur in persons working in industries involving the use of inorganic arsenic and its compounds, such as wood preservation, glass production, nonferrous metal alloys, and electronic semiconductor manufacturing. Inorganic arsenic is also found in coke oven emissions associated with the smelter industry.
The conversion between As(III) and As(V) is a large factor in arsenic environmental contamination. According to Croal, Gralnick, Malasarn and Newman, "[the] understanding [of] what stimulates As(III) oxidation and/or limits As(V) reduction is relevant for bioremediation of contaminated sites (Croal). The study of chemolithoautotrophic As(III) oxidizers and the heterotrophic As(V) reducers can help the understanding of the oxidation and/or reduction of arsenic.
Treatment
Treatment of chronic arsenic poisoning is possible. British anti-lewisite (dimercaprol) is prescribed in doses of 5 mg/kg up to 300 mg every 4 hours for the first day, then every 6 hours for the second day, and finally every 8 hours for 8 additional days. However the USA's Agency for Toxic Substances and Disease Registry (ATSDR) states that the long-term effects of arsenic exposure cannot be predicted. Blood, urine, hair, and nails may be tested for arsenic; however, these tests cannot foresee possible health outcomes from the exposure. Long-term exposure and consequent excretion through urine has been linked to bladder and kidney cancer in addition to cancer of the liver, prostate, skin, lungs, and nasal cavity. | Arsenic | Wikipedia | 499 | 897 | https://en.wikipedia.org/wiki/Arsenic | Physical sciences | Chemical elements_2 | null |
Antimony is a chemical element; it has symbol Sb () and atomic number 51. A lustrous grey metal or metalloid, it is found in nature mainly as the sulfide mineral stibnite (Sb2S3). Antimony compounds have been known since ancient times and were powdered for use as medicine and cosmetics, often known by the Arabic name kohl. The earliest known description of this metalloid in the West was written in 1540 by Vannoccio Biringuccio.
China is the largest producer of antimony and its compounds, with most production coming from the Xikuangshan Mine in Hunan. The industrial methods for refining antimony from stibnite are roasting followed by reduction with carbon, or direct reduction of stibnite with iron.
The most common applications for metallic antimony are in alloys with lead and tin, which have improved properties for solders, bullets, and plain bearings. It improves the rigidity of lead-alloy plates in lead–acid batteries. Antimony trioxide is a prominent additive for halogen-containing flame retardants. Antimony is used as a dopant in semiconductor devices.
Characteristics
Properties
Antimony is a member of group 15 of the periodic table, one of the elements called pnictogens, and has an electronegativity of 2.05. In accordance with periodic trends, it is more electronegative than tin or bismuth, and less electronegative than tellurium or arsenic. Antimony is stable in air at room temperature but, if heated, it reacts with oxygen to produce antimony trioxide, Sb2O3.
Antimony is a silvery, lustrous gray metalloid with a Mohs scale hardness of 3, which is too soft to mark hard objects. Coins of antimony were issued in China's Guizhou in 1931; durability was poor, and minting was soon discontinued because of its softness and toxicity. Antimony is resistant to attack by acids. | Antimony | Wikipedia | 409 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
The only stable allotrope of antimony under standard conditions is metallic, brittle, silver-white, and shiny. It crystallises in a trigonal cell, isomorphic with bismuth and the gray allotrope of arsenic, and is formed when molten antimony is cooled slowly. Amorphous black antimony is formed upon rapid cooling of antimony vapor, and is only stable as a thin film (thickness in nanometres); thicker samples spontaneously transform into the metallic form. It oxidizes in air and may ignite spontaneously. At 100 °C, it gradually transforms into the stable form. The supposed yellow allotrope of antimony, generated only by oxidation of stibine (SbH3) at −90 °C, is also impure and not a true allotrope; above this temperature and in ambient light, it transforms into the more stable black allotrope. A rare explosive form of antimony can be formed from the electrolysis of antimony trichloride, but it always contains appreciable chlorine and is not really an antimony allotrope. When scratched with a sharp implement, an exothermic reaction occurs and white fumes are given off as metallic antimony forms; when rubbed with a pestle in a mortar, a strong detonation occurs.
Elemental antimony adopts a layered structure (space group Rm No. 166) whose layers consist of fused, ruffled, six-membered rings. The nearest and next-nearest neighbors form an irregular octahedral complex, with the three atoms in each double layer slightly closer than the three atoms in the next. This relatively close packing leads to a high density of 6.697 g/cm3, but the weak bonding between the layers leads to the low hardness and brittleness of antimony.
Isotopes | Antimony | Wikipedia | 383 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
Antimony has two stable isotopes: 121Sb with a natural abundance of 57.36% and 123Sb with a natural abundance of 42.64%. It also has 35 radioisotopes, of which the longest-lived is 125Sb with a half-life of 2.75 years. In addition, 29 metastable states have been characterized. The most stable of these is 120m1Sb with a half-life of 5.76 days. Isotopes that are lighter than the stable 123Sb tend to decay by β+ decay, and those that are heavier tend to decay by β− decay, with some exceptions. Antimony is the lightest element to have an isotope with an alpha decay branch, excluding 8Be and other light nuclides with beta-delayed alpha emission.
Occurrence
The abundance of antimony in the Earth's crust is estimated at 0.2 parts per million, comparable to thallium at 0.5 ppm and silver at 0.07 ppm. It is the 63rd most abundant element in the crust. Even though this element is not abundant, it is found in more than 100 mineral species. Antimony is sometimes found natively (e.g. on Antimony Peak), but more frequently it is found in the sulfide stibnite (Sb2S3) which is the predominant ore mineral.
Compounds
Antimony compounds are often classified according to their oxidation state: Sb(III) and Sb(V). The +5 oxidation state is more common.
Oxides and hydroxides
Antimony trioxide is formed when antimony is burnt in air. In the gas phase, the molecule of the compound is , but it polymerizes upon condensing. Antimony pentoxide () can be formed only by oxidation with concentrated nitric acid. Antimony also forms a mixed-valence oxide, antimony tetroxide (), which features both Sb(III) and Sb(V). Unlike oxides of phosphorus and arsenic, these oxides are amphoteric, do not form well-defined oxoacids, and react with acids to form antimony salts. | Antimony | Wikipedia | 440 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
Antimonous acid is unknown, but the conjugate base sodium antimonite () forms upon fusing sodium oxide and . Transition metal antimonites are also known. Antimonic acid exists only as the hydrate , forming salts as the antimonate anion . When a solution containing this anion is dehydrated, the precipitate contains mixed oxides.
The most important antimony ore is stibnite (). Other sulfide minerals include pyrargyrite (), zinkenite, jamesonite, and boulangerite. Antimony pentasulfide is non-stoichiometric, which features antimony in the +3 oxidation state and S–S bonds. Several thioantimonides are known, such as and .
Halides
Antimony forms two series of halides: and . The trihalides , , , and are all molecular compounds having trigonal pyramidal molecular geometry.
The trifluoride is prepared by the reaction of with HF:
+ 6 HF → 2 + 3
It is Lewis acidic and readily accepts fluoride ions to form the complex anions and . Molten is a weak electrical conductor. The trichloride is prepared by dissolving in hydrochloric acid:
+ 6 HCl → 2 + 3
Arsenic sulfides are not readily attacked by the hydrochloric acid, so this method offers a route to As-free Sb.
The pentahalides and have trigonal bipyramidal molecular geometry in the gas phase, but in the liquid phase, is polymeric, whereas is monomeric. is a powerful Lewis acid used to make the superacid fluoroantimonic acid ("H2SbF7").
Oxyhalides are more common for antimony than for arsenic and phosphorus. Antimony trioxide dissolves in concentrated acid to form oxoantimonyl compounds such as SbOCl and . | Antimony | Wikipedia | 403 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
Antimonides, hydrides, and organoantimony compounds
Compounds in this class generally are described as derivatives of Sb3−. Antimony forms antimonides with metals, such as indium antimonide (InSb) and silver antimonide (). The alkali metal and zinc antimonides, such as Na3Sb and Zn3Sb2, are more reactive. Treating these antimonides with acid produces the highly unstable gas stibine, :
+ 3 →
Stibine can also be produced by treating salts with hydride reagents such as sodium borohydride. Stibine decomposes spontaneously at room temperature. Because stibine has a positive heat of formation, it is thermodynamically unstable and thus antimony does not react with hydrogen directly.
Organoantimony compounds are typically prepared by alkylation of antimony halides with Grignard reagents. A large variety of compounds are known with both Sb(III) and Sb(V) centers, including mixed chloro-organic derivatives, anions, and cations. Examples include triphenylstibine (Sb(C6H5)3) and pentaphenylantimony (Sb(C6H5)5).
History
Antimony(III) sulfide, Sb2S3, was recognized in predynastic Egypt as an eye cosmetic (kohl) as early as about 3100 BC, when the cosmetic palette was invented.
An artifact, said to be part of a vase, made of antimony dating to about 3000 BC was found at Telloh, Chaldea (part of present-day Iraq), and a copper object plated with antimony dating between 2500 BC and 2200 BC has been found in Egypt. Austen, at a lecture by Herbert Gladstone in 1892, commented that "we only know of antimony at the present day as a highly brittle and crystalline metal, which could hardly be fashioned into a useful vase, and therefore this remarkable 'find' (artifact mentioned above) must represent the lost art of rendering antimony malleable." | Antimony | Wikipedia | 442 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
The British archaeologist Roger Moorey was unconvinced the artifact was indeed a vase, mentioning that Selimkhanov, after his analysis of the Tello object (published in 1975), "attempted to relate the metal to Transcaucasian natural antimony" (i.e. native metal) and that "the antimony objects from Transcaucasia are all small personal ornaments." This weakens the evidence for a lost art "of rendering antimony malleable".
The Roman scholar Pliny the Elder described several ways of preparing antimony sulfide for medical purposes in his treatise Natural History, around 77 AD. Pliny the Elder also made a distinction between "male" and "female" forms of antimony; the male form is probably the sulfide, while the female form, which is superior, heavier, and less friable, has been suspected to be native metallic antimony.
The Greek naturalist Pedanius Dioscorides mentioned that antimony sulfide could be roasted by heating by a current of air. It is thought that this produced metallic antimony.
Antimony was frequently described in alchemical manuscripts, including the Summa Perfectionis of Pseudo-Geber, written around the 14th century. A description of a procedure for isolating antimony is later given in the 1540 book De la pirotechnia by Vannoccio Biringuccio, predating the more famous 1556 book by Agricola, De re metallica. In this context Agricola has been often incorrectly credited with the discovery of metallic antimony. The book Currus Triumphalis Antimonii (The Triumphal Chariot of Antimony), describing the preparation of metallic antimony, was published in Germany in 1604. It was purported to be written by a Benedictine monk, writing under the name Basilius Valentinus in the 15th century; if it were authentic, which it is not, it would predate Biringuccio.
The metal antimony was known to German chemist Andreas Libavius in 1615 who obtained it by adding iron to a molten mixture of antimony sulfide, salt and potassium tartrate. This procedure produced antimony with a crystalline or starred surface.
With the advent of challenges to phlogiston theory, it was recognized that antimony is an element forming sulfides, oxides, and other compounds, as do other metals. | Antimony | Wikipedia | 486 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
The first discovery of naturally occurring pure antimony in the Earth's crust was described by the Swedish scientist and local mine district engineer Anton von Swab in 1783; the type-sample was collected from the Sala Silver Mine in the Bergslagen mining district of Sala, Västmanland, Sweden.
Etymology
The medieval Latin form, from which the modern languages and late Byzantine Greek take their names for antimony, is . The origin of that is uncertain, and all suggestions have some difficulty either of form or interpretation. The popular etymology, from ἀντίμοναχός anti-monachos or French , would mean "monk-killer", which is explained by the fact that many early alchemists were monks, and some antimony compounds were poisonous.
Another popular etymology is the hypothetical Greek word ἀντίμόνος antimonos, "against aloneness", explained as "not found as metal", or "not found unalloyed". However, ancient Greek would more naturally express the pure negative as α- ("not"). Edmund Oscar von Lippmann conjectured a hypothetical Greek word ανθήμόνιον anthemonion, which would mean "floret", and cites several examples of related Greek words (but not that one) which describe chemical or biological efflorescence.
The early uses of antimonium include the translations, in 1050–1100, by Constantine the African of Arabic medical treatises. Several authorities believe antimonium is a scribal corruption of some Arabic form; Meyerhof derives it from ithmid; other possibilities include athimar, the Arabic name of the metalloid, and a hypothetical as-stimmi, derived from or parallel to the Greek.
The standard chemical symbol for antimony (Sb) is credited to Jöns Jakob Berzelius, who derived the abbreviation from stibium.
The ancient words for antimony mostly have, as their chief meaning, kohl, the sulfide of antimony.
The Egyptians called antimony mśdmt or stm. | Antimony | Wikipedia | 433 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
The Arabic word for the substance, as opposed to the cosmetic, can appear as ithmid, athmoud, othmod, or uthmod. Littré suggests the first form, which is the earliest, derives from stimmida, an accusative for stimmi. The Greek word στίμμι (stimmi) is used by Attic tragic poets of the 5th century BC, and is possibly a loan word from Arabic or from Egyptian stm.
Production
Process
The extraction of antimony from ores depends on the quality and composition of the ore. Most antimony is mined as the sulfide; lower-grade ores are concentrated by froth flotation, while higher-grade ores are heated to 500–600 °C, the temperature at which stibnite melts and separates from the gangue minerals. Antimony can be isolated from the crude antimony sulfide by reduction with scrap iron:
+ 3 Fe → 2 Sb + 3 FeS
The sulfide is converted to an oxide by roasting. The product is further purified by vaporizing the volatile antimony(III) oxide, which is recovered. This sublimate is often used directly for the main applications, impurities being arsenic and sulfide. Antimony is isolated from the oxide by a carbothermal reduction:
2 + 3 C → 4 Sb + 3
The lower-grade ores are reduced in blast furnaces while the higher-grade ores are reduced in reverberatory furnaces.
Top producers and production volumes
In 2022, according to the US Geological Survey, China accounted for 54.5% of total antimony production, followed in second place by Russia with 18.2% and Tajikistan with 15.5%.
Chinese production of antimony is expected to decline in the future as mines and smelters are closed down by the government as part of pollution control. Especially due to an environmental protection law having gone into effect in January 2015 and revised "Emission Standards of Pollutants for Stanum, Antimony, and Mercury" having gone into effect, hurdles for economic production are higher.
Reported production of antimony in China has fallen and is unlikely to increase in the coming years, according to the Roskill report. No significant antimony deposits in China have been developed for about ten years, and the remaining economic reserves are being rapidly depleted.
Reserves | Antimony | Wikipedia | 489 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
Supply risk
For antimony-importing regions, such as Europe and the U.S., antimony is considered to be a critical mineral for industrial manufacturing that is at risk of supply chain disruption. With global production coming mainly from China (74%), Tajikistan (8%), and Russia (4%), these sources are critical to supply.
European Union: Antimony is considered a critical raw material for defense, automotive, construction and textiles. The E.U. sources are 100% imported, coming mainly from Turkey (62%), Bolivia (20%) and Guatemala (7%).
United Kingdom: The British Geological Survey's 2015 risk list ranks antimony second highest (after rare earth elements) on the relative supply risk index.
United States: Antimony is a mineral commodity considered critical to the economic and national security. In 2022, no antimony was mined in the U.S.
Applications
Approximately 48% of antimony is consumed in flame retardants, 33% in lead–acid batteries, and 8% in plastics.
Flame retardants
Antimony is mainly used as the trioxide for flame-proofing compounds, always in combination with halogenated flame retardants except in halogen-containing polymers. The flame retarding effect of antimony trioxide is produced by the formation of halogenated antimony compounds, which react with hydrogen atoms, and probably also with oxygen atoms and OH radicals, thus inhibiting fire. Markets for these flame-retardants include children's clothing, toys, aircraft, and automobile seat covers. They are also added to polyester resins in fiberglass composites for such items as light aircraft engine covers. The resin will burn in the presence of an externally generated flame, but will extinguish when the external flame is removed.
Alloys | Antimony | Wikipedia | 378 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
Antimony forms a highly useful alloy with lead, increasing its hardness and mechanical strength. When casting it increases fluidity of the melt and reduces shrinkage during cooling. For most applications involving lead, varying amounts of antimony are used as alloying metal. In lead–acid batteries, this addition improves plate strength and charging characteristics. For sailboats, lead keels are used to provide righting moment, ranging from 600 lbs to over 200 tons for the largest sailing superyachts; to improve hardness and tensile strength of the lead keel, antimony is mixed with lead between 2% and 5% by volume. Antimony is used in antifriction alloys (such as Babbitt metal), in bullets and lead shot, electrical cable sheathing, type metal (for example, for linotype printing machines), solder (some "lead-free" solders contain 5% Sb), in pewter, and in hardening alloys with low tin content in the manufacturing of organ pipes.
Other applications
Three other applications consume nearly all the rest of the world's supply. One application is as a stabilizer and catalyst for the production of polyethylene terephthalate. Another is as a fining agent to remove microscopic bubbles in glass, mostly for TV screens antimony ions interact with oxygen, suppressing the tendency of the latter to form bubbles. The third application is pigments.
In the 1990s antimony was increasingly being used in semiconductors as a dopant in n-type silicon wafers for diodes, infrared detectors, and Hall-effect devices. In the 1950s, the emitters and collectors of n-p-n alloy junction transistors were doped with tiny beads of a lead-antimony alloy. Indium antimonide (InSb) is used as a material for mid-infrared detectors.
The material Ge2Sb2Te5 is used as for phase-change memory, a type of computer memory. | Antimony | Wikipedia | 401 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
Biology and medicine have few uses for antimony. Treatments containing antimony, known as antimonials, are used as emetics. Antimony compounds are used as antiprotozoan drugs. Potassium antimonyl tartrate, or tartar emetic, was once used as an anti-schistosomal drug from 1919 on. It was subsequently replaced by praziquantel. Antimony and its compounds are used in several veterinary preparations, such as anthiomaline and lithium antimony thiomalate, as a skin conditioner in ruminants. Antimony has a nourishing or conditioning effect on keratinized tissues in animals.
Antimony-based drugs, such as meglumine antimoniate, are also considered the drugs of choice for treatment of leishmaniasis. Early treatments used antimony(III) species (trivalent antimonials), but in 1922 Upendranath Brahmachari invented a much safer antimony(V) drug, and since then so-called pentavalent antimonials have been the standard first-line treatment. However, Leishmania strains in Bihar and neighboring regions have developed resistance to antimony. Elemental antimony as an antimony pill was once used as a medicine. It could be reused by others after ingestion and elimination.
Antimony(III) sulfide is used in the heads of some safety matches. Antimony sulfides help to stabilize the friction coefficient in automotive brake pad materials. Antimony is used in bullets, bullet tracers, paint, glass art, and as an opacifier in enamel. Antimony-124 is used together with beryllium in neutron sources; the gamma rays emitted by antimony-124 initiate the photodisintegration of beryllium. The emitted neutrons have an average energy of 24 keV. Natural antimony is used in startup neutron sources.
The powder derived from crushed antimony sulfide (kohl) has been used for millennia as an eye cosmetic. Historically it was applied to the eyes with a metal rod and with one's spittle, and was thought by the ancients to aid in curing eye infections. The practice is still seen in Yemen and in other Muslim countries.
Precautions | Antimony | Wikipedia | 463 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
Antimony and many of its compounds are toxic, and the effects of antimony poisoning are similar to arsenic poisoning. The toxicity of antimony is far lower than that of arsenic; this might be caused by the significant differences of uptake, metabolism and excretion between arsenic and antimony. The uptake of antimony(III) or antimony(V) in the gastrointestinal tract is at most 20%. Antimony(V) is not quantitatively reduced to antimony(III) in the cell (in fact antimony(III) is oxidised to antimony(V) instead).
Since methylation of antimony does not occur, the excretion of antimony(V) in urine is the main way of elimination. Like arsenic, the most serious effect of acute antimony poisoning is cardiotoxicity and the resulting myocarditis; however, it can also manifest as Adams–Stokes syndrome, which arsenic does not. Reported cases of intoxication by antimony equivalent to 90 mg antimony potassium tartrate dissolved from enamel has been reported to show only short term effects. An intoxication with 6 g of antimony potassium tartrate was reported to result in death after three days.
Inhalation of antimony dust is harmful and in certain cases may be fatal; in small doses, antimony causes headaches, dizziness, and depression. Larger doses such as prolonged skin contact may cause dermatitis, or damage the kidneys and the liver, causing violent and frequent vomiting, leading to death in a few days.
Antimony is incompatible with strong oxidizing agents, strong acids, halogen acids, chlorine, or fluorine. It should be kept away from heat.
Antimony leaches from polyethylene terephthalate (PET) bottles into liquids. While levels observed for bottled water are below drinking water guidelines, fruit juice concentrates (for which no guidelines are established) produced in the UK were found to contain up to 44.7 μg/L of antimony, well above the EU limits for tap water of 5 μg/L. The guidelines are:
World Health Organization: 20 μg/L
Japan: 15 μg/L
United States Environmental Protection Agency, Health Canada and the Ontario Ministry of Environment: 6 μg/L
EU and German Federal Ministry of Environment: 5 μg/L | Antimony | Wikipedia | 493 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
The tolerable daily intake (TDI) proposed by WHO is 6 μg antimony per kilogram of body weight. The immediately dangerous to life or health (IDLH) value for antimony is 50 mg/m3.
Toxicity
Certain compounds of antimony appear to be toxic, particularly antimony trioxide and antimony potassium tartrate. Effects may be similar to arsenic poisoning. Occupational exposure may cause respiratory irritation, pneumoconiosis, antimony spots on the skin, gastrointestinal symptoms, and cardiac arrhythmias. In addition, antimony trioxide is potentially carcinogenic to humans.
Adverse health effects have been observed in humans and animals following inhalation, oral, or dermal exposure to antimony and antimony compounds. Antimony toxicity typically occurs either due to occupational exposure, during therapy or from accidental ingestion. It is unclear if antimony can enter the body through the skin. The presence of low levels of antimony in saliva may also be associated with dental decay. | Antimony | Wikipedia | 211 | 898 | https://en.wikipedia.org/wiki/Antimony | Physical sciences | Chemical elements_2 | null |
Actinium is a chemical element; it has symbol Ac and atomic number 89. It was first isolated by Friedrich Oskar Giesel in 1902, who gave it the name emanium; the element got its name by being wrongly identified with a substance André-Louis Debierne found in 1899 and called actinium. The actinide series, a set of 15 elements between actinium and lawrencium in the periodic table, are named for actinium. Together with polonium, radium, and radon, actinium was one of the first non-primordial radioactive elements to be isolated.
A soft, silvery-white radioactive metal, actinium reacts rapidly with oxygen and moisture in air forming a white coating of actinium oxide that prevents further oxidation. As with most lanthanides and many actinides, actinium assumes oxidation state +3 in nearly all its chemical compounds. Actinium is found only in traces in uranium and thorium ores as the isotope 227Ac, which decays with a half-life of 21.772 years, predominantly emitting beta and sometimes alpha particles, and 228Ac, which is beta active with a half-life of 6.15 hours. One tonne of natural uranium in ore contains about 0.2 milligrams of actinium-227, and one tonne of thorium contains about 5 nanograms of actinium-228. The close similarity of physical and chemical properties of actinium and lanthanum makes separation of actinium from the ore impractical. Instead, the element is prepared, in milligram amounts, by the neutron irradiation of in a nuclear reactor. Owing to its scarcity, high price and radioactivity, actinium has no significant industrial use. Its current applications include a neutron source and an agent for radiation therapy. | Actinium | Wikipedia | 374 | 899 | https://en.wikipedia.org/wiki/Actinium | Physical sciences | Chemical elements_2 | null |
History
André-Louis Debierne, a French chemist, announced the discovery of a new element in 1899. He separated it from pitchblende residues left by Marie and Pierre Curie after they had extracted radium. In 1899, Debierne described the substance as similar to titanium and (in 1900) as similar to thorium. Friedrich Oskar Giesel found in 1902 a substance similar to lanthanum and called it "emanium" in 1904. After a comparison of the substances' half-lives determined by Debierne, Harriet Brooks in 1904, and Otto Hahn and Otto Sackur in 1905, Debierne's chosen name for the new element was retained because it had seniority, despite the contradicting chemical properties he claimed for the element at different times.
Articles published in the 1970s and later suggest that Debierne's results published in 1904 conflict with those reported in 1899 and 1900. Furthermore, the now-known chemistry of actinium precludes its presence as anything other than a minor constituent of Debierne's 1899 and 1900 results; in fact, the chemical properties he reported make it likely that he had, instead, accidentally identified protactinium, which would not be discovered for another fourteen years, only to have it disappear due to its hydrolysis and adsorption onto his laboratory equipment. This has led some authors to advocate that Giesel alone should be credited with the discovery. A less confrontational vision of scientific discovery is proposed by Adloff. He suggests that hindsight criticism of the early publications should be mitigated by the then nascent state of radiochemistry: highlighting the prudence of Debierne's claims in the original papers, he notes that nobody can contend that Debierne's substance did not contain actinium. Debierne, who is now considered by the vast majority of historians as the discoverer, lost interest in the element and left the topic. Giesel, on the other hand, can rightfully be credited with the first preparation of radiochemically pure actinium and with the identification of its atomic number 89.
The name actinium originates from the Ancient Greek aktis, aktinos (ακτίς, ακτίνος), meaning beam or ray. Its symbol Ac is also used in abbreviations of other compounds that have nothing to do with actinium, such as acetyl, acetate and sometimes acetaldehyde. | Actinium | Wikipedia | 503 | 899 | https://en.wikipedia.org/wiki/Actinium | Physical sciences | Chemical elements_2 | null |
Properties
Actinium is a soft, silvery-white, radioactive, metallic element. Its estimated shear modulus is similar to that of lead. Owing to its strong radioactivity, actinium glows in the dark with a pale blue light, which originates from the surrounding air ionized by the emitted energetic particles. Actinium has similar chemical properties to lanthanum and other lanthanides, and therefore these elements are difficult to separate when extracting from uranium ores. Solvent extraction and ion chromatography are commonly used for the separation.
The first element of the actinides, actinium gave the set its name, much as lanthanum had done for the lanthanides. The actinides are much more diverse than the lanthanides and therefore it was not until 1945 that the most significant change to Dmitri Mendeleev's periodic table since the recognition of the lanthanides, the introduction of the actinides, was generally accepted after Glenn T. Seaborg's research on the transuranium elements (although it had been proposed as early as 1892 by British chemist Henry Bassett).
Actinium reacts rapidly with oxygen and moisture in air forming a white coating of actinium oxide that impedes further oxidation. As with most lanthanides and actinides, actinium exists in the oxidation state +3, and the Ac3+ ions are colorless in solutions. The oxidation state +3 originates from the [Rn] 6d17s2 electronic configuration of actinium, with three valence electrons that are easily donated to give the stable closed-shell structure of the noble gas radon. Although the 5f orbitals are unoccupied in an actinium atom, it can be used as a valence orbital in actinium complexes and hence it is generally considered the first 5f element by authors working on it. Ac3+ is the largest of all known tripositive ions and its first coordination sphere contains approximately 10.9 ± 0.5 water molecules.
Chemical compounds | Actinium | Wikipedia | 409 | 899 | https://en.wikipedia.org/wiki/Actinium | Physical sciences | Chemical elements_2 | null |
Due to actinium's intense radioactivity, only a limited number of actinium compounds are known. These include: AcF3, AcCl3, AcBr3, AcOF, AcOCl, AcOBr, Ac2S3, Ac2O3, AcPO4 and Ac(NO3)3. They all contain actinium in the oxidation state +3. In particular, the lattice constants of the analogous lanthanum and actinium compounds differ by only a few percent.
Here a, b and c are lattice constants, No is space group number and Z is the number of formula units per unit cell. Density was not measured directly but calculated from the lattice parameters.
Oxides
Actinium oxide (Ac2O3) can be obtained by heating the hydroxide at or the oxalate at , in vacuum. Its crystal lattice is isotypic with the oxides of most trivalent rare-earth metals.
Halides
Actinium trifluoride can be produced either in solution or in solid reaction. The former reaction is carried out at room temperature, by adding hydrofluoric acid to a solution containing actinium ions. In the latter method, actinium metal is treated with hydrogen fluoride vapors at in an all-platinum setup. Treating actinium trifluoride with ammonium hydroxide at yields oxyfluoride AcOF. Whereas lanthanum oxyfluoride can be easily obtained by burning lanthanum trifluoride in air at for an hour, similar treatment of actinium trifluoride yields no AcOF and only results in melting of the initial product.
AcF3 + 2 NH3 + H2O → AcOF + 2 NH4F
Actinium trichloride is obtained by reacting actinium hydroxide or oxalate with carbon tetrachloride vapors at temperatures above . Similarly to the oxyfluoride, actinium oxychloride can be prepared by hydrolyzing actinium trichloride with ammonium hydroxide at . However, in contrast to the oxyfluoride, the oxychloride could well be synthesized by igniting a solution of actinium trichloride in hydrochloric acid with ammonia.
Reaction of aluminium bromide and actinium oxide yields actinium tribromide:
Ac2O3 + 2 AlBr3 → 2 AcBr3 + Al2O3 | Actinium | Wikipedia | 499 | 899 | https://en.wikipedia.org/wiki/Actinium | Physical sciences | Chemical elements_2 | null |
and treating it with ammonium hydroxide at results in the oxybromide AcOBr.
Other compounds
Actinium hydride was obtained by reduction of actinium trichloride with potassium at , and its structure was deduced by analogy with the corresponding LaH2 hydride. The source of hydrogen in the reaction was uncertain.
Mixing monosodium phosphate (NaH2PO4) with a solution of actinium in hydrochloric acid yields white-colored actinium phosphate hemihydrate (AcPO4·0.5H2O), and heating actinium oxalate with hydrogen sulfide vapors at for a few minutes results in a black actinium sulfide Ac2S3. It may possibly be produced by acting with a mixture of hydrogen sulfide and carbon disulfide on actinium oxide at .
Isotopes
Naturally occurring actinium is principally composed of two radioactive isotopes; (from the radioactive family of ) and (a granddaughter of ). decays mainly as a beta emitter with a very small energy, but in 1.38% of cases it emits an alpha particle, so it can readily be identified through alpha spectrometry. Thirty-three radioisotopes have been identified, the most stable being with a half-life of 21.772 years, with a half-life of 10.0 days and with a half-life of 29.37 hours. All remaining radioactive isotopes have half-lives that are less than 10 hours and the majority of them have half-lives shorter than one minute. The shortest-lived known isotope of actinium is (half-life of 69 nanoseconds) which decays through alpha decay. Actinium also has two known meta states. The most significant isotopes for chemistry are 225Ac, 227Ac, and 228Ac.
Purified comes into equilibrium with its decay products after about a half of year. It decays according to its 21.772-year half-life emitting mostly beta (98.62%) and some alpha particles (1.38%); the successive decay products are part of the actinium series. Owing to the low available amounts, low energy of its beta particles (maximum 44.8 keV) and low intensity of alpha radiation, is difficult to detect directly by its emission and it is therefore traced via its decay products. The isotopes of actinium range in atomic weight from 203 u () to 236 u ().
Occurrence and synthesis | Actinium | Wikipedia | 511 | 899 | https://en.wikipedia.org/wiki/Actinium | Physical sciences | Chemical elements_2 | null |
Actinium is found only in traces in uranium ores – one tonne of uranium in ore contains about 0.2 milligrams of 227Ac – and in thorium ores, which contain about 5 nanograms of 228Ac per one tonne of thorium. The actinium isotope 227Ac is a transient member of the uranium-actinium series decay chain, which begins with the parent isotope 235U (or 239Pu) and ends with the stable lead isotope 207Pb. The isotope 228Ac is a transient member of the thorium series decay chain, which begins with the parent isotope 232Th and ends with the stable lead isotope 208Pb. Another actinium isotope (225Ac) is transiently present in the neptunium series decay chain, beginning with 237Np (or 233U) and ending with thallium (205Tl) and near-stable bismuth (209Bi); even though all primordial 237Np has decayed away, it is continuously produced by neutron knock-out reactions on natural 238U. | Actinium | Wikipedia | 217 | 899 | https://en.wikipedia.org/wiki/Actinium | Physical sciences | Chemical elements_2 | null |
The low natural concentration, and the close similarity of physical and chemical properties to those of lanthanum and other lanthanides, which are always abundant in actinium-bearing ores, render separation of actinium from the ore impractical. The most concentrated actinium sample prepared from raw material consisted of 7 micrograms of 227Ac in less than 0.1 milligrams of La2O3, and complete separation was never achieved. Instead, actinium is prepared, in milligram amounts, by the neutron irradiation of in a nuclear reactor.
^{226}_{88}Ra + ^{1}_{0}n -> ^{227}_{88}Ra ->[\beta^-][42.2 \ \ce{min}] ^{227}_{89}Ac
The reaction yield is about 2% of the radium weight. 227Ac can further capture neutrons resulting in small amounts of 228Ac. After the synthesis, actinium is separated from radium and from the products of decay and nuclear fusion, such as thorium, polonium, lead and bismuth. The extraction can be performed with thenoyltrifluoroacetone-benzene solution from an aqueous solution of the radiation products, and the selectivity to a certain element is achieved by adjusting the pH (to about 6.0 for actinium). An alternative procedure is anion exchange with an appropriate resin in nitric acid, which can result in a separation factor of 1,000,000 for radium and actinium vs. thorium in a two-stage process. Actinium can then be separated from radium, with a ratio of about 100, using a low cross-linking cation exchange resin and nitric acid as eluant.
225Ac was first produced artificially at the Institute for Transuranium Elements (ITU) in Germany using a cyclotron and at St George Hospital in Sydney using a linac in 2000. This rare isotope has potential applications in radiation therapy and is most efficiently produced by bombarding a radium-226 target with 20–30 MeV deuterium ions. This reaction also yields 226Ac which however decays with a half-life of 29 hours and thus does not contaminate 225Ac. | Actinium | Wikipedia | 478 | 899 | https://en.wikipedia.org/wiki/Actinium | Physical sciences | Chemical elements_2 | null |
Actinium metal has been prepared by the reduction of actinium fluoride with lithium vapor in vacuum at a temperature between . Higher temperatures resulted in evaporation of the product and lower ones lead to an incomplete transformation. Lithium was chosen among other alkali metals because its fluoride is most volatile.
Applications
Owing to its scarcity, high price and radioactivity, 227Ac currently has no significant industrial use, but 225Ac is currently being studied for use in cancer treatments such as targeted alpha therapies.
227Ac is highly radioactive and was therefore studied for use as an active element of radioisotope thermoelectric generators, for example in spacecraft. The oxide of 227Ac pressed with beryllium is also an efficient neutron source with the activity exceeding that of the standard americium-beryllium and radium-beryllium pairs. In all those applications, 227Ac (a beta source) is merely a progenitor which generates alpha-emitting isotopes upon its decay. Beryllium captures alpha particles and emits neutrons owing to its large cross-section for the (α,n) nuclear reaction:
^{9}_{4}Be + ^{4}_{2}He -> ^{12}_{6}C + ^{1}_{0}n + \gamma
The 227AcBe neutron sources can be applied in a neutron probe – a standard device for measuring the quantity of water present in soil, as well as moisture/density for quality control in highway construction. Such probes are also used in well logging applications, in neutron radiography, tomography and other radiochemical investigations. | Actinium | Wikipedia | 341 | 899 | https://en.wikipedia.org/wiki/Actinium | Physical sciences | Chemical elements_2 | null |
225Ac is applied in medicine to produce in a reusable generator or can be used alone as an agent for radiation therapy, in particular targeted alpha therapy (TAT). This isotope has a half-life of 10 days, making it much more suitable for radiation therapy than 213Bi (half-life 46 minutes). Additionally, 225Ac decays to nontoxic 209Bi rather than toxic lead, which is the final product in the decay chains of several other candidate isotopes, namely 227Th, 228Th, and 230U. Not only 225Ac itself, but also its daughters, emit alpha particles which kill cancer cells in the body. The major difficulty with application of 225Ac was that intravenous injection of simple actinium complexes resulted in their accumulation in the bones and liver for a period of tens of years. As a result, after the cancer cells were quickly killed by alpha particles from 225Ac, the radiation from the actinium and its daughters might induce new mutations. To solve this problem, 225Ac was bound to a chelating agent, such as citrate, ethylenediaminetetraacetic acid (EDTA) or diethylene triamine pentaacetic acid (DTPA). This reduced actinium accumulation in the bones, but the excretion from the body remained slow. Much better results were obtained with such chelating agents as HEHA () or DOTA () coupled to trastuzumab, a monoclonal antibody that interferes with the HER2/neu receptor. The latter delivery combination was tested on mice and proved to be effective against leukemia, lymphoma, breast, ovarian, neuroblastoma and prostate cancers. | Actinium | Wikipedia | 354 | 899 | https://en.wikipedia.org/wiki/Actinium | Physical sciences | Chemical elements_2 | null |
The medium half-life of 227Ac (21.77 years) makes it a very convenient radioactive isotope in modeling the slow vertical mixing of oceanic waters. The associated processes cannot be studied with the required accuracy by direct measurements of current velocities (of the order 50 meters per year). However, evaluation of the concentration depth-profiles for different isotopes allows estimating the mixing rates. The physics behind this method is as follows: oceanic waters contain homogeneously dispersed 235U. Its decay product, 231Pa, gradually precipitates to the bottom, so that its concentration first increases with depth and then stays nearly constant. 231Pa decays to 227Ac; however, the concentration of the latter isotope does not follow the 231Pa depth profile, but instead increases toward the sea bottom. This occurs because of the mixing processes which raise some additional 227Ac from the sea bottom. Thus analysis of both 231Pa and 227Ac depth profiles allows researchers to model the mixing behavior.
There are theoretical predictions that AcHx hydrides (in this case with very high pressure) are a candidate for a near room-temperature superconductor as they have Tc significantly higher than H3S, possibly near 250 K.
Precautions
227Ac is highly radioactive and experiments with it are carried out in a specially designed laboratory equipped with a tight glove box. When actinium trichloride is administered intravenously to rats, about 33% of actinium is deposited into the bones and 50% into the liver. Its toxicity is comparable to, but slightly lower, than that of americium and plutonium. For trace quantities, fume hoods with good aeration suffice; for gram amounts, hot cells with shielding from the intense gamma radiation emitted by 227Ac are necessary. | Actinium | Wikipedia | 362 | 899 | https://en.wikipedia.org/wiki/Actinium | Physical sciences | Chemical elements_2 | null |
Americium is a synthetic chemical element; it has symbol Am and atomic number 95. It is radioactive and a transuranic member of the actinide series in the periodic table, located under the lanthanide element europium and was thus named after the Americas by analogy.
Americium was first produced in 1944 by the group of Glenn T. Seaborg from Berkeley, California, at the Metallurgical Laboratory of the University of Chicago, as part of the Manhattan Project. Although it is the third element in the transuranic series, it was discovered fourth, after the heavier curium. The discovery was kept secret and only released to the public in November 1945. Most americium is produced by uranium or plutonium being bombarded with neutrons in nuclear reactors – one tonne of spent nuclear fuel contains about 100 grams of americium. It is widely used in commercial ionization chamber smoke detectors, as well as in neutron sources and industrial gauges. Several unusual applications, such as nuclear batteries or fuel for space ships with nuclear propulsion, have been proposed for the isotope 242mAm, but they are as yet hindered by the scarcity and high price of this nuclear isomer.
Americium is a relatively soft radioactive metal with a silvery appearance. Its most common isotopes are 241Am and 243Am. In chemical compounds, americium usually assumes the oxidation state +3, especially in solutions. Several other oxidation states are known, ranging from +2 to +7, and can be identified by their characteristic optical absorption spectra. The crystal lattices of solid americium and its compounds contain small intrinsic radiogenic defects, due to metamictization induced by self-irradiation with alpha particles, which accumulates with time; this can cause a drift of some material properties over time, more noticeable in older samples.
History | Americium | Wikipedia | 378 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
Although americium was likely produced in previous nuclear experiments, it was first intentionally synthesized, isolated and identified in late autumn 1944, at the University of California, Berkeley, by Glenn T. Seaborg, Leon O. Morgan, Ralph A. James, and Albert Ghiorso. They used a 60-inch cyclotron at the University of California, Berkeley. The element was chemically identified at the Metallurgical Laboratory (now Argonne National Laboratory) of the University of Chicago. Following the lighter neptunium, plutonium, and heavier curium, americium was the fourth transuranium element to be discovered. At the time, the periodic table had been restructured by Seaborg to its present layout, containing the actinide row below the lanthanide one. This led to americium being located right below its twin lanthanide element europium; it was thus by analogy named after the Americas: "The name americium (after the Americas) and the symbol Am are suggested for the element on the basis of its position as the sixth member of the actinide rare-earth series, analogous to europium, Eu, of the lanthanide series."
The new element was isolated from its oxides in a complex, multi-step process. First plutonium-239 nitrate (239PuNO3) solution was coated on a platinum foil of about 0.5 cm2 area, the solution was evaporated and the residue was converted into plutonium dioxide (PuO2) by calcining. After cyclotron irradiation, the coating was dissolved with nitric acid, and then precipitated as the hydroxide using concentrated aqueous ammonia solution. The residue was dissolved in perchloric acid. Further separation was carried out by ion exchange, yielding a certain isotope of curium. The separation of curium and americium was so painstaking that those elements were initially called by the Berkeley group as pandemonium (from Greek for all demons or hell) and delirium (from Latin for madness).
Initial experiments yielded four americium isotopes: 241Am, 242Am, 239Am and 238Am. Americium-241 was directly obtained from plutonium upon absorption of two neutrons. It decays by emission of a α-particle to 237Np; the half-life of this decay was first determined as years but then corrected to 432.2 years. | Americium | Wikipedia | 508 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
The times are half-lives
The second isotope 242Am was produced upon neutron bombardment of the already-created 241Am. Upon rapid β-decay, 242Am converts into the isotope of curium 242Cm (which had been discovered previously). The half-life of this decay was initially determined at 17 hours, which was close to the presently accepted value of 16.02 h.
The discovery of americium and curium in 1944 was closely related to the Manhattan Project; the results were confidential and declassified only in 1945. Seaborg leaked the synthesis of the elements 95 and 96 on the U.S. radio show for children Quiz Kids five days before the official presentation at an American Chemical Society meeting on 11 November 1945, when one of the listeners asked whether any new transuranium element besides plutonium and neptunium had been discovered during the war. After the discovery of americium isotopes 241Am and 242Am, their production and compounds were patented listing only Seaborg as the inventor. The initial americium samples weighed a few micrograms; they were barely visible and were identified by their radioactivity. The first substantial amounts of metallic americium weighing 40–200 micrograms were not prepared until 1951 by reduction of americium(III) fluoride with barium metal in high vacuum at 1100 °C.
Occurrence
The longest-lived and most common isotopes of americium, 241Am and 243Am, have half-lives of 432.2 and 7,370 years, respectively. Therefore, any primordial americium (americium that was present on Earth during its formation) should have decayed by now. Trace amounts of americium probably occur naturally in uranium minerals as a result of neutron capture and beta decay (238U → 239Pu → 240Pu → 241Am), though the quantities would be tiny and this has not been confirmed. Extraterrestrial long-lived 247Cm is probably also deposited on Earth and has 243Am as one of its intermediate decay products, but again this has not been confirmed. | Americium | Wikipedia | 424 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
Existing americium is concentrated in the areas used for the atmospheric nuclear weapons tests conducted between 1945 and 1980, as well as at the sites of nuclear incidents, such as the Chernobyl disaster. For example, the analysis of the debris at the testing site of the first U.S. hydrogen bomb, Ivy Mike, (1 November 1952, Enewetak Atoll), revealed high concentrations of various actinides including americium; but due to military secrecy, this result was not published until later, in 1956. Trinitite, the glassy residue left on the desert floor near Alamogordo, New Mexico, after the plutonium-based Trinity nuclear bomb test on 16 July 1945, contains traces of americium-241. Elevated levels of americium were also detected at the crash site of a US Boeing B-52 bomber aircraft, which carried four hydrogen bombs, in 1968 in Greenland.
In other regions, the average radioactivity of surface soil due to residual americium is only about 0.01 picocuries per gram (0.37 mBq/g). Atmospheric americium compounds are poorly soluble in common solvents and mostly adhere to soil particles. Soil analysis revealed about 1,900 times higher concentration of americium inside sandy soil particles than in the water present in the soil pores; an even higher ratio was measured in loam soils.
Americium is produced mostly artificially in small quantities, for research purposes. A tonne of spent nuclear fuel contains about 100 grams of various americium isotopes, mostly 241Am and 243Am. Their prolonged radioactivity is undesirable for the disposal, and therefore americium, together with other long-lived actinides, must be neutralized. The associated procedure may involve several steps, where americium is first separated and then converted by neutron bombardment in special reactors to short-lived nuclides. This procedure is well known as nuclear transmutation, but it is still being developed for americium. The transuranic elements from americium to fermium occurred naturally in the natural nuclear fission reactor at Oklo, but no longer do so.
Americium is also one of the elements that have theoretically been detected in Przybylski's Star.
Synthesis and extraction
Isotope nucleosynthesis | Americium | Wikipedia | 480 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
Americium has been produced in small quantities in nuclear reactors for decades, and kilograms of its 241Am and 243Am isotopes have been accumulated by now. Nevertheless, since it was first offered for sale in 1962, its price, about of 241Am, remains almost unchanged owing to the very complex separation procedure. The heavier isotope 243Am is produced in much smaller amounts; it is thus more difficult to separate, resulting in a higher cost of the order .
Americium is not synthesized directly from uranium – the most common reactor material – but from the plutonium isotope 239Pu. The latter needs to be produced first, according to the following nuclear process:
^{238}_{92}U ->[\ce{(n,\gamma)}] ^{239}_{92}U ->[\beta^-][23.5 \ \ce{min}] ^{239}_{93}Np ->[\beta^-][2.3565 \ \ce{d}] ^{239}_{94}Pu
The capture of two neutrons by 239Pu (a so-called (n,γ) reaction), followed by a β-decay, results in 241Am:
^{239}_{94}Pu ->[\ce{2(n,\gamma)}] ^{241}_{94}Pu ->[\beta^-][14.35 \ \ce{yr}] ^{241}_{95}Am
The plutonium present in spent nuclear fuel contains about 12% of 241Pu. Because it beta-decays to 241Am, 241Pu can be extracted and may be used to generate further 241Am. However, this process is rather slow: half of the original amount of 241Pu decays to 241Am after about 15 years, and the 241Am amount reaches a maximum after 70 years.
The obtained 241Am can be used for generating heavier americium isotopes by further neutron capture inside a nuclear reactor. In a light water reactor (LWR), 79% of 241Am converts to 242Am and 10% to its nuclear isomer 242mAm:
Americium-242 has a half-life of only 16 hours, which makes its further conversion to 243Am extremely inefficient. The latter isotope is produced instead in a process where 239Pu captures four neutrons under high neutron flux: | Americium | Wikipedia | 502 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
^{239}_{94}Pu ->[\ce{4(n,\gamma)}] \ ^{243}_{94}Pu ->[\beta^-][4.956 \ \ce{h}] ^{243}_{95}Am
Metal generation
Most synthesis routines yield a mixture of different actinide isotopes in oxide forms, from which isotopes of americium can be separated. In a typical procedure, the spent reactor fuel (e.g. MOX fuel) is dissolved in nitric acid, and the bulk of uranium and plutonium is removed using a PUREX-type extraction (Plutonium–URanium EXtraction) with tributyl phosphate in a hydrocarbon. The lanthanides and remaining actinides are then separated from the aqueous residue (raffinate) by a diamide-based extraction, to give, after stripping, a mixture of trivalent actinides and lanthanides. Americium compounds are then selectively extracted using multi-step chromatographic and centrifugation techniques with an appropriate reagent. A large amount of work has been done on the solvent extraction of americium. For example, a 2003 EU-funded project codenamed "EUROPART" studied triazines and other compounds as potential extraction agents. A bis-triazinyl bipyridine complex was proposed in 2009 as such a reagent is highly selective to americium (and curium). Separation of americium from the highly similar curium can be achieved by treating a slurry of their hydroxides in aqueous sodium bicarbonate with ozone, at elevated temperatures. Both Am and Cm are mostly present in solutions in the +3 valence state; whereas curium remains unchanged, americium oxidizes to soluble Am(IV) complexes which can be washed away.
Metallic americium is obtained by reduction from its compounds. Americium(III) fluoride was first used for this purpose. The reaction was conducted using elemental barium as reducing agent in a water- and oxygen-free environment inside an apparatus made of tantalum and tungsten.
An alternative is the reduction of americium dioxide by metallic lanthanum or thorium:
Physical properties | Americium | Wikipedia | 477 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
In the periodic table, americium is located to the right of plutonium, to the left of curium, and below the lanthanide europium, with which it shares many physical and chemical properties. Americium is a highly radioactive element. When freshly prepared, it has a silvery-white metallic lustre, but then slowly tarnishes in air. With a density of 12 g/cm3, americium is less dense than both curium (13.52 g/cm3) and plutonium (19.8 g/cm3); but has a higher density than europium (5.264 g/cm3)—mostly because of its higher atomic mass. Americium is relatively soft and easily deformable and has a significantly lower bulk modulus than the actinides before it: Th, Pa, U, Np and Pu. Its melting point of 1173 °C is significantly higher than that of plutonium (639 °C) and europium (826 °C), but lower than for curium (1340 °C). | Americium | Wikipedia | 224 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
At ambient conditions, americium is present in its most stable α form which has a hexagonal crystal symmetry, and a space group P63/mmc with cell parameters a = 346.8 pm and c = 1124 pm, and four atoms per unit cell. The crystal consists of a double-hexagonal close packing with the layer sequence ABAC and so is isotypic with α-lanthanum and several actinides such as α-curium. The crystal structure of americium changes with pressure and temperature. When compressed at room temperature to 5 GPa, α-Am transforms to the β modification, which has a face-centered cubic (fcc) symmetry, space group Fmm and lattice constant a = 489 pm. This fcc structure is equivalent to the closest packing with the sequence ABC. Upon further compression to 23 GPa, americium transforms to an orthorhombic γ-Am structure similar to that of α-uranium. There are no further transitions observed up to 52 GPa, except for an appearance of a monoclinic phase at pressures between 10 and 15 GPa. There is no consistency on the status of this phase in the literature, which also sometimes lists the α, β and γ phases as I, II and III. The β-γ transition is accompanied by a 6% decrease in the crystal volume; although theory also predicts a significant volume change for the α-β transition, it is not observed experimentally. The pressure of the α-β transition decreases with increasing temperature, and when α-americium is heated at ambient pressure, at 770 °C it changes into an fcc phase which is different from β-Am, and at 1075 °C it converts to a body-centered cubic structure. The pressure-temperature phase diagram of americium is thus rather similar to those of lanthanum, praseodymium and neodymium. | Americium | Wikipedia | 393 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
As with many other actinides, self-damage of the crystal structure due to alpha-particle irradiation is intrinsic to americium. It is especially noticeable at low temperatures, where the mobility of the produced structure defects is relatively low, by broadening of X-ray diffraction peaks. This effect makes somewhat uncertain the temperature of americium and some of its properties, such as electrical resistivity. So for americium-241, the resistivity at 4.2 K increases with time from about 2 μOhm·cm to 10 μOhm·cm after 40 hours, and saturates at about 16 μOhm·cm after 140 hours. This effect is less pronounced at room temperature, due to annihilation of radiation defects; also heating to room temperature the sample which was kept for hours at low temperatures restores its resistivity. In fresh samples, the resistivity gradually increases with temperature from about 2 μOhm·cm at liquid helium to 69 μOhm·cm at room temperature; this behavior is similar to that of neptunium, uranium, thorium and protactinium, but is different from plutonium and curium which show a rapid rise up to 60 K followed by saturation. The room temperature value for americium is lower than that of neptunium, plutonium and curium, but higher than for uranium, thorium and protactinium.
Americium is paramagnetic in a wide temperature range, from that of liquid helium, to room temperature and above. This behavior is markedly different from that of its neighbor curium which exhibits antiferromagnetic transition at 52 K. The thermal expansion coefficient of americium is slightly anisotropic and amounts to along the shorter a axis and for the longer c hexagonal axis. The enthalpy of dissolution of americium metal in hydrochloric acid at standard conditions is , from which the standard enthalpy change of formation (ΔfH°) of aqueous Am3+ ion is . The standard potential Am3+/Am0 is . | Americium | Wikipedia | 435 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
Chemical properties
Americium metal readily reacts with oxygen and dissolves in aqueous acids. The most stable oxidation state for americium is +3. The chemistry of americium(III) has many similarities to the chemistry of lanthanide(III) compounds. For example, trivalent americium forms insoluble fluoride, oxalate, iodate, hydroxide, phosphate and other salts. Compounds of americium in oxidation states +2, +4, +5, +6 and +7 have also been studied. This is the widest range that has been observed with actinide elements. The color of americium compounds in aqueous solution is as follows: Am3+ (yellow-reddish), Am4+ (yellow-reddish), ; (yellow), (brown) and (dark green). The absorption spectra have sharp peaks, due to f-f transitions' in the visible and near-infrared regions. Typically, Am(III) has absorption maxima at ca. 504 and 811 nm, Am(V) at ca. 514 and 715 nm, and Am(VI) at ca. 666 and 992 nm.
Americium compounds with oxidation state +4 and higher are strong oxidizing agents, comparable in strength to the permanganate ion () in acidic solutions. Whereas the Am4+ ions are unstable in solutions and readily convert to Am3+, compounds such as americium dioxide (AmO2) and americium(IV) fluoride (AmF4) are stable in the solid state.
The pentavalent oxidation state of americium was first observed in 1951. In acidic aqueous solution the ion is unstable with respect to disproportionation. The reaction
is typical. The chemistry of Am(V) and Am(VI) is comparable to the chemistry of uranium in those oxidation states. In particular, compounds like and are comparable to uranates and the ion is comparable to the uranyl ion, . Such compounds can be prepared by oxidation of Am(III) in dilute nitric acid with ammonium persulfate. Other oxidising agents that have been used include silver(I) oxide, ozone and sodium persulfate.
Chemical compounds | Americium | Wikipedia | 479 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
Oxygen compounds
Three americium oxides are known, with the oxidation states +2 (AmO), +3 (Am2O3) and +4 (AmO2). Americium(II) oxide was prepared in minute amounts and has not been characterized in detail. Americium(III) oxide is a red-brown solid with a melting point of 2205 °C. Americium(IV) oxide is the main form of solid americium which is used in nearly all its applications. As most other actinide dioxides, it is a black solid with a cubic (fluorite) crystal structure.
The oxalate of americium(III), vacuum dried at room temperature, has the chemical formula Am2(C2O4)3·7H2O. Upon heating in vacuum, it loses water at 240 °C and starts decomposing into AmO2 at 300 °C, the decomposition completes at about 470 °C. The initial oxalate dissolves in nitric acid with the maximum solubility of 0.25 g/L.
Halides
Halides of americium are known for the oxidation states +2, +3 and +4, where the +3 is most stable, especially in solutions.
Reduction of Am(III) compounds with sodium amalgam yields Am(II) salts – the black halides AmCl2, AmBr2 and AmI2. They are very sensitive to oxygen and oxidize in water, releasing hydrogen and converting back to the Am(III) state. Specific lattice constants are:
Orthorhombic AmCl2: a = , b = and c =
Tetragonal AmBr2: a = and c = . They can also be prepared by reacting metallic americium with an appropriate mercury halide HgX2, where X = Cl, Br or I:
{Am} + \underset{mercury\ halide}{HgX2} ->[{} \atop 400 - 500 ^\circ \ce C] {AmX2} + {Hg}
Americium(III) fluoride (AmF3) is poorly soluble and precipitates upon reaction of Am3+ and fluoride ions in weak acidic solutions:
Am^3+ + 3F^- -> AmF3(v) | Americium | Wikipedia | 487 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
The tetravalent americium(IV) fluoride (AmF4) is obtained by reacting solid americium(III) fluoride with molecular fluorine:
2AmF3 + F2 -> 2AmF4
Another known form of solid tetravalent americium fluoride is KAmF5. Tetravalent americium has also been observed in the aqueous phase. For this purpose, black Am(OH)4 was dissolved in 15-M NH4F with the americium concentration of 0.01 M. The resulting reddish solution had a characteristic optical absorption spectrum which is similar to that of AmF4 but differed from other oxidation states of americium. Heating the Am(IV) solution to 90 °C did not result in its disproportionation or reduction, however a slow reduction was observed to Am(III) and assigned to self-irradiation of americium by alpha particles.
Most americium(III) halides form hexagonal crystals with slight variation of the color and exact structure between the halogens. So, chloride (AmCl3) is reddish and has a structure isotypic to uranium(III) chloride (space group P63/m) and the melting point of 715 °C. The fluoride is isotypic to LaF3 (space group P63/mmc) and the iodide to BiI3 (space group R). The bromide is an exception with the orthorhombic PuBr3-type structure and space group Cmcm. Crystals of americium(III) chloride hexahydrate (AmCl3·6H2O) can be prepared by dissolving americium dioxide in hydrochloric acid and evaporating the liquid. Those crystals are hygroscopic and have yellow-reddish color and a monoclinic crystal structure.
Oxyhalides of americium in the form AmVIO2X2, AmVO2X, AmIVOX2 and AmIIIOX can be obtained by reacting the corresponding americium halide with oxygen or Sb2O3, and AmOCl can also be produced by vapor phase hydrolysis:
AmCl3 + H2O -> AmOCl + 2HCl | Americium | Wikipedia | 478 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
Chalcogenides and pnictides
The known chalcogenides of americium include the sulfide AmS2, selenides AmSe2 and Am3Se4, and tellurides Am2Te3 and AmTe2. The pnictides of americium (243Am) of the AmX type are known for the elements phosphorus, arsenic, antimony and bismuth. They crystallize in the rock-salt lattice.
Silicides and borides
Americium monosilicide (AmSi) and "disilicide" (nominally AmSix with: 1.87 < x < 2.0) were obtained by reduction of americium(III) fluoride with elementary silicon in vacuum at 1050 °C (AmSi) and 1150−1200 °C (AmSix). AmSi is a black solid isomorphic with LaSi, it has an orthorhombic crystal symmetry. AmSix has a bright silvery lustre and a tetragonal crystal lattice (space group I41/amd), it is isomorphic with PuSi2 and ThSi2. Borides of americium include AmB4 and AmB6. The tetraboride can be obtained by heating an oxide or halide of americium with magnesium diboride in vacuum or inert atmosphere.
Organoamericium compounds
Analogous to uranocene, americium is predicted to form the organometallic compound amerocene with two cyclooctatetraene ligands, with the chemical formula (η8-C8H8)2Am. A cyclopentadienyl complex is also known that is likely to be stoichiometrically AmCp3.
Formation of the complexes of the type Am(n-C3H7-BTP)3, where BTP stands for 2,6-di(1,2,4-triazin-3-yl)pyridine, in solutions containing n-C3H7-BTP and Am3+ ions has been confirmed by EXAFS. Some of these BTP-type complexes selectively interact with americium and therefore are useful in its selective separation from lanthanides and another actinides. | Americium | Wikipedia | 474 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
Biological aspects
Americium is an artificial element of recent origin, and thus does not have a biological requirement. It is harmful to life. It has been proposed to use bacteria for removal of americium and other heavy metals from rivers and streams. Thus, Enterobacteriaceae of the genus Citrobacter precipitate americium ions from aqueous solutions, binding them into a metal-phosphate complex at their cell walls. Several studies have been reported on the biosorption and bioaccumulation of americium by bacteria and fungi. In the laboratory, both americium and curium were found to support the growth of methylotrophs.
Fission
The isotope 242mAm (half-life 141 years) has the largest cross sections for absorption of thermal neutrons (5,700 barns), that results in a small critical mass for a sustained nuclear chain reaction. The critical mass for a bare 242mAm sphere is about 9–14 kg (the uncertainty results from insufficient knowledge of its material properties). It can be lowered to 3–5 kg with a metal reflector and should become even smaller with a water reflector. Such small critical mass is favorable for portable nuclear weapons, but those based on 242mAm are not known yet, probably because of its scarcity and high price. The critical masses of the two readily available isotopes, 241Am and 243Am, are relatively high – 57.6 to 75.6 kg for 241Am and 209 kg for 243Am. Scarcity and high price yet hinder application of americium as a nuclear fuel in nuclear reactors.
There are proposals of very compact 10-kW high-flux reactors using as little as 20 grams of 242mAm. Such low-power reactors would be relatively safe to use as neutron sources for radiation therapy in hospitals.
Isotopes | Americium | Wikipedia | 376 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
About 18 isotopes and 11 nuclear isomers are known for americium, having mass numbers 229, 230, and 232 through 247. There are two long-lived alpha-emitters; 243Am has a half-life of 7,370 years and is the most stable isotope, and 241Am has a half-life of 432.2 years. The most stable nuclear isomer is 242m1Am; it has a long half-life of 141 years. The half-lives of other isotopes and isomers range from 0.64 microseconds for 245m1Am to 50.8 hours for 240Am. As with most other actinides, the isotopes of americium with odd number of neutrons have relatively high rate of nuclear fission and low critical mass.
Americium-241 decays to 237Np emitting alpha particles of 5 different energies, mostly at 5.486 MeV (85.2%) and 5.443 MeV (12.8%). Because many of the resulting states are metastable, they also emit gamma rays with the discrete energies between 26.3 and 158.5 keV.
Americium-242 is a short-lived isotope with a half-life of 16.02 h. It mostly (82.7%) converts by β-decay to 242Cm, but also by electron capture to 242Pu (17.3%). Both 242Cm and 242Pu transform via nearly the same decay chain through 238Pu down to 234U.
Nearly all (99.541%) of 242m1Am decays by internal conversion to 242Am and the remaining 0.459% by α-decay to 238Np. The latter subsequently decays to 238Pu and then to 234U.
Americium-243 transforms by α-emission into 239Np, which converts by β-decay to 239Pu, and the 239Pu changes into 235U by emitting an α-particle.
Applications
Ionization-type smoke detector
Americium is used in the most common type of household smoke detector, which uses 241Am in the form of americium dioxide as its source of ionizing radiation. This isotope is preferred over 226Ra because it emits 5 times more alpha particles and relatively little harmful gamma radiation. | Americium | Wikipedia | 474 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
The amount of americium in a typical new smoke detector is 1 microcurie (37 kBq) or 0.29 microgram. This amount declines slowly as the americium decays into neptunium-237, a different transuranic element with a much longer half-life (about 2.14 million years). With its half-life of 432.2 years, the americium in a smoke detector includes about 3% neptunium after 19 years, and about 5% after 32 years. The radiation passes through an ionization chamber, an air-filled space between two electrodes, and permits a small, constant current between the electrodes. Any smoke that enters the chamber absorbs the alpha particles, which reduces the ionization and affects this current, triggering the alarm. Compared to the alternative optical smoke detector, the ionization smoke detector is cheaper and can detect particles which are too small to produce significant light scattering; however, it is more prone to false alarms.
Radionuclide
As 241Am has a roughly similar half-life to 238Pu (432.2 years vs. 87 years), it has been proposed as an active element of radioisotope thermoelectric generators, for example in spacecraft. Although americium produces less heat and electricity – the power yield is 114.7 mW/g for 241Am and 6.31 mW/g for 243Am (cf. 390 mW/g for 238Pu) – and its radiation poses more threat to humans owing to neutron emission, the European Space Agency is considering using americium for its space probes.
Another proposed space-related application of americium is a fuel for space ships with nuclear propulsion. It relies on the very high rate of nuclear fission of 242mAm, which can be maintained even in a micrometer-thick foil. Small thickness avoids the problem of self-absorption of emitted radiation. This problem is pertinent to uranium or plutonium rods, in which only surface layers provide alpha-particles. The fission products of 242mAm can either directly propel the spaceship or they can heat a thrusting gas. They can also transfer their energy to a fluid and generate electricity through a magnetohydrodynamic generator. | Americium | Wikipedia | 462 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
One more proposal which utilizes the high nuclear fission rate of 242mAm is a nuclear battery. Its design relies not on the energy of the emitted by americium alpha particles, but on their charge, that is the americium acts as the self-sustaining "cathode". A single 3.2 kg 242mAm charge of such battery could provide about 140 kW of power over a period of 80 days. Even with all the potential benefits, the current applications of 242mAm are as yet hindered by the scarcity and high price of this particular nuclear isomer.
In 2019, researchers at the UK National Nuclear Laboratory and the University of Leicester demonstrated the use of heat generated by americium to illuminate a small light bulb. This technology could lead to systems to power missions with durations up to 400 years into interstellar space, where solar panels do not function.
Neutron source
The oxide of 241Am pressed with beryllium is an efficient neutron source. Here americium acts as the alpha source, and beryllium produces neutrons owing to its large cross-section for the (α,n) nuclear reaction:
^{241}_{95}Am -> ^{237}_{93}Np + ^{4}_{2}He + \gamma
^{9}_{4}Be + ^{4}_{2}He -> ^{12}_{6}C + ^{1}_{0}n + \gamma
The most widespread use of 241AmBe neutron sources is a neutron probe – a device used to measure the quantity of water present in soil, as well as moisture/density for quality control in highway construction. 241Am neutron sources are also used in well logging applications, as well as in neutron radiography, tomography and other radiochemical investigations.
Production of other elements
Americium is a starting material for the production of other transuranic elements and transactinides – for example, 82.7% of 242Am decays to 242Cm and 17.3% to 242Pu. In the nuclear reactor, 242Am is also up-converted by neutron capture to 243Am and 244Am, which transforms by β-decay to 244Cm: | Americium | Wikipedia | 460 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
^{243}_{95}Am ->[\ce{(n,\gamma)}] ^{244}_{95}Am ->[\beta^-][10.1 \ \ce{h}] ^{244}_{96}Cm
Irradiation of 241Am by 12C or 22Ne ions yields the isotopes 247Es (einsteinium) or 260Db (dubnium), respectively. Furthermore, the element berkelium (243Bk isotope) had been first intentionally produced and identified by bombarding 241Am with alpha particles, in 1949, by the same Berkeley group, using the same 60-inch cyclotron. Similarly, nobelium was produced at the Joint Institute for Nuclear Research, Dubna, Russia, in 1965 in several reactions, one of which included irradiation of 243Am with 15N ions. Besides, one of the synthesis reactions for lawrencium, discovered by scientists at Berkeley and Dubna, included bombardment of 243Am with 18O.
Spectrometer
Americium-241 has been used as a portable source of both gamma rays and alpha particles for a number of medical and industrial uses. The 59.5409 keV gamma ray emissions from 241Am in such sources can be used for indirect analysis of materials in radiography and X-ray fluorescence spectroscopy, as well as for quality control in fixed nuclear density gauges and nuclear densometers. For example, the element has been employed to gauge glass thickness to help create flat glass. Americium-241 is also suitable for calibration of gamma-ray spectrometers in the low-energy range, since its spectrum consists of nearly a single peak and negligible Compton continuum (at least three orders of magnitude lower intensity). Americium-241 gamma rays were also used to provide passive diagnosis of thyroid function. This medical application is however obsolete.
Health concerns
As a highly radioactive element, americium and its compounds must be handled only in an appropriate laboratory under special arrangements. Although most americium isotopes predominantly emit alpha particles which can be blocked by thin layers of common materials, many of the daughter products emit gamma-rays and neutrons which have a long penetration depth. | Americium | Wikipedia | 462 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
If consumed, most of the americium is excreted within a few days, with only 0.05% absorbed in the blood, of which roughly 45% goes to the liver and 45% to the bones, and the remaining 10% is excreted. The uptake to the liver depends on the individual and increases with age. In the bones, americium is first deposited over cortical and trabecular surfaces and slowly redistributes over the bone with time. The biological half-life of 241Am is 50 years in the bones and 20 years in the liver, whereas in the gonads (testicles and ovaries) it remains permanently; in all these organs, americium promotes formation of cancer cells as a result of its radioactivity.
Americium often enters landfills from discarded smoke detectors. The rules associated with the disposal of smoke detectors are relaxed in most jurisdictions. In 1994, 17-year-old David Hahn extracted the americium from about 100 smoke detectors in an attempt to build a breeder nuclear reactor. There have been a few cases of exposure to americium, the worst case being that of chemical operations technician Harold McCluskey, who at the age of 64 was exposed to 500 times the occupational standard for americium-241 as a result of an explosion in his lab. McCluskey died at the age of 75 of unrelated pre-existing disease. | Americium | Wikipedia | 294 | 900 | https://en.wikipedia.org/wiki/Americium | Physical sciences | Chemical elements_2 | null |
Astatine is a chemical element; it has symbol At and atomic number 85. It is the rarest naturally occurring element in the Earth's crust, occurring only as the decay product of various heavier elements. All of astatine's isotopes are short-lived; the most stable is astatine-210, with a half-life of 8.1 hours. Consequently, a solid sample of the element has never been seen, because any macroscopic specimen would be immediately vaporized by the heat of its radioactivity.
The bulk properties of astatine are not known with certainty. Many of them have been estimated from its position on the periodic table as a heavier analog of fluorine, chlorine, bromine, and iodine, the four stable halogens. However, astatine also falls roughly along the dividing line between metals and nonmetals, and some metallic behavior has also been observed and predicted for it. Astatine is likely to have a dark or lustrous appearance and may be a semiconductor or possibly a metal. Chemically, several anionic species of astatine are known and most of its compounds resemble those of iodine, but it also sometimes displays metallic characteristics and shows some similarities to silver.
The first synthesis of astatine was in 1940 by Dale R. Corson, Kenneth Ross MacKenzie, and Emilio G. Segrè at the University of California, Berkeley. They named it from the Ancient Greek () 'unstable'. Four isotopes of astatine were subsequently found to be naturally occurring, although much less than one gram is present at any given time in the Earth's crust. Neither the most stable isotope, astatine-210, nor the medically useful astatine-211 occur naturally; they are usually produced by bombarding bismuth-209 with alpha particles.
Characteristics
Astatine is an extremely radioactive element; all its isotopes have half-lives of 8.1 hours or less, decaying into other astatine isotopes, bismuth, polonium, or radon. Most of its isotopes are very unstable, with half-lives of seconds or less. Of the first 101 elements in the periodic table, only francium is less stable, and all the astatine isotopes more stable than the longest-lived francium isotopes (205–211At) are in any case synthetic and do not occur in nature. | Astatine | Wikipedia | 500 | 901 | https://en.wikipedia.org/wiki/Astatine | Physical sciences | Chemical elements_2 | null |
The bulk properties of astatine are not known with any certainty. Research is limited by its short half-life, which prevents the creation of weighable quantities. A visible piece of astatine would immediately vaporize itself because of the heat generated by its intense radioactivity. It remains to be seen if, with sufficient cooling, a macroscopic quantity of astatine could be deposited as a thin film. Astatine is usually classified as either a nonmetal or a metalloid; metal formation has also been predicted.
Physical
Most of the physical properties of astatine have been estimated (by interpolation or extrapolation), using theoretically or empirically derived methods. For example, halogens get darker with increasing atomic weight – fluorine is nearly colorless, chlorine is yellow-green, bromine is red-brown, and iodine is dark gray/violet. Astatine is sometimes described as probably being a black solid (assuming it follows this trend), or as having a metallic appearance (if it is a metalloid or a metal).
Astatine sublimes less readily than iodine, having a lower vapor pressure. Even so, half of a given quantity of astatine will vaporize in approximately an hour if put on a clean glass surface at room temperature. The absorption spectrum of astatine in the middle ultraviolet region has lines at 224.401 and 216.225 nm, suggestive of 6p to 7s transitions.
The structure of solid astatine is unknown. As an analog of iodine it may have an orthorhombic crystalline structure composed of diatomic astatine molecules, and be a semiconductor (with a band gap of 0.7 eV). Alternatively, if condensed astatine forms a metallic phase, as has been predicted, it may have a monatomic face-centered cubic structure; in this structure, it may well be a superconductor, like the similar high-pressure phase of iodine. Metallic astatine is expected to have a density of 8.91–8.95 g/cm3. | Astatine | Wikipedia | 433 | 901 | https://en.wikipedia.org/wiki/Astatine | Physical sciences | Chemical elements_2 | null |
Evidence for (or against) the existence of diatomic astatine (At2) is sparse and inconclusive. Some sources state that it does not exist, or at least has never been observed, while other sources assert or imply its existence. Despite this controversy, many properties of diatomic astatine have been predicted; for example, its bond length would be , dissociation energy <, and heat of vaporization (∆Hvap) 54.39 kJ/mol. Many values have been predicted for the melting and boiling points of astatine, but only for At2.
Chemical
The chemistry of astatine is "clouded by the extremely low concentrations at which astatine experiments have been conducted, and the possibility of reactions with impurities, walls and filters, or radioactivity by-products, and other unwanted nano-scale interactions". Many of its apparent chemical properties have been observed using tracer studies on extremely dilute astatine solutions, typically less than 10−10 mol·L−1. Some properties, such as anion formation, align with other halogens. Astatine has some metallic characteristics as well, such as plating onto a cathode, and coprecipitating with metal sulfides in hydrochloric acid. It forms complexes with EDTA, a metal chelating agent, and is capable of acting as a metal in antibody radiolabeling; in some respects, astatine in the +1 state is akin to silver in the same state. Most of the organic chemistry of astatine is, however, analogous to that of iodine. It has been suggested that astatine can form a stable monatomic cation in aqueous solution. | Astatine | Wikipedia | 360 | 901 | https://en.wikipedia.org/wiki/Astatine | Physical sciences | Chemical elements_2 | null |
Astatine has an electronegativity of 2.2 on the revised Pauling scale – lower than that of iodine (2.66) and the same as hydrogen. In hydrogen astatide (HAt), the negative charge is predicted to be on the hydrogen atom, implying that this compound could be referred to as astatine hydride according to certain nomenclatures. That would be consistent with the electronegativity of astatine on the Allred–Rochow scale (1.9) being less than that of hydrogen (2.2). However, official IUPAC stoichiometric nomenclature is based on an idealized convention of determining the relative electronegativities of the elements by the mere virtue of their position within the periodic table. According to this convention, astatine is handled as though it is more electronegative than hydrogen, irrespective of its true electronegativity. The electron affinity of astatine, at 233 kJ mol−1, is 21% less than that of iodine. In comparison, the value of Cl (349) is 6.4% higher than F (328); Br (325) is 6.9% less than Cl; and I (295) is 9.2% less than Br. The marked reduction for At was predicted as being due to spin–orbit interactions. The first ionization energy of astatine is about 899 kJ mol−1, which continues the trend of decreasing first ionization energies down the halogen group (fluorine, 1681; chlorine, 1251; bromine, 1140; iodine, 1008).
Compounds
Less reactive than iodine, astatine is the least reactive of the halogens; the chemical properties of tennessine, the next-heavier group 17 element, have not yet been investigated, however. Astatine compounds have been synthesized in nano-scale amounts and studied as intensively as possible before their radioactive disintegration. The reactions involved have been typically tested with dilute solutions of astatine mixed with larger amounts of iodine. Acting as a carrier, the iodine ensures there is sufficient material for laboratory techniques (such as filtration and precipitation) to work. Like iodine, astatine has been shown to adopt odd-numbered oxidation states ranging from −1 to +7. | Astatine | Wikipedia | 492 | 901 | https://en.wikipedia.org/wiki/Astatine | Physical sciences | Chemical elements_2 | null |
Only a few compounds with metals have been reported, in the form of astatides of sodium, palladium, silver, thallium, and lead. Some characteristic properties of silver and sodium astatide, and the other hypothetical alkali and alkaline earth astatides, have been estimated by extrapolation from other metal halides.
The formation of an astatine compound with hydrogen – usually referred to as hydrogen astatide – was noted by the pioneers of astatine chemistry. As mentioned, there are grounds for instead referring to this compound as astatine hydride. It is easily oxidized; acidification by dilute nitric acid gives the At0 or At+ forms, and the subsequent addition of silver(I) may only partially, at best, precipitate astatine as silver(I) astatide (AgAt). Iodine, in contrast, is not oxidized, and precipitates readily as silver(I) iodide.
Astatine is known to bind to boron, carbon, and nitrogen. Various boron cage compounds have been prepared with At–B bonds, these being more stable than At–C bonds. Astatine can replace a hydrogen atom in benzene to form astatobenzene C6H5At; this may be oxidized to C6H5AtCl2 by chlorine. By treating this compound with an alkaline solution of hypochlorite, C6H5AtO2 can be produced. The dipyridine-astatine(I) cation, [At(C5H5N)2]+, forms ionic compounds with perchlorate (a non-coordinating anion) and with nitrate, [At(C5H5N)2]NO3. This cation exists as a coordination complex in which two dative covalent bonds separately link the astatine(I) centre with each of the pyridine rings via their nitrogen atoms. | Astatine | Wikipedia | 421 | 901 | https://en.wikipedia.org/wiki/Astatine | Physical sciences | Chemical elements_2 | null |
With oxygen, there is evidence of the species AtO− and AtO+ in aqueous solution, formed by the reaction of astatine with an oxidant such as elemental bromine or (in the last case) by sodium persulfate in a solution of perchloric acid. The species previously thought to be has since been determined to be , a hydrolysis product of AtO+ (another such hydrolysis product being AtOOH). The well characterized anion can be obtained by, for example, the oxidation of astatine with potassium hypochlorite in a solution of potassium hydroxide. Preparation of lanthanum triastatate La(AtO3)3, following the oxidation of astatine by a hot Na2S2O8 solution, has been reported. Further oxidation of , such as by xenon difluoride (in a hot alkaline solution) or periodate (in a neutral or alkaline solution), yields the perastatate ion ; this is only stable in neutral or alkaline solutions. Astatine is also thought to be capable of forming cations in salts with oxyanions such as iodate or dichromate; this is based on the observation that, in acidic solutions, monovalent or intermediate positive states of astatine coprecipitate with the insoluble salts of metal cations such as silver(I) iodate or thallium(I) dichromate.
Astatine may form bonds to the other chalcogens; these include S7At+ and with sulfur, a coordination selenourea compound with selenium, and an astatine–tellurium colloid with tellurium. | Astatine | Wikipedia | 363 | 901 | https://en.wikipedia.org/wiki/Astatine | Physical sciences | Chemical elements_2 | null |
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