Incandescent light bulb
An incandescent light bulb, incandescent lamp or incandescent light globe is an electric light with a wire filament heated to such a high temperature that it glows with visible light. The filament is protected from oxidation with a glass or fused quartz bulb, filled with inert gas or a vacuum. In a halogen lamp, filament evaporation is slowed by a chemical process that redeposits metal vapor onto the filament, thereby extending its life; the light bulb is supplied with electric current by feed-through terminals or wires embedded in the glass. Most bulbs are used in a socket which provides electrical connections. Incandescent bulbs are manufactured in a wide range of sizes, light output, voltage ratings, from 1.5 volts to about 300 volts. They require no external regulating equipment, have low manufacturing costs, work well on either alternating current or direct current; as a result, the incandescent bulb is used in household and commercial lighting, for portable lighting such as table lamps, car headlamps, flashlights, for decorative and advertising lighting.
Incandescent bulbs are much less efficient than other types of electric lighting. The remaining energy is converted into heat; the luminous efficacy of a typical incandescent bulb for 120 V operation is 16 lumens per watt, compared with 60 lm/W for a compact fluorescent bulb or 150 lm/W for some white LED lamps. Some applications of the incandescent bulb deliberately use the heat generated by the filament; such applications include incubators, brooding boxes for poultry, heat lights for reptile tanks, infrared heating for industrial heating and drying processes, lava lamps, the Easy-Bake Oven toy. Incandescent bulbs have short lifetimes compared with other types of lighting. Incandescent bulbs have been replaced in many applications by other types of electric light, such as fluorescent lamps, compact fluorescent lamps, cold cathode fluorescent lamps, high-intensity discharge lamps, light-emitting diode lamps; some jurisdictions, such as the European Union, China and United States, are in the process of phasing out the use of incandescent light bulbs while others, including Colombia, Cuba and Brazil, have prohibited them already.
In addressing the question of who invented the incandescent lamp, historians Robert Friedel and Paul Israel list 22 inventors of incandescent lamps prior to Joseph Swan and Thomas Edison. They conclude that Edison's version was able to outstrip the others because of a combination of three factors: an effective incandescent material, a higher vacuum than others were able to achieve and a high resistance that made power distribution from a centralized source economically viable. Historian Thomas Hughes has attributed Edison's success to his development of an entire, integrated system of electric lighting; the lamp was a small component in his system of electric lighting, no more critical to its effective functioning than the Edison Jumbo generator, the Edison main and feeder, the parallel-distribution system. Other inventors with generators and incandescent lamps, with comparable ingenuity and excellence, have long been forgotten because their creators did not preside over their introduction in a system of lighting.
In 1761 Ebenezer Kinnersley demonstrated heating a wire to incandescence. In 1802, Humphry Davy used what he described as "a battery of immense size", consisting of 2,000 cells housed in the basement of the Royal Institution of Great Britain, to create an incandescent light by passing the current through a thin strip of platinum, chosen because the metal had an high melting point, it was not bright enough nor did it last long enough to be practical, but it was the precedent behind the efforts of scores of experimenters over the next 75 years. Over the first three-quarters of the 19th century, many experimenters worked with various combinations of platinum or iridium wires, carbon rods, evacuated or semi-evacuated enclosures. Many of these devices were demonstrated and some were patented. In 1835, James Bowman Lindsay demonstrated a constant electric light at a public meeting in Dundee, Scotland, he stated that he could "read a book at a distance of one and a half feet". Lindsay, a lecturer at the Watt Institution in Dundee, Scotland, at the time, had developed a light, not combustible, created no smoke or smell and was less expensive to produce than Davy's platinum-dependent bulb.
However, having perfected the device to his own satisfaction, he turned to the problem of wireless telegraphy and did not develop the electric light any further. His claims are not well documented, although he is credited in Challoner et al. with being the inventor of the "Incandescent Light Bulb". In 1838, Belgian lithographer Marcellin Jobard invented an incandescent light bulb with a vacuum atmosphere using a carbon filament. In 1840, British scientist Warren de la Rue enclosed a coiled platinum filament in a vacuum tube and passed an electric current through it; the design was based on the concept that the high melting point of platinum would allow it to operate at high temperatures and that the evacuated chamber would contain fewer gas molecules to react with the platinum, improving its longevity. Although a workable design, the cost of the platinum made it impractical for commercial use. In 1841, Frederick de Moleyns of England was granted the first patent for an incandescent lamp, with a design using platinum wires contained within a vacuum
Density
The density, or more the volumetric mass density, of a substance is its mass per unit volume. The symbol most used for density is ρ, although the Latin letter D can be used. Mathematically, density is defined as mass divided by volume: ρ = m V where ρ is the density, m is the mass, V is the volume. In some cases, density is loosely defined as its weight per unit volume, although this is scientifically inaccurate – this quantity is more called specific weight. For a pure substance the density has the same numerical value as its mass concentration. Different materials have different densities, density may be relevant to buoyancy and packaging. Osmium and iridium are the densest known elements at standard conditions for temperature and pressure but certain chemical compounds may be denser. To simplify comparisons of density across different systems of units, it is sometimes replaced by the dimensionless quantity "relative density" or "specific gravity", i.e. the ratio of the density of the material to that of a standard material water.
Thus a relative density less than one means. The density of a material varies with pressure; this variation is small for solids and liquids but much greater for gases. Increasing the pressure on an object decreases the volume of the object and thus increases its density. Increasing the temperature of a substance decreases its density by increasing its volume. In most materials, heating the bottom of a fluid results in convection of the heat from the bottom to the top, due to the decrease in the density of the heated fluid; this causes it to rise relative to more dense unheated material. The reciprocal of the density of a substance is called its specific volume, a term sometimes used in thermodynamics. Density is an intensive property in that increasing the amount of a substance does not increase its density. In a well-known but apocryphal tale, Archimedes was given the task of determining whether King Hiero's goldsmith was embezzling gold during the manufacture of a golden wreath dedicated to the gods and replacing it with another, cheaper alloy.
Archimedes knew that the irregularly shaped wreath could be crushed into a cube whose volume could be calculated and compared with the mass. Baffled, Archimedes is said to have taken an immersion bath and observed from the rise of the water upon entering that he could calculate the volume of the gold wreath through the displacement of the water. Upon this discovery, he leapt from his bath and ran naked through the streets shouting, "Eureka! Eureka!". As a result, the term "eureka" entered common parlance and is used today to indicate a moment of enlightenment; the story first appeared in written form in Vitruvius' books of architecture, two centuries after it took place. Some scholars have doubted the accuracy of this tale, saying among other things that the method would have required precise measurements that would have been difficult to make at the time. From the equation for density, mass density has units of mass divided by volume; as there are many units of mass and volume covering many different magnitudes there are a large number of units for mass density in use.
The SI unit of kilogram per cubic metre and the cgs unit of gram per cubic centimetre are the most used units for density. One g/cm3 is equal to one thousand kg/m3. One cubic centimetre is equal to one millilitre. In industry, other larger or smaller units of mass and or volume are more practical and US customary units may be used. See below for a list of some of the most common units of density. A number of techniques as well as standards exist for the measurement of density of materials; such techniques include the use of a hydrometer, Hydrostatic balance, immersed body method, air comparison pycnometer, oscillating densitometer, as well as pour and tap. However, each individual method or technique measures different types of density, therefore it is necessary to have an understanding of the type of density being measured as well as the type of material in question; the density at all points of a homogeneous object equals its total mass divided by its total volume. The mass is measured with a scale or balance.
To determine the density of a liquid or a gas, a hydrometer, a dasymeter or a Coriolis flow meter may be used, respectively. Hydrostatic weighing uses the displacement of water due to a submerged object to determine the density of the object. If the body is not homogeneous its density varies between different regions of the object. In that case the density around any given location is determined by calculating the density of a small volume around that location. In the limit of an infinitesimal volume the density of an inhomogeneous object at a point becomes: ρ = d m / d V, where d V is an elementary volume at position r; the mass of the body t
Rayleigh scattering
Rayleigh scattering, named after the British physicist Lord Rayleigh, is the predominantly elastic scattering of light or other electromagnetic radiation by particles much smaller than the wavelength of the radiation. Rayleigh scattering is, hence, a parametric process; the particles may be individual molecules. It can occur when light travels through transparent solids and liquids, is most prominently seen in gases. Rayleigh scattering results from the electric polarizability of the particles; the oscillating electric field of a light wave acts on the charges within a particle, causing them to move at the same frequency. The particle therefore becomes a small radiating dipole; this radiation is an integral part of the photon and no excitation or deexcitation occurs. Rayleigh scattering of sunlight in Earth's atmosphere causes diffuse sky radiation, the reason for the blue color of the daytime and twilight sky, as well as the yellowish to reddish hue of the low Sun. For wave frequencies well below the resonance frequency of the scattering particle, the amount of scattering is inversely proportional to the fourth power of the wavelength.
Rayleigh scattering of molecular nitrogen and oxygen in the atmosphere includes elastic scattering as well as the inelastic contribution from rotational Raman scattering in air, since the changes in wavenumber of the scattered photon are smaller than 50 cm−1. This can lead to changes in the rotational state of the molecules. Furthermore, the inelastic contribution has the same wavelengths dependency as the elastic part. Scattering by particles similar to, or larger than, the wavelength of light is treated by the Mie theory, the discrete dipole approximation and other computational techniques. Rayleigh scattering applies to particles that are small with respect to wavelengths of light, that are optically "soft". On the other hand, anomalous diffraction theory applies to larger particles. In 1859, while attempting to determine whether any contaminants remained in the purified air he used for infrared experiments, John Tyndall discovered that bright light scattering off nanoscopic particulates was faintly blue-tinted.
He conjectured that a similar scattering of sunlight gave the sky its blue hue, but he could not explain the preference for blue light, nor could atmospheric dust explain the intensity of the sky's color. In 1871, Lord Rayleigh published two papers on the color and polarization of skylight to quantify Tyndall's effect in water droplets in terms of the tiny particulates' volumes and refractive indices. In 1881 with the benefit of James Clerk Maxwell's 1865 proof of the electromagnetic nature of light, he showed that his equations followed from electromagnetism. In 1899, he showed that they applied to individual molecules, with terms containing particulate volumes and refractive indices replaced with terms for molecular polarizability; the size of a scattering particle is parameterized by the ratio where r is its characteristic length and λ is the wavelength of the light. The amplitude of light scattered from within any transparent dielectric is proportional to the inverse square of its wavelength and to the volume of material, to the cube of its characteristic length.
The wavelength dependence is characteristic of dipole scattering and the volume dependence will apply to any scattering mechanism. Objects with x ≫ 1 act as geometric shapes. At the intermediate x ≃ 1 of Mie scattering, interference effects develop through phase variations over the object's surface. Rayleigh scattering applies to the case when the scattering particle is small and the whole surface re-radiates with the same phase; because the particles are randomly positioned, the scattered light arrives at a particular point with a random collection of phases. In detail, the intensity I of light scattered by any one of the small spheres of diameter d and refractive index n from a beam of unpolarized light of wavelength λ and intensity I0 is given by where R is the distance to the particle and θ is the scattering angle. Averaging this over all angles gives the Rayleigh scattering cross-sectionThe fraction of light scattered by a group of scattering particles is the number of particles per unit volume N times the cross-section.
For example, the major constituent of the atmosphere, has a Rayleigh cross section of 5.1×10−31 m2 at a wavelength of 532 nm. This means that at atmospheric pressure, where there are about 2×1025 molecules per cubic meter, about a fraction 10−5 of the light will be scattered for every meter of travel; the strong wavelength dependence of the scattering means that shorter wavelengths are scattered more than longer wavelengths. The expression above can be written in terms of individual molecules by expressing the dependence on refractive index in terms of the molecular polarizability α, proportional to the dipole moment induced by the electric field of the light. In this case, the Rayleigh scattering intensity for a single particle is given in CGS-units by When the dielectric constant ϵ of a certain region of volume V is different from the average dielectric constant of the medium ϵ ¯ {\displaystyle
Visible spectrum
The visible spectrum is the portion of the electromagnetic spectrum, visible to the human eye. Electromagnetic radiation in this range of wavelengths is called visible light or light. A typical human eye will respond to wavelengths from about 380 to 740 nanometers. In terms of frequency, this corresponds to a band in the vicinity of 430–770 THz; the spectrum does not contain all the colors. Unsaturated colors such as pink, or purple variations like magenta, for example, are absent because they can only be made from a mix of multiple wavelengths. Colors containing only one wavelength are called pure colors or spectral colors. Visible wavelengths pass unattenuated through the Earth's atmosphere via the "optical window" region of the electromagnetic spectrum. An example of this phenomenon is when clean air scatters blue light more than red light, so the midday sky appears blue; the optical window is referred to as the "visible window" because it overlaps the human visible response spectrum. The near infrared window lies just out of the human vision, as well as the medium wavelength infrared window, the long wavelength or far infrared window, although other animals may experience them.
In the 13th century, Roger Bacon theorized that rainbows were produced by a similar process to the passage of light through glass or crystal. In the 17th century, Isaac Newton discovered that prisms could disassemble and reassemble white light, described the phenomenon in his book Opticks, he was the first to use the word spectrum in this sense in print in 1671 in describing his experiments in optics. Newton observed that, when a narrow beam of sunlight strikes the face of a glass prism at an angle, some is reflected and some of the beam passes into and through the glass, emerging as different-colored bands. Newton hypothesized light to be made up of "corpuscles" of different colors, with the different colors of light moving at different speeds in transparent matter, red light moving more than violet in glass; the result is that red light is bent less than violet as it passes through the prism, creating a spectrum of colors. Newton divided the spectrum into six named colors: red, yellow, green and violet.
He added indigo as the seventh color since he believed that seven was a perfect number as derived from the ancient Greek sophists, of there being a connection between the colors, the musical notes, the known objects in the solar system, the days of the week. The human eye is insensitive to indigo's frequencies, some people who have otherwise-good vision cannot distinguish indigo from blue and violet. For this reason, some commentators, including Isaac Asimov, have suggested that indigo should not be regarded as a color in its own right but as a shade of blue or violet. Evidence indicates that what Newton meant by "indigo" and "blue" does not correspond to the modern meanings of those color words. Comparing Newton's observation of prismatic colors to a color image of the visible light spectrum shows that "indigo" corresponds to what is today called blue, whereas "blue" corresponds to cyan. In the 18th century, Johann Wolfgang von Goethe wrote about optical spectra in his Theory of Colours. Goethe used the word spectrum to designate a ghostly optical afterimage, as did Schopenhauer in On Vision and Colors.
Goethe argued. Where Newton narrowed the beam of light to isolate the phenomenon, Goethe observed that a wider aperture produces not a spectrum but rather reddish-yellow and blue-cyan edges with white between them; the spectrum appears only. In the early 19th century, the concept of the visible spectrum became more definite, as light outside the visible range was discovered and characterized by William Herschel and Johann Wilhelm Ritter, Thomas Young, Thomas Johann Seebeck, others. Young was the first to measure the wavelengths of different colors of light, in 1802; the connection between the visible spectrum and color vision was explored by Thomas Young and Hermann von Helmholtz in the early 19th century. Their theory of color vision proposed that the eye uses three distinct receptors to perceive color. Many species can see light within frequencies outside the human "visible spectrum". Bees and many other insects can detect ultraviolet light. Plant species that depend on insect pollination may owe reproductive success to their appearance in ultraviolet light rather than how colorful they appear to humans.
Birds, can see into the ultraviolet, some have sex-dependent markings on their plumage that are visible only in the ultraviolet range. Many animals that can see into the ultraviolet range cannot see red light or any other reddish wavelengths. Bees' visible spectrum ends at about 590 nm. Birds can see some red wavelengths; the popular belief that the common goldfish is the only animal that can see both infrared and ultraviolet light is incorrect, because goldfish cannot see infrared light. Dogs are thought to be color blind but they have been shown to be sensitive to colors, though not as many as humans; some snakes can "see" radiant heat at wavelengths between 5 and 30 μm to a degree of accuracy such that a blind rattlesnake can target vulnerable body parts of the prey at which it strikes, other snakes with the organ may detect warm bodies from a meter away. It may be used in thermoregulation and predator detection. (See
Absorption spectroscopy
Absorption spectroscopy refers to spectroscopic techniques that measure the absorption of radiation, as a function of frequency or wavelength, due to its interaction with a sample. The sample absorbs i.e. photons, from the radiating field. The intensity of the absorption varies as a function of frequency, this variation is the absorption spectrum. Absorption spectroscopy is performed across the electromagnetic spectrum. Absorption spectroscopy is employed as an analytical chemistry tool to determine the presence of a particular substance in a sample and, in many cases, to quantify the amount of the substance present. Infrared and ultraviolet-visible spectroscopy are common in analytical applications. Absorption spectroscopy is employed in studies of molecular and atomic physics, astronomical spectroscopy and remote sensing. There are a wide range of experimental approaches for measuring absorption spectra; the most common arrangement is to direct a generated beam of radiation at a sample and detect the intensity of the radiation that passes through it.
The transmitted energy can be used to calculate the absorption. The source, sample arrangement and detection technique vary depending on the frequency range and the purpose of the experiment. A material's absorption spectrum is the fraction of incident radiation absorbed by the material over a range of frequencies; the absorption spectrum is determined by the atomic and molecular composition of the material. Radiation is more to be absorbed at frequencies that match the energy difference between two quantum mechanical states of the molecules; the absorption that occurs due to a transition between two states is referred to as an absorption line and a spectrum is composed of many lines. The frequencies where absorption lines occur, as well as their relative intensities depend on the electronic and molecular structure of the sample; the frequencies will depend on the interactions between molecules in the sample, the crystal structure in solids, on several environmental factors. The lines will have a width and shape that are determined by the spectral density or the density of states of the system.
Absorption lines are classified by the nature of the quantum mechanical change induced in the molecule or atom. Rotational lines, for instance, occur. Rotational lines are found in the microwave spectral region. Vibrational lines correspond to changes in the vibrational state of the molecule and are found in the infrared region. Electronic lines correspond to a change in the electronic state of an atom or molecule and are found in the visible and ultraviolet region. X-ray absorptions are associated with the excitation of inner shell electrons in atoms; these changes can be combined, leading to new absorption lines at the combined energy of the two changes. The energy associated with the quantum mechanical change determines the frequency of the absorption line but the frequency can be shifted by several types of interactions. Electric and magnetic fields can cause a shift. Interactions with neighboring molecules can cause shifts. For instance, absorption lines of the gas phase molecule can shift when that molecule is in a liquid or solid phase and interacting more with neighboring molecules.
The width and shape of absorption lines are determined by the instrument used for the observation, the material absorbing the radiation and the physical environment of that material. It is common for lines to have the shape of a Lorentzian distribution, it is common for a line to be described by its intensity and width instead of the entire shape being characterized. The integrated intensity—obtained by integrating the area under the absorption line—is proportional to the amount of the absorbing substance present; the intensity is related to the temperature of the substance and the quantum mechanical interaction between the radiation and the absorber. This interaction is quantified by the transition moment and depends on the particular lower state the transition starts from, the upper state it is connected to; the width of absorption lines may be determined by the spectrometer used to record it. A spectrometer has an inherent limit on how narrow a line it can resolve and so the observed width may be at this limit.
If the width is larger than the resolution limit it is determined by the environment of the absorber. A liquid or solid absorber, in which neighboring molecules interact with one another, tends to have broader absorption lines than a gas. Increasing the temperature or pressure of the absorbing material will tend to increase the line width, it is common for several neighboring transitions to be close enough to one another that their lines overlap and the resulting overall line is therefore broader yet. Absorption and transmission spectra represent equivalent information and one can be calculated from the other through a mathematical transformation. A transmission spectrum will have its maximum intensities at wavelengths where the absorption is weakest because more light is transmitted through the sample. An absorption spectrum will have its maximum intensities at wavelengths where the absorption is strongest. Emission is a process. Emission can occur at any frequency at which absorption can occur, this allows the absorption lines to be determined from an emission spectrum.
The emission spectrum will have a quite different intensity pattern from the absorption spectrum, though
Thallium
Thallium is a chemical element with symbol Tl and atomic number 81. It is a gray post-transition metal, not found free in nature; when isolated, thallium discolors when exposed to air. Chemists William Crookes and Claude-Auguste Lamy discovered thallium independently in 1861, in residues of sulfuric acid production. Both used the newly developed method of flame spectroscopy, in which thallium produces a notable green spectral line. Thallium, from Greek θαλλός, thallós, meaning "a green twig", was named by Crookes, it was isolated by both Lamy and Crookes in 1862. Crookes exhibited it as a powder precipitated by zinc at the International exhibition, which opened on 1 May that year. Thallium tends to oxidize to the +1 oxidation states as ionic salts; the +3 state resembles that of the other elements in group 13. However, the +1 state, far more prominent in thallium than the elements above it, recalls the chemistry of alkali metals, thallium ions are found geologically in potassium-based ores, are handled in many ways like potassium ions by ion pumps in living cells.
Commercially, thallium is produced not from potassium ores, but as a byproduct from refining of heavy-metal sulfide ores. 60–70% of thallium production is used in the electronics industry, the remainder is used in the pharmaceutical industry and in glass manufacturing. It is used in infrared detectors; the radioisotope thallium-201 is used in small, nontoxic amounts as an agent in a nuclear medicine scan, during one type of nuclear cardiac stress test. Soluble thallium salts are toxic, they were used in rat poisons and insecticides. Use of these compounds has been restricted or banned in many countries, because of their nonselective toxicity. Thallium poisoning results in hair loss, although this characteristic symptom does not always surface; because of its historic popularity as a murder weapon, thallium has gained notoriety as "the poisoner's poison" and "inheritance powder". A thallium atom has 81 electrons, arranged in the electron configuration 4f145d106s26p1. Due to the inert pair effect, the 6s electron pair is relativistically stabilised and it is more difficult to get them involved in chemical bonding than for the heavier elements.
Thus few electrons are available for metallic bonding, similar to the neighboring elements mercury and lead, hence thallium, like its congeners, is a soft electrically conducting metal with a low melting point of 304 °C. A number of standard electrode potentials, depending on the reaction under study, are reported for thallium, reflecting the decreased stability of the +3 oxidation state: Thallium is the first element in group 13 where the reduction of the +3 oxidation state to the +1 oxidation state is spontaneous under standard conditions. Since bond energies decrease down the group, with thallium, the energy released in forming two additional bonds and attaining the +3 state is not always enough to outweigh the energy needed to involve the 6s-electrons. Accordingly, thallium oxide and hydroxide are more basic and thallium oxide and hydroxide are more acidic, showing that thallium conforms to the general rule of elements being more electropositive in their lower oxidation states. Thallium is sectile enough to be cut with a knife at room temperature.
It has a metallic luster that, when exposed to air tarnishes to a bluish-gray tinge, resembling lead. It may be preserved by immersion in oil. A heavy layer of oxide builds up on thallium. In the presence of water, thallium hydroxide is formed. Sulfuric and nitric acids dissolve thallium to make the sulfate and nitrate salts, while hydrochloric acid forms an insoluble thallium chloride layer. Thallium has 25 isotopes which have atomic masses that range from 184 to 210. 203Tl and 205Tl make up nearly all of natural thallium. 204Tl is the most stable radioisotope, with a half-life of 3.78 years. It is made by the neutron activation of stable thallium in a nuclear reactor; the most useful radioisotope, 201Tl, decays by electron capture, emitting X-rays, photons of 135 and 167 keV in 10% total abundance. It is the most popular isotope used for thallium nuclear cardiac stress tests. Thallium compounds resemble the corresponding aluminium compounds, they are moderately strong oxidizing agents and are unstable, as illustrated by the positive reduction potential for the Tl3+/Tl couple.
Some mixed-valence compounds are known, such as Tl4O3 and TlCl2, which contain both thallium and thallium. Thallium oxide, Tl2O3, is a black solid which decomposes above 800 °C, forming the thallium oxide and oxygen; the simplest possible thallium compound, thallane, is too unstable to exist in bulk, both due to the instability of the +3 oxidation state as well as poor overlap of the valence 6s and 6p orbitals of thallium with the 1s orbital of hydrogen. The trihalides are more stable, although they are chemically distinct from those of the lighter group 13 elements and are still the least stable in the whole group. For instance, thallium fluoride, TlF3, has the β-BiF3 structure rather than that of the lighter group 13 trifluorides, does not form the TlF−4 complex anion in aqueous solution; the trichloride and tribromide disproportionate just above room
Caesium
Caesium or cesium is a chemical element with symbol Cs and atomic number 55. It is a soft, silvery-golden alkali metal with a melting point of 28.5 °C, which makes it one of only five elemental metals that are liquid at or near room temperature. Caesium has chemical properties similar to those of rubidium and potassium; the most reactive of all metals, it is pyrophoric and reacts with water at −116 °C. It is the least electronegative element, with a value of 0.79 on the Pauling scale. It has only one stable isotope, caesium-133. Caesium is mined from pollucite, while the radioisotopes caesium-137, a fission product, are extracted from waste produced by nuclear reactors; the German chemist Robert Bunsen and physicist Gustav Kirchhoff discovered caesium in 1860 by the newly developed method of flame spectroscopy. The first small-scale applications for caesium were as a "getter" in vacuum tubes and in photoelectric cells. In 1967, acting on Einstein's proof that the speed of light is the most constant dimension in the universe, the International System of Units used two specific wave counts from an emission spectrum of caesium-133 to co-define the second and the metre.
Since caesium has been used in accurate atomic clocks. Since the 1990s, the largest application of the element has been as caesium formate for drilling fluids, but it has a range of applications in the production of electricity, in electronics, in chemistry; the radioactive isotope caesium-137 has a half-life of about 30 years and is used in medical applications, industrial gauges, hydrology. Nonradioactive caesium compounds are only mildly toxic, but the pure metal's tendency to react explosively with water means that caesium is considered a hazardous material, the radioisotopes present a significant health and ecological hazard in the environment. Caesium is the softest element, it is a ductile, pale metal, which darkens in the presence of trace amounts of oxygen. When in the presence of mineral oil, it loses its metallic lustre and takes on a duller, grey appearance, it has a melting point of 28.5 °C, making it one of the few elemental metals that are liquid near room temperature. Mercury is the only elemental metal with a known melting point lower than caesium.
In addition, the metal has a rather low boiling point, 641 °C, the lowest of all metals other than mercury. Its compounds burn with a blue or violet colour. Caesium forms alloys with the other alkali metals and mercury. At temperatures below 650 °C, it does not alloy with cobalt, molybdenum, platinum, tantalum, or tungsten, it forms well-defined intermetallic compounds with antimony, gallium and thorium, which are photosensitive. It mixes with all the other alkali metals. A few amalgams have been studied: CsHg2 is black with a purple metallic lustre, while CsHg is golden-coloured with a metallic lustre; the golden colour of caesium comes from the decreasing frequency of light required to excite electrons of the alkali metals as the group is descended. For lithium through rubidium this frequency is in the ultraviolet, but for caesium it enters the blue–violet end of the spectrum, thus caesium transmits and absorbs violet light preferentially while other colours are reflected. Caesium metal is reactive and pyrophoric.
It ignites spontaneously in air, reacts explosively with water at low temperatures, more so than the other alkali metals. It reacts with solid water at temperatures as low as −116 °C; because of this high reactivity, caesium metal is classified as a hazardous material. It is shipped in dry, saturated hydrocarbons such as mineral oil, it can be handled only under inert gas, such as argon. However, a caesium-water explosion is less powerful than a sodium-water explosion with a similar amount of sodium; this is because caesium explodes upon contact with water, leaving little time for hydrogen to accumulate. Caesium can be stored in vacuum-sealed borosilicate glass ampoules. In quantities of more than about 100 grams, caesium is shipped in hermetically sealed, stainless steel containers; the chemistry of caesium is similar to that of other alkali metals, in particular rubidium, the element above caesium in the periodic table. As expected for an alkali metal, the only common oxidation state is +1; some small differences arise from the fact that it has a higher atomic mass and is more electropositive than other alkali metals.
Caesium is the most electropositive chemical element. The caesium ion is larger and less "hard" than those of the lighter alkali metals. Most caesium compounds contain the element as the cation Cs+, which binds ionically to a wide variety of anions. One noteworthy exception is the caeside anion, others are the several suboxides. Salts of Cs+ are colourless unless the anion itself is coloured. Many of the simple salts are hygroscopic, but less so than the corresponding salts of lighter alkali metals; the phosphate, carbonate, oxide and sulfate salts are water-soluble. Double salts are less soluble, the low solubility of caesium aluminium sulfate is exploited in refining Cs from ores; the double salt with antimony (such as
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