Jeremiah called the "weeping prophet", was one of the major prophets of the Hebrew Bible. According to Jewish tradition, Jeremiah authored the Book of Jeremiah, the Books of Kings and the Book of Lamentations, with the assistance and under the editorship of Baruch ben Neriah, his scribe and disciple. Greater detail is known about Jeremiah's life than for that of any other prophet. However, no biography of him can be written. Judaism considers the Book of Jeremiah part of its canon, regards Jeremiah as the second of the major prophets. Christianity regards Jeremiah as a prophet, he is quoted in the New Testament. Islam considers Jeremiah a prophet, his narrative is given in Islamic tradition. Jeremiah's ministry was active from the thirteenth year of Josiah, king of Judah, until after the fall of Jerusalem and the destruction of Solomon's Temple in 587 BC; this period spanned the reigns of five kings of Judah: Josiah, Jehoiakim and Zedekiah. Jeremiah was the son of a kohen from the Benjamite village of Anathoth.
The difficulties he encountered, as described in the books of Jeremiah and Lamentations, have prompted scholars to refer to him as "the weeping prophet". Jeremiah was called to prophetic ministry c. 626 BC by YHWH to give prophecy of Jerusalem's destruction that would occur by invaders from the north. This was because Israel had been unfaithful to the laws of the covenant and had forsaken God by worshiping Baal. Jeremiah condemned people burning their children as offerings to Moloch; this nation had deviated so far from God that they had broken the covenant, causing God to withdraw his blessings. Jeremiah was guided by God to proclaim that the nation of Judah would be faced with famine and taken captive by foreigners who would exile them to a foreign land; the prophetess Huldah was a relative and contemporary of Jeremiah while the prophets Zephaniah and Isaiah were his mentors. According to Jeremiah 1:2–3, Yahweh called Jeremiah to prophetic ministry in about 626 BC, about five years before Josiah king of Judah turned the nation toward repentance from idolatrous practices.
According to the Books of Kings, Jeremiah, Josiah's reforms were insufficient to save Judah and Jerusalem from destruction, because of the sins of Manasseh, Josiah's grandfather, Judah's return to Idolatry. Such was the lust of the nation for false gods that after Josiah's death, the nation would return to the gods of the surrounding nations. Jeremiah was said to have been appointed to reveal the sins of the people and the coming consequences. Jeremiah did not know how to speak. However, the Lord insisted that Jeremiah go and speak, he touched Jeremiah's mouth to place the word of the Lord there. God told Jeremiah to "Get yourself ready!" The character traits and practices Jeremiah was to acquire are specified in Jeremiah 1 and include not being afraid, standing up to speak, speaking as told, going where sent. Since Jeremiah is described as emerging well trained and literate from his earliest preaching, the relationship between him and the Shaphan family has been used to suggest that he may have trained at the scribal school in Jerusalem over which Shaphan presided.
In his early ministry, Jeremiah was a preaching prophet, preaching throughout Israel. He condemned idolatry, the greed of priests, false prophets. Many years God instructed Jeremiah to write down these early oracles and his other messages. Jeremiah's ministry prompted plots against him. Unhappy with Jeremiah's message for concern that it would shut down the Anathoth sanctuary, his priestly kin and the men of Anathoth conspired to kill him. However, the Lord revealed the conspiracy to Jeremiah, protected his life, declared disaster for the men of Anathoth; when Jeremiah complains to the Lord about this persecution, he is told that the attacks on him will become worse. A priest Pashur the son of ben Immer, a temple official in Jerusalem, had Jeremiah beaten and put in the stocks at the Upper Gate of Benjamin for a day. After this, Jeremiah expresses lament over the difficulty that speaking God's word has caused him and regrets becoming a laughingstock and the target of mockery, he recounts how if he tries to shut the word of the Lord inside and not mention God's name, the word becomes like fire in his heart and he is unable to hold it in.
Whilst Jeremiah was prophesying the coming destruction, a number of other prophets were prophesying peace. Jeremiah spoke against these other prophets. According to the book of Jeremiah, during the reign of King Zedekiah, The Lord instructed Jeremiah to make a yoke of the message that the nation would be subject to the king of Babylon; the prophet Hananiah opposed Jeremiah's message. He took the yoke off of Jeremiah's neck, broke it, prophesied to the priests and all the people that within two years the Lord would break the yoke of the king of Babylon, but the Lord spoke to Jeremiah saying "Go and speak to Hananiah saying, you have broken the yoke of wood, but you have made instead a yoke of iron." Jeremiah was sympathetic to as well as descended from the Northern Kingdom. Many of his first reported oracles are about, addressed to, the Israelites at Samaria, he resembles the northern prophet Hosea, in his use of language, examples of God's relationship to Israel. Hosea seems to have been the first prophet to describe the desired relationship as an example of ancient Israelite marriage, where a man might be polygynous, while a woman was only permitted one husband.
Jeremiah repeats Hosea's marital imagery (Jeremiah 2:2b–2:3.
Transparency and translucency
In the field of optics, transparency is the physical property of allowing light to pass through the material without being scattered. On a macroscopic scale, the photons can be said to follow Snell's Law. Translucency is a superset of transparency: it allows light to pass through, but does not follow Snell's law. In other words, a translucent medium allows the transport of light while a transparent medium not only allows the transport of light but allows for image formation. Transparent materials appear clear, with the overall appearance of one color, or any combination leading up to a brilliant spectrum of every color; the opposite property of translucency is opacity. When light encounters a material, it can interact with it in several different ways; these interactions depend on the nature of the material. Photons interact with an object by some combination of reflection and transmission; some materials, such as plate glass and clean water, transmit much of the light that falls on them and reflect little of it.
Many liquids and aqueous solutions are transparent. Absence of structural defects and molecular structure of most liquids are responsible for excellent optical transmission. Materials which do not transmit light are called opaque. Many such substances have a chemical composition which includes what are referred to as absorption centers. Many substances are selective in their absorption of white light frequencies, they absorb certain portions of the visible spectrum while reflecting others. The frequencies of the spectrum which are not absorbed are either reflected or transmitted for our physical observation; this is. The attenuation of light of all frequencies and wavelengths is due to the combined mechanisms of absorption and scattering. Transparency can provide perfect camouflage for animals able to achieve it; this is easier in turbid seawater than in good illumination. Many marine animals such as jellyfish are transparent. With regard to the absorption of light, primary material considerations include: At the electronic level, absorption in the ultraviolet and visible portions of the spectrum depends on whether the electron orbitals are spaced such that they can absorb a quantum of light of a specific frequency, does not violate selection rules.
For example, in most glasses, electrons have no available energy levels above them in range of that associated with visible light, or if they do, they violate selection rules, meaning there is no appreciable absorption in pure glasses, making them ideal transparent materials for windows in buildings. At the atomic or molecular level, physical absorption in the infrared portion of the spectrum depends on the frequencies of atomic or molecular vibrations or chemical bonds, on selection rules. Nitrogen and oxygen are not greenhouse gases because there is no absorption, but because there is no molecular dipole moment. With regard to the scattering of light, the most critical factor is the length scale of any or all of these structural features relative to the wavelength of the light being scattered. Primary material considerations include: Crystalline structure: whether or not the atoms or molecules exhibit the'long-range order' evidenced in crystalline solids. Glassy structure: scattering centers include fluctuations in density or composition.
Microstructure: scattering centers include internal surfaces such as grain boundaries, crystallographic defects and microscopic pores. Organic materials: scattering centers include fiber and cell structures and boundaries. Diffuse reflection - Generally, when light strikes the surface of a solid material, it bounces off in all directions due to multiple reflections by the microscopic irregularities inside the material, by its surface, if it is rough. Diffuse reflection is characterized by omni-directional reflection angles. Most of the objects visible to the naked eye are identified via diffuse reflection. Another term used for this type of reflection is "light scattering". Light scattering from the surfaces of objects is our primary mechanism of physical observation. Light scattering in liquids and solids depends on the wavelength of the light being scattered. Limits to spatial scales of visibility therefore arise, depending on the frequency of the light wave and the physical dimension of the scattering center.
Visible light has a wavelength scale on the order of a half a micrometer. Scattering centers as small. Optical transparency in polycrystalline materials is limited by the amount of light, scattered by their microstructural features. Light scattering depends on the wavelength of the light. Limits to spatial scales of visibility therefore arise, depending on the frequency of the light wave and the physical dimension of the scattering center. For example, since visible light has a wavelength scale on the order of a micrometer, scattering centers will have dimensions on a similar spatial scale. Primary scattering centers in polycrystalline materi
Specific gravity is the ratio of the density of a substance to the density of a reference substance. Apparent specific gravity is the ratio of the weight of a volume of the substance to the weight of an equal volume of the reference substance; the reference substance for liquids is nearly always water at its densest. Nonetheless, the temperature and pressure must be specified for the reference. Pressure is nearly always 1 atm. Temperatures for both sample and reference vary from industry to industry. In British beer brewing, the practice for specific gravity as specified above is to multiply it by 1,000. Specific gravity is used in industry as a simple means of obtaining information about the concentration of solutions of various materials such as brines, antifreeze coolants, sugar solutions and acids. Being a ratio of densities, specific gravity is a dimensionless quantity; the reason for the specific gravity being dimensionless is to provide a global consistency between the U. S. and Metric Systems, since various units for density may be used such as pounds per cubic feet or grams per cubic centimeter, etc.
Specific gravity varies with pressure. Substances with a specific gravity of 1 are neutrally buoyant in water; those with SG greater than 1 are denser than water and will, disregarding surface tension effects, sink in it. Those with an SG less than 1 will float on it. In scientific work, the relationship of mass to volume is expressed directly in terms of the density of the substance under study, it is in industry where specific gravity finds wide application for historical reasons. True specific gravity can be expressed mathematically as: S G true = ρ sample ρ H 2 O where ρsample is the density of the sample and ρH2O is the density of water; the apparent specific gravity is the ratio of the weights of equal volumes of sample and water in air: S G apparent = W A, sample W A, H 2 O where WA,sample represents the weight of the sample measured in air and WA,H2O the weight of water measured in air. It can be shown that true specific gravity can be computed from different properties: S G true = ρ sample ρ H 2 O = m sample V m H 2 O V = m sample m H 2 O g g = W V, sample W V, H 2 O where g is the local acceleration due to gravity, V is the volume of the sample and of water, ρsample is the density of the sample, ρH2O is the density of water and WV represents a weight obtained in vacuum.
The density of water varies with pressure as does the density of the sample. So it is necessary to specify the temperatures and pressures at which the densities or weights were determined, it is nearly always the case. But as specific gravity refers to incompressible aqueous solutions or other incompressible substances, variations in density caused by pressure are neglected at least where apparent specific gravity is being measured. For true specific gravity calculations, air pressure must be considered. Temperatures are specified by the notation, with Ts representing the temperature at which the sample's density was determined and Tr the temperature at which the reference density is specified. For example, SG would be understood to mean that the density of the sample was determined at 20 °C and of the water at 4 °C. Taking into account different sample and reference temperatures, we note that, while SGH2O = 1.000000, it is the case that SGH2O = 0.998203⁄0.999840 = 0.998363. Here, temperature is being specified using the current ITS-90 scale and the densities used here and in the rest of this article are based on that scale.
On the previous IPTS-68 scale, the densities at 20 °C and 4 °C are 0.9982071 and 0.9999720 respective
Gunpowder known as black powder to distinguish it from modern smokeless powder, is the earliest known chemical explosive. It consists of a mixture of sulfur and potassium nitrate; the sulfur and charcoal act as fuels. Because of its incendiary properties and the amount of heat and gas volume that it generates, gunpowder has been used as a propellant in firearms, artillery and fireworks, as a blasting powder in quarrying and road building. Gunpowder was invented in 9th-century China and spread throughout most parts of Eurasia by the end of the 13th century. Developed by the Taoists for medicinal purposes, gunpowder was first used for warfare about 1000 AD. Gunpowder is classified as a low explosive because of its slow decomposition rate and low brisance. Low explosives deflagrate at subsonic speeds, whereas high explosives detonate, producing a supersonic wave. Ignition of gunpowder packed behind a projectile generates enough pressure to force the shot from the muzzle at high speed, but not enough force to rupture the gun barrel.
Gunpowder thus makes a good propellant, but is less suitable for shattering rock or fortifications with its low-yield explosive power. However, by transferring enough energy a bombardier may wear down an opponent's fortified defenses. Gunpowder was used to fill fused artillery shells until the second half of the 19th century, when the first high explosives were put into use. Gunpowder is no longer used in modern weapons, nor is it used for industrial purposes, due to its inefficient cost compared to newer alternatives such as dynamite and ammonium nitrate/fuel oil. Today gunpowder firearms are limited to hunting, target shooting, bulletless historical reenactments. Based on a 9th-century Taoist text, the invention of gunpowder by Chinese alchemists was an accidental byproduct from experiments seeking to create the elixir of life; this experimental medicine origin of gunpowder is reflected in its Chinese name huoyao, which means "fire medicine". The first military applications of gunpowder were developed around 1000 AD.
The earliest chemical formula for gunpowder appeared in the 11th century Song dynasty text, Wujing Zongyao, however gunpowder had been used for fire arrows since at least the 10th century. In the following centuries various gunpowder weapons such as bombs, fire lances, the gun appeared in China. Saltpeter was known to the Chinese by the mid-1st century AD and was produced in the provinces of Sichuan and Shandong. There is strong evidence of the use of sulfur in various medicinal combinations. A Chinese alchemical text dated 492 noted saltpeter burnt with a purple flame, providing a practical and reliable means of distinguishing it from other inorganic salts, thus enabling alchemists to evaluate and compare purification techniques; the first reference to the incendiary properties of such mixtures is the passage of the Zhenyuan miaodao yaolüe, a Taoist text tentatively dated to the mid-9th century: "Some have heated together sulfur and saltpeter with honey. The Chinese word for "gunpowder" is Chinese: 火药/火藥.
In the following centuries a variety of gunpowder weapons such as rockets and land mines appeared before the first metal barrel firearms were invented. Explosive weapons such as bombs have been discovered in a shipwreck off the shore of Japan dated from 1281, during the Mongol invasions of Japan; the Chinese Wujing Zongyao, written by Zeng Gongliang between 1040 and 1044, provides encyclopedia references to a variety of mixtures that included petrochemicals—as well as garlic and honey. A slow match for flame throwing mechanisms using the siphon principle and for fireworks and rockets is mentioned; the mixture formulas in this book do not contain enough saltpeter to create an explosive however. The Essentials was however written by a Song dynasty court bureaucrat, there is little evidence that it had any immediate impact on warfare. However, by 1083 the Song court was producing hundreds of thousands of fire arrows for their garrisons. Bombs and the first proto-guns, known as "fire lances", became prominent during the 12th century and were used by the Song during the Jin-Song Wars.
Fire lances were first recorded to have been used at the Siege of De'an in 1132 by Song forces against the Jin. In the early 13th century the Jin utilized iron-casing bombs. Projectiles were added to fire lances, re-usable fire lance barrels were developed, first out of hardened paper, metal. By 1257 some fire lances were firing wads of bullets. In the late 13th century metal fire lances became'eruptors', proto-cannons firing co-viative projectiles, by 1287 at the latest, had become true guns, the hand cannon; the earliest Western accounts of gunpowder appear in texts written by English philosopher Roger Bacon in the 13th century. Several sources men
Solubility is the property of a solid, liquid or gaseous chemical substance called solute to dissolve in a solid, liquid or gaseous solvent. The solubility of a substance fundamentally depends on the physical and chemical properties of the solute and solvent as well as on temperature and presence of other chemicals of the solution; the extent of the solubility of a substance in a specific solvent is measured as the saturation concentration, where adding more solute does not increase the concentration of the solution and begins to precipitate the excess amount of solute. Insolubility is the inability to dissolve in a liquid or gaseous solvent. Most the solvent is a liquid, which can be a pure substance or a mixture. One may speak of solid solution, but of solution in a gas. Under certain conditions, the equilibrium solubility can be exceeded to give a so-called supersaturated solution, metastable. Metastability of crystals can lead to apparent differences in the amount of a chemical that dissolves depending on its crystalline form or particle size.
A supersaturated solution crystallises when'seed' crystals are introduced and rapid equilibration occurs. Phenylsalicylate is one such simple observable substance when melted and cooled below its fusion point. Solubility is not to be confused with the ability to'dissolve' a substance, because the solution might occur because of a chemical reaction. For example, zinc'dissolves' in hydrochloric acid as a result of a chemical reaction releasing hydrogen gas in a displacement reaction; the zinc ions are soluble in the acid. The solubility of a substance is an different property from the rate of solution, how fast it dissolves; the smaller a particle is, the faster it dissolves although there are many factors to add to this generalization. Crucially solubility applies to all areas of chemistry, inorganic, physical and biochemistry. In all cases it will depend on the physical conditions and the enthalpy and entropy directly relating to the solvents and solutes concerned. By far the most common solvent in chemistry is water, a solvent for most ionic compounds as well as a wide range of organic substances.
This is a crucial factor in much environmental and geochemical work. According to the IUPAC definition, solubility is the analytical composition of a saturated solution expressed as a proportion of a designated solute in a designated solvent. Solubility may be stated in various units of concentration such as molarity, mole fraction, mole ratio, mass per volume and other units; the extent of solubility ranges from infinitely soluble such as ethanol in water, to poorly soluble, such as silver chloride in water. The term insoluble is applied to poorly or poorly soluble compounds. A number of other descriptive terms are used to qualify the extent of solubility for a given application. For example, U. S. Pharmacopoeia gives the following terms: The thresholds to describe something as insoluble, or similar terms, may depend on the application. For example, one source states that substances are described as "insoluble" when their solubility is less than 0.1 g per 100 mL of solvent. Solubility occurs under dynamic equilibrium, which means that solubility results from the simultaneous and opposing processes of dissolution and phase joining.
The solubility equilibrium occurs. The term solubility is used in some fields where the solute is altered by solvolysis. For example, many metals and their oxides are said to be "soluble in hydrochloric acid", although in fact the aqueous acid irreversibly degrades the solid to give soluble products, it is true that most ionic solids are dissolved by polar solvents, but such processes are reversible. In those cases where the solute is not recovered upon evaporation of the solvent, the process is referred to as solvolysis; the thermodynamic concept of solubility does not apply straightforwardly to solvolysis. When a solute dissolves, it may form several species in the solution. For example, an aqueous suspension of ferrous hydroxide, Fe2, will contain the series + as well as other species. Furthermore, the solubility of ferrous hydroxide and the composition of its soluble components depend on pH. In general, solubility in the solvent phase can be given only for a specific solute, thermodynamically stable, the value of the solubility will include all the species in the solution.
Solubility is defined for specific phases. For example, the solubility of aragonite and calcite in water are expected to differ though they are both polymorphs of calcium carbonate and have the same chemical formula; the solubility of one substance in another is determined by the balance of intermolecular forces between the solvent and solute, the entropy change that accompanies the solvation. Factors such as temperature and pressure will alter this balance. Solubility may strongly depend on the presence of other species dissolved in the solvent, for example, complex-forming anions in liquids. Solubility will depend on the excess or deficiency of a common ion in the solution, a phenomenon known as the common-ion effect. To a lesser extent, solubility will depend on the ionic strength of solutions; the last two effects can be quantified using the equation for solubility equilibrium. For a solid that dissolves in a redox reaction, solubility is expe
Borates are the name for a large number of boron-containing oxyanions. The term "borates" may refer to tetrahedral boron anions, or more loosely to chemical compounds which contain borate anions of either description. Larger borates are composed of trigonal planar BO3 or tetrahedral BO4 structural units, joined together via shared oxygen atoms and may be cyclic or linear in structure. Boron most occurs in nature as borates, such as borate minerals and borosilicates; the simplest borate anion, the orthoborate ion, 3-, is known in the solid state, for example in Ca32. In this it adopts a near trigonal planar structure, it is a structural analogue of the carbonate anion 2 -. Simple bonding theories point to the trigonal planar structure. In terms of valence bond theory the bonds are formed by using sp2 hybrid orbitals on boron; some compounds termed orthoborates do not contain the trigonal planar ion, for example gadolinium orthoborate, GdBO3 contains the polyborate 9- ion, whereas the high temperature form contains planar 3-.
All borates can be considered derivatives of boric acid, B3. Boric acid is a weak proton donor in the sense of Brønsted acid, but is a Lewis acid, i.e. it can accept an electron pair. In water, it behaves as a Lewis acid accepting the electron pair of a hydroxyl ion produced by the water autoprotolysis. So, B3 is acidic because of its reaction with OH– from water, forming the tetrahydroxyborate complex − and releasing the corresponding proton left by the water autoprotolysis: B3 + 2H2O ⇌ − + + In the presence of cis-diols such as mannitol, glucose and glycerol the pKa is lowered to about 4. At neutral pH boric acid undergoes condensation reactions to form polymeric oxyanions. Well-known polyborate anions include the triborate and pentaborate anions; the condensation reaction for the formation of tetraborate is as follows: 2 B3 + 2 − ⇌ 2- + 5 H2OThe tetraborate anion includes two tetrahedral and two trigonal boron atoms symmetrically assembled in a fused bicyclic structure. The two tetrahedral boron atoms are linked together by a common oxygen atom and each bears a negative net charge brought by the supplementary OH− groups laterally attached to them.
This intricate molecular anion exhibits three rings: two fused distorted hexagonal rings and one distorted octagonal ring. Each ring is made of a succession of alternate oxygen atoms. Boroxole rings are a common structural motif in polyborate ions; the tetraborate anion occurs in the mineral borax, as an octahydrate, Na2·8H2O. The borax chemical formula is commonly written in a more compact notation as Na2B4O7·10H2O. Sodium borate can be obtained in high purity and so can be used to make a standard solution in titrimetric analysis. A number of metal borates are known, they arise by treating boric boron oxides with metal oxides. Examples hereafter include linear chains of 2, 3 or 4 trigonal BO3 structural units, each sharing only one oxygen atom with adjacent unit: diborate 4-, found in Mg2B2O5 triborate 5-, found in CaAlB3O7 tetraborate 6-, found in Li6B4O9Metaborates, such as LiBO2 contain chains of trigonal BO3 structural units, each sharing two oxygen atoms with adjacent units, whereas NaBO2 and KBO2 contain the cyclic 2- ion.
Borosilicate glass known as pyrex, can be viewed as a silicate in which some 4- units are replaced by 5- centers, together with additional cations to compensate for the difference in valence states of Si and B. Because of this substitution leads to imperfections, the material is slow to crystallise and forms a glass with low coefficient of thermal expansion and is resistant to cracking when heated, unlike soda glass. Common borate salts include sodium metaborate, NaBO2, borax. Borax is soluble in water, so mineral deposits only occur in places with low rainfall. Extensive deposits were found in Death Valley and transported out using the famous twenty-mule teams. Deposits were found at Boron, California on the edge of the Mojave Desert; the Atacama Desert in Chile contains mineable borate concentrations. Lithium metaborate or lithium tetraborate, or a mixture of both, can be used in borate fusion sample preparation of various samples for analysis by XRF, AAS, ICP-OES, ICP-AES and ICP-MS. Borate fusion and energy dispersive X-ray fluorescence spectrometry with polarized excitation have been used in the analysis of contaminated soils.
Disodium octaborate tetrahydrate is used as fungicide. Zinc borate is used as a flame retardant. Borate esters are organic compounds which are conveniently prepared by the stoichiometric condensation reaction of boric acid with alcohols. Borax Nanochannel glass materials Porous glass Vycor glass Tris borate Suanite at webmineral Johachidolite at webmineral Non-CCA Wood Preservatives: Guide to Selected Resources - National Pesticide Information Center
Nitrogen is a chemical element with symbol N and atomic number 7. It was first discovered and isolated by Scottish physician Daniel Rutherford in 1772. Although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is accorded the credit because his work was published first; the name nitrogène was suggested by French chemist Jean-Antoine-Claude Chaptal in 1790, when it was found that nitrogen was present in nitric acid and nitrates. Antoine Lavoisier suggested instead the name azote, from the Greek ἀζωτικός "no life", as it is an asphyxiant gas. Nitrogen is the lightest member of group 15 of the periodic table called the pnictogens; the name comes from the Greek πνίγειν "to choke", directly referencing nitrogen's asphyxiating properties. It is a common element in the universe, estimated at about seventh in total abundance in the Milky Way and the Solar System. At standard temperature and pressure, two atoms of the element bind to form dinitrogen, a colourless and odorless diatomic gas with the formula N2.
Dinitrogen forms about 78 % of Earth's atmosphere. Nitrogen occurs in all organisms in amino acids, in the nucleic acids and in the energy transfer molecule adenosine triphosphate; the human body contains about 3% nitrogen by mass, the fourth most abundant element in the body after oxygen and hydrogen. The nitrogen cycle describes movement of the element from the air, into the biosphere and organic compounds back into the atmosphere. Many industrially important compounds, such as ammonia, nitric acid, organic nitrates, cyanides, contain nitrogen; the strong triple bond in elemental nitrogen, the second strongest bond in any diatomic molecule after carbon monoxide, dominates nitrogen chemistry. This causes difficulty for both organisms and industry in converting N2 into useful compounds, but at the same time means that burning, exploding, or decomposing nitrogen compounds to form nitrogen gas releases large amounts of useful energy. Synthetically produced ammonia and nitrates are key industrial fertilisers, fertiliser nitrates are key pollutants in the eutrophication of water systems.
Apart from its use in fertilisers and energy-stores, nitrogen is a constituent of organic compounds as diverse as Kevlar used in high-strength fabric and cyanoacrylate used in superglue. Nitrogen is a constituent including antibiotics. Many drugs are mimics or prodrugs of natural nitrogen-containing signal molecules: for example, the organic nitrates nitroglycerin and nitroprusside control blood pressure by metabolizing into nitric oxide. Many notable nitrogen-containing drugs, such as the natural caffeine and morphine or the synthetic amphetamines, act on receptors of animal neurotransmitters. Nitrogen compounds have a long history, ammonium chloride having been known to Herodotus, they were well known by the Middle Ages. Alchemists knew nitric acid as aqua fortis, as well as other nitrogen compounds such as ammonium salts and nitrate salts; the mixture of nitric and hydrochloric acids was known as aqua regia, celebrated for its ability to dissolve gold, the king of metals. The discovery of nitrogen is attributed to the Scottish physician Daniel Rutherford in 1772, who called it noxious air.
Though he did not recognise it as an different chemical substance, he distinguished it from Joseph Black's "fixed air", or carbon dioxide. The fact that there was a component of air that does not support combustion was clear to Rutherford, although he was not aware that it was an element. Nitrogen was studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as "mephitic air" or azote, from the Greek word άζωτικός, "no life". In an atmosphere of pure nitrogen, animals died and flames were extinguished. Though Lavoisier's name was not accepted in English, since it was pointed out that all gases are mephitic, it is used in many languages and still remains in English in the common names of many nitrogen compounds, such as hydrazine and compounds of the azide ion, it led to the name "pnictogens" for the group headed by nitrogen, from the Greek πνίγειν "to choke".
The English word nitrogen entered the language from the French nitrogène, coined in 1790 by French chemist Jean-Antoine Chaptal, from the French nitre and the French suffix -gène, "producing", from the Greek -γενής. Chaptal's meaning was that nitrogen is the essential part of nitric acid, which in turn was produced from nitre. In earlier times, niter had been confused with Egyptian "natron" – called νίτρον in Greek – which, despite the name, contained no nitrate; the earliest military and agricultural applications of nitrogen compounds used saltpeter, most notably in gunpowder, as fertiliser. In 1910, Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", a monatomic allotrope of nitrogen; the "whirling cloud of brilliant yellow light