Hydrogen peroxide is a chemical compound with the formula H2O2. In its pure form, it is a pale blue, clear liquid more viscous than water. Hydrogen peroxide is the simplest peroxide, it is used as bleaching agent and antiseptic. Concentrated hydrogen peroxide, or "high-test peroxide", is a reactive oxygen species and has been used as a propellant in rocketry, its chemistry is dominated by the nature of its unstable peroxide bond. Hydrogen peroxide is unstable and decomposes in the presence of light; because of its instability, hydrogen peroxide is stored with a stabilizer in a weakly acidic solution. Hydrogen peroxide is found in biological systems including the human body. Enzymes that use or decompose hydrogen peroxide are classified as peroxidases; the boiling point of H2O2 has been extrapolated as being 150.2 °C 50 °C higher than water. In practice, hydrogen peroxide will undergo explosive thermal decomposition if heated to this temperature, it may be safely distilled at lower temperatures under reduced pressure.
In aqueous solutions hydrogen peroxide differs from the pure substance due to the effects of hydrogen bonding between water and hydrogen peroxide molecules. Hydrogen peroxide and water form a eutectic mixture; the boiling point of the same mixtures is depressed in relation with the mean of both boiling points. It occurs at 114 °C; this boiling point is 14 °C greater than that of pure water and 36.2 °C less than that of pure hydrogen peroxide. Hydrogen peroxide is a nonplanar molecule as shown by Paul-Antoine Giguère in 1950 using infrared spectroscopy, with C2 symmetry. Although the O−O bond is a single bond, the molecule has a high rotational barrier of 2460 cm−1; the increased barrier is ascribed to repulsion between the lone pairs of the adjacent oxygen atoms and results in hydrogen peroxide displaying atropisomerism. The molecular structures of gaseous and crystalline H2O2 are different; this difference is attributed to the effects of hydrogen bonding, absent in the gaseous state. Crystals of H2O2 are tetragonal with the space group D44P4121.
Hydrogen peroxide has several structural analogues with Hm−X−X−Hn bonding arrangements. It has the highest boiling point of this series, its melting point is fairly high, being comparable to that of hydrazine and water, with only hydroxylamine crystallising more indicative of strong hydrogen bonding. Diphosphane and hydrogen disulfide exhibit only weak hydrogen bonding and have little chemical similarity to hydrogen peroxide. All of these analogues are thermodynamically unstable. Structurally, the analogues all adopt similar skewed structures, due to repulsion between adjacent lone pairs. Alexander von Humboldt synthesized one of the first synthetic peroxides, barium peroxide, in 1799 as a by-product of his attempts to decompose air. Nineteen years Louis Jacques Thénard recognized that this compound could be used for the preparation of a unknown compound, which he described as eau oxygénée – subsequently known as hydrogen peroxide. An improved version of Thénard's process used hydrochloric acid, followed by addition of sulfuric acid to precipitate the barium sulfate byproduct.
This process was used from the end of the 19th century until the middle of the 20th century. Thénard and Joseph Louis Gay-Lussac synthesized sodium peroxide in 1811; the bleaching effect of peroxides and their salts on natural dyes became known around that time, but early attempts of industrial production of peroxides failed, the first plant producing hydrogen peroxide was built in 1873 in Berlin. The discovery of the synthesis of hydrogen peroxide by electrolysis with sulfuric acid introduced the more efficient electrochemical method, it was first implemented into industry in 1908 in Weißenstein, Austria. The anthraquinone process, still used, was developed during the 1930s by the German chemical manufacturer IG Farben in Ludwigshafen; the increased demand and improvements in the synthesis methods resulted in the rise of the annual production of hydrogen peroxide from 35,000 tonnes in 1950, to over 100,000 tonnes in 1960, to 300,000 tonnes by 1970. Pure hydrogen peroxide was long believed to be unstable, as early attempts to separate it from the water, present during synthesis, all failed.
This instability was due to traces of impurities, which catalyze the decomposition of the hydrogen peroxide. Pure hydrogen peroxide was first obtained in 1894—almost 80 years after its discovery—by Richard Wolffenstein, who produced it by vacuum distillation. Determination of the molecular structure of hydrogen peroxide proved to be difficult. In 1892 the Italian physical chemist Giacomo Carrara determined its molecular mass by freezing-point depression, which confirmed that its molecular formula is H2O2. At least half a dozen hypothetical molecular structures seemed to be consistent with the available evidence. In 1934, the English mathematical physicist William Penney and the Scottish physicist Gordon Sutherland proposed a molecular structure for hydrogen peroxide, similar to the presently accepted one. Hydrogen peroxide was prepared industrially by hydrolysis of ammonium persulfate, itself obtained by the electrolysis of a solution
A rocket is a missile, aircraft or other vehicle that obtains thrust from a rocket engine. Rocket engine exhaust is formed from propellant carried within the rocket before use. Rocket engines work by action and reaction and push rockets forward by expelling their exhaust in the opposite direction at high speed, can therefore work in the vacuum of space. In fact, rockets work more efficiently in space than in an atmosphere. Multistage rockets are capable of attaining escape velocity from Earth and therefore can achieve unlimited maximum altitude. Compared with airbreathing engines, rockets are lightweight and powerful and capable of generating large accelerations. To control their flight, rockets rely on momentum, auxiliary reaction engines, gimballed thrust, momentum wheels, deflection of the exhaust stream, propellant flow, spin, or gravity. Rockets for military and recreational uses date back to at least 13th-century China. Significant scientific and industrial use did not occur until the 20th century, when rocketry was the enabling technology for the Space Age, including setting foot on the Earth's moon.
Rockets are now used for fireworks, ejection seats, launch vehicles for artificial satellites, human spaceflight, space exploration. Chemical rockets are the most common type of high power rocket creating a high speed exhaust by the combustion of fuel with an oxidizer; the stored propellant can be a simple pressurized gas or a single liquid fuel that disassociates in the presence of a catalyst, two liquids that spontaneously react on contact, two liquids that must be ignited to react, a solid combination of fuel with oxidizer, or solid fuel with liquid oxidizer. Chemical rockets store a large amount of energy in an released form, can be dangerous. However, careful design, testing and use minimizes risks; the first gunpowder-powered rockets evolved in medieval China under the Song dynasty by the 13th century. The Mongols adopted Chinese rocket technology and the invention spread via the Mongol invasions to the Middle East and to Europe in the mid-13th century. Rockets are recorded in use by the Song navy in a military exercise dated to 1245.
Internal-combustion rocket propulsion is mentioned in a reference to 1264, recording that the "ground-rat", a type of firework, had frightened the Empress-Mother Gongsheng at a feast held in her honor by her son the Emperor Lizong. Subsequently, rockets are included in the military treatise Huolongjing known as the Fire Drake Manual, written by the Chinese artillery officer Jiao Yu in the mid-14th century; this text mentions the first known multistage rocket, the'fire-dragon issuing from the water', thought to have been used by the Chinese navy. Medieval and early modern rockets were used militarily as incendiary weapons in sieges. Between 1270 and 1280, Hasan al-Rammah wrote al-furusiyyah wa al-manasib al-harbiyya, which included 107 gunpowder recipes, 22 of them for rockets. In Europe, Konrad Kyeser described rockets in his military treatise Bellifortis around 1405; the name "rocket" comes from the Italian rocchetta, meaning "bobbin" or "little spindle", given due to the similarity in shape to the bobbin or spool used to hold the thread to be fed to a spinning wheel.
Leonhard Fronsperger and Conrad Haas adopted the Italian term into German in the mid-16th century. Artis Magnae Artilleriae pars prima, an important early modern work on rocket artillery, by Kazimierz Siemienowicz, was first printed in Amsterdam in 1650; the Mysorean rockets were the first successful iron-cased rockets, developed in the late 18th century in the Kingdom of Mysore by Tipu Sultan. The Congreve rocket was a British weapon designed and developed by Sir William Congreve in 1804; this rocket was based directly on the Mysorean rockets, used compressed powder and was fielded in the Napoleonic Wars. It was Congreve rockets that Francis Scott Key was referring to when he wrote of the "rockets' red glare" while held captive on a British ship, laying siege to Fort McHenry in 1814. Together, the Mysorean and British innovations increased the effective range of military rockets from 100 to 2,000 yards; the first mathematical treatment of the dynamics of rocket propulsion is due to William Moore.
In 1815 Alexander Dmitrievich Zasyadko constructed rocket-launching platforms, which allowed rockets to be fired in salvos, gun-laying devices. William Hale in 1844 increased the accuracy of rocket artillery. Edward Mounier Boxer further improved the Congreve rocket in 1865. William Leitch first proposed the concept of using rockets to enable human spaceflight in 1861. Konstantin Tsiolkovsky also conceived this idea, extensively developed a body of theory that has provided the foundation for subsequent spaceflight development. Robert Goddard in 1920 published proposed improvements to rocket technology in A Method of Reaching Extreme Altitudes. In 1923, Hermann Oberth published Die Rakete zu den Planetenräumen Modern rockets originated in 1926 when Goddard attached a supersonic nozzle to the combustion chamber of a liquid-propellant rocket; these nozzles turn the hot gas from the combustion chamber into a cooler, hypersonic directed jet of gas, more than doubling the thrust and raising the engine efficiency from 2% to 64%.
Use of liquid propellants instead of gunpowder improved the effectiveness of rocket artillery in World War II, opened up the p
Mars Reconnaissance Orbiter
Mars Reconnaissance Orbiter is a multipurpose spacecraft designed to conduct reconnaissance and exploration of Mars from orbit. The US$720 million spacecraft was built by Lockheed Martin under the supervision of the Jet Propulsion Laboratory; the mission is managed by the California Institute of Technology, at the JPL, in Pasadena, for the NASA Science Mission Directorate, Washington, D. C, it was launched August 12, 2005, attained Martian orbit on March 10, 2006. In November 2006, after five months of aerobraking, it entered its final science orbit and began its primary science phase; as MRO entered orbit, it joined five other active spacecraft that were either in orbit or on the planet's surface: Mars Global Surveyor, Mars Express, 2001 Mars Odyssey, the two Mars Exploration Rovers. Mars Global Surveyor and the rover Spirit have since ceased to function. Opportunity has remained silent since June 10, 2018, NASA declared its mission complete on February 13, 2019; as of that date, 2001 Mars Odyssey and MRO continue to remain operational.
MRO contains a host of scientific instruments such as cameras and radar, which are used to analyze the landforms, stratigraphy and ice of Mars. It paves the way for future spacecraft by monitoring Mars' daily weather and surface conditions, studying potential landing sites, hosting a new telecommunications system. MRO's telecommunications system will transfer more data back to Earth than all previous interplanetary missions combined, MRO will serve as a capable relay satellite for future missions, it has enough propellant to keep functioning into the 2030s. One of two missions considered for the 2003 Mars launch window, the MRO proposal lost against what became known as the Mars Exploration Rovers; the orbiter mission was rescheduled for launch in 2005, NASA announced its final name, Mars Reconnaissance Orbiter, on October 26, 2000. MRO is modeled after NASA's successful Mars Global Surveyor to conduct surveillance of Mars from orbit. Early specifications of the satellite included a large camera to take high resolution pictures of Mars.
In this regard, James B. Garvin, the Mars exploration program scientist for NASA, proclaimed that MRO would be a "microscope in orbit"; the satellite was to include a visible-near-infrared spectrograph. On October 3, 2001, NASA chose Lockheed Martin as the primary contractor for the spacecraft's fabrication. By the end of 2001 all of the mission's instruments were selected. There were no major setbacks during MRO's construction, the spacecraft was moved to John F. Kennedy Space Center on May 1, 2005 to prepare it for launch. MRO science operations were scheduled to last two Earth years, from November 2006 to November 2008. One of the mission's main goals is to map the Martian landscape with its high-resolution cameras in order to choose landing sites for future surface missions; the MRO played an important role in choosing the landing site of the Phoenix Lander, which explored the Martian Arctic in Green Valley. The initial site chosen by scientists was imaged with the HiRISE camera and found to be littered with boulders.
After analysis with HiRISE and the Mars Odyssey's THEMIS instrument a new site was chosen. Mars Science Laboratory, a maneuverable rover had its landing site inspected; the MRO provided critical navigation data during their landings and acts as a telecommunications relay. MRO is using its onboard scientific equipment to study the Martian climate, weather and geology, to search for signs of liquid water in the polar caps and underground. In addition, MRO was tasked with looking for the remains of the lost Mars Polar Lander and Beagle 2 spacecraft. Beagle 2 was found by the orbiter at the beginning of 2015. After its main science operations are completed, the probe's extended mission is to be the communication and navigation system for landers and rover probes. On August 12, 2005, MRO was launched aboard an Atlas V-401 rocket from Space Launch Complex 41 at Cape Canaveral Air Force Station; the Centaur upper stage of the rocket completed its burns over a fifty-six-minute period and placed MRO into an interplanetary transfer orbit towards Mars.
MRO cruised through interplanetary space for seven and a half months before reaching Mars. While en route most of the scientific instruments and experiments were tested and calibrated. To ensure proper orbital insertion upon reaching Mars, four trajectory correction maneuvers were planned and a fifth emergency maneuver was discussed. However, only three trajectory correction maneuvers were necessary, which saved 60 pounds fuel that would be usable during MRO's extended mission. MRO began orbital insertion by approaching Mars on March 10, 2006, passing above its southern hemisphere at an altitude of 370–400 kilometers. All six of MRO's main engines burned for 27 minutes to slow the probe from 2,900 to 1,900 meters per second; the helium pressurization tank was colder than expected, which reduced the pressure in the fuel tank by about 21 kilopascals. The reduced pressure caused the engine thrust to be diminished by 2%, but MRO automatically compensated by extending the burn time by 33 seconds. Completion of the orbital insertion placed the orbiter in a elliptical polar orbit with a period of 35.5 hours.
Shortly after insertion, the periapsis – the point in the orbit closest to Mars – was 426 km from the surface. The apoapsis – the point in the orbit farthest from Mars – was 44,500 km from the surface. On March
In thermodynamics, the term exothermic process describes a process or reaction that releases energy from the system to its surroundings in the form of heat, but in a form of light, electricity, or sound. Its etymology stems from the Greek prefix έξω and the Greek word θερμικός; the term exothermic was first coined by Marcellin Berthelot. The opposite of an exothermic process is an endothermic process, one that absorbs energy in the form of heat; the concept is applied in the physical sciences to chemical reactions, where as in chemical bond energy that will be converted to thermal energy. Exothermic describe two types of chemical systems found in nature, as follows. Stated, after an exothermic reaction, more energy has been released to the surroundings than was absorbed to initiate and maintain the reaction. An example would be the burning of a candle, wherein the sum of calories produced by combustion exceeds the number of calories absorbed in lighting the flame and in the flame maintaining itself..
On the other hand, in an endothermic reaction or system, energy is taken from the surroundings in the course of the reaction. An example of an endothermic reaction is a first aid cold pack, in which the reaction of two chemicals, or dissolving of one in another, requires calories from the surroundings, the reaction cools the pouch and surroundings by absorbing heat from them. An endothermic system is seen in the production of wood: trees absorb radiant energy, from the sun, use it in endothermic reactions such as taking apart CO2 and H2O and combining the carbon and hydrogen generated to produce cellulose and other organic chemicals; these products, in the form of wood, may be burned in a fireplace, producing CO2 and water, releasing energy in the form of heat and light to their surroundings, e.g. to a home's interior and chimney gasses. Exothermic refers to a transformation in which a system releases energy to the surroundings, expressed by Q < 0. When the transformation occurs at constant pressure, one has for the enthalpy ∆H < 0,and constant volume, one has for the internal energy ∆U < 0.
In an adiabatic system, an exothermic process results in an increase in temperature of the system. In exothermic chemical reactions, the heat, released by the reaction takes the form of electromagnetic energy; the transition of electrons from one quantum energy level to another causes light to be released. This light is equivalent in energy to the stabilization energy of the energy for the chemical reaction, i.e. the bond energy. This light, released can be absorbed by other molecules in solution to give rise to molecular vibrations or rotytions, which gives rise to the classical understanding of heat. In contrast, when endothermic reactions occur, energy is absorbed to place an electron in a higher energy state, such that the electron can associate with another atom to form a chemical complex. Net energy is absorbed by an endothermic reaction. In an exothermic reaction, the energy needed to start the reaction is less than el energy, subsequently released, so there is a net release of energy.
This is the physical understanding of endothermic reactions. Some examples of exothermic processes are: Combustion of fuels such as wood and oil petroleum Thermite reaction Reaction of alkali metals and other electropositive metals with water Condensation of rain from water vapor Mixing water and strong acids or strong bases Mixing acids and bases Dehydration of carbohydrates by sulfuric acid The setting of cement and concrete Some polymerisation reactions such as the setting of epoxy resin Reaction of most metals with halogens or oxygen Nuclear fusion in hydrogen bombs and in stellar cores Nuclear fission of heavy elements Chemical exothermic reactions are more spontaneous than their counterparts, endothermic reactions. In a thermochemical reaction, exothermic, the heat may be listed among the products of the reaction; because of historical accident, students encounter a source of possible confusion between the terminology of physics and biology. Whereas the thermodynamic terms "exothermic" and "endothermic" refer to processes that give out heat energy and processes that absorb heat energy, in biology the sense is inverted.
The metabolic terms "ectothermic" and "endothermic" refer to organisms that rely on external heat to achieve a full working temperature, to organisms that produce heat from within as a major factor in controlling their bodily temperature. Http://chemistry.about.com/b/a/184556.htm Observe exothermic reactions in a simple experiment
Catalysis is the process of increasing the rate of a chemical reaction by adding a substance known as a catalyst, not consumed in the catalyzed reaction and can continue to act repeatedly. Because of this, only small amounts of catalyst are required to alter the reaction rate in principle. In general, chemical reactions occur faster in the presence of a catalyst because the catalyst provides an alternative reaction pathway with a lower activation energy than the non-catalyzed mechanism. In catalyzed mechanisms, the catalyst reacts to form a temporary intermediate, which regenerates the original catalyst in a cyclic process. A substance which provides a mechanism with a higher activation energy does not decrease the rate because the reaction can still occur by the non-catalyzed route. An added substance which does reduce the reaction rate is not considered a catalyst but a reaction inhibitor. Catalysts may be classified as either heterogeneous. A homogeneous catalyst is one whose molecules are dispersed in the same phase as the reactant's molecules.
A heterogeneous catalyst is one whose molecules are not in the same phase as the reactant's, which are gases or liquids that are adsorbed onto the surface of the solid catalyst. Enzymes and other biocatalysts are considered as a third category. In the presence of a catalyst, less free energy is required to reach the transition state, but the total free energy from reactants to products does not change. A catalyst may participate in multiple chemical transformations; the effect of a catalyst may vary due to the presence of other substances known as inhibitors or poisons or promoters. Catalyzed reactions have a lower activation energy than the corresponding uncatalyzed reaction, resulting in a higher reaction rate at the same temperature and for the same reactant concentrations. However, the detailed mechanics of catalysis is complex. Catalysts may bind to the reagents to polarize bonds, e.g. acid catalysts for reactions of carbonyl compounds, or form specific intermediates that are not produced such as osmate esters in osmium tetroxide-catalyzed dihydroxylation of alkenes, or cause dissociation of reagents to reactive forms, such as chemisorbed hydrogen in catalytic hydrogenation.
Kinetically, catalytic reactions are typical chemical reactions. The catalyst participates in this slowest step, rates are limited by amount of catalyst and its "activity". In heterogeneous catalysis, the diffusion of reagents to the surface and diffusion of products from the surface can be rate determining. A nanomaterial-based catalyst is an example of a heterogeneous catalyst. Analogous events associated with substrate binding and product dissociation apply to homogeneous catalysts. Although catalysts are not consumed by the reaction itself, they may be inhibited, deactivated, or destroyed by secondary processes. In heterogeneous catalysis, typical secondary processes include coking where the catalyst becomes covered by polymeric side products. Additionally, heterogeneous catalysts can dissolve into the solution in a solid–liquid system or sublimate in a solid–gas system; the production of most industrially important chemicals involves catalysis. Most biochemically significant processes are catalysed.
Research into catalysis is a major field in applied science and involves many areas of chemistry, notably organometallic chemistry and materials science. Catalysis is relevant to many aspects of environmental science, e.g. the catalytic converter in automobiles and the dynamics of the ozone hole. Catalytic reactions are preferred in environmentally friendly green chemistry due to the reduced amount of waste generated, as opposed to stoichiometric reactions in which all reactants are consumed and more side products are formed. Many transition metals and transition metal complexes are used in catalysis as well. Catalysts called. A catalyst works by providing an alternative reaction pathway to the reaction product; the rate of the reaction is increased as this alternative route has a lower activation energy than the reaction route not mediated by the catalyst. The disproportionation of hydrogen peroxide creates oxygen, as shown below. 2 H2O2 → 2 H2O + O2This reaction is preferable in the sense that the reaction products are more stable than the starting material, though the uncatalysed reaction is slow.
In fact, the decomposition of hydrogen peroxide is so slow that hydrogen peroxide solutions are commercially available. This reaction is affected by catalysts such as manganese dioxide, or the enzyme peroxidase in organisms. Upon the addition of a small amount of manganese dioxide, the hydrogen peroxide reacts rapidly; this effect is seen by the effervescence of oxygen. The manganese dioxide is not consumed in the reaction, thus may be recovered unchanged, re-used indefinitely. Accordingly, manganese dioxide catalyses this reaction. Catalytic activity is denoted by the symbol z and measured in mol/s, a unit, called katal and defined the SI unit for catalytic activity since 1999. Catalytic activity is not a kind of reaction rate, but a property of the catalyst under certain conditions, in relation to a specific chemical reaction. Catalytic activity of one katal of a catalyst means one mole of that catalyst will catalyse 1 mole of the reactant to product in one second. A catalyst may and will have different catalytic activity for di
Hydrogen is a chemical element with symbol H and atomic number 1. With a standard atomic weight of 1.008, hydrogen is the lightest element in the periodic table. Hydrogen is the most abundant chemical substance in the Universe, constituting 75% of all baryonic mass. Non-remnant stars are composed of hydrogen in the plasma state; the most common isotope of hydrogen, termed protium, has no neutrons. The universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, tasteless, non-toxic, nonmetallic combustible diatomic gas with the molecular formula H2. Since hydrogen forms covalent compounds with most nonmetallic elements, most of the hydrogen on Earth exists in molecular forms such as water or organic compounds. Hydrogen plays a important role in acid–base reactions because most acid-base reactions involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a negative charge when it is known as a hydride, or as a positively charged species denoted by the symbol H+.
The hydrogen cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds are always more complex. As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics. Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, that it produces water when burned, the property for which it was named: in Greek, hydrogen means "water-former". Industrial production is from steam reforming natural gas, less from more energy-intensive methods such as the electrolysis of water. Most hydrogen is used near the site of its production, the two largest uses being fossil fuel processing and ammonia production for the fertilizer market. Hydrogen is a concern in metallurgy as it can embrittle many metals, complicating the design of pipelines and storage tanks.
Hydrogen gas is flammable and will burn in air at a wide range of concentrations between 4% and 75% by volume. The enthalpy of combustion is −286 kJ/mol: 2 H2 + O2 → 2 H2O + 572 kJ Hydrogen gas forms explosive mixtures with air in concentrations from 4–74% and with chlorine at 5–95%; the explosive reactions may be triggered by heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C. Pure hydrogen-oxygen flames emit ultraviolet light and with high oxygen mix are nearly invisible to the naked eye, as illustrated by the faint plume of the Space Shuttle Main Engine, compared to the visible plume of a Space Shuttle Solid Rocket Booster, which uses an ammonium perchlorate composite; the detection of a burning hydrogen leak may require a flame detector. Hydrogen flames in other conditions are blue; the destruction of the Hindenburg airship was a notorious example of hydrogen combustion and the cause is still debated. The visible orange flames in that incident were the result of a rich mixture of hydrogen to oxygen combined with carbon compounds from the airship skin.
H2 reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding hydrogen halides, hydrogen chloride and hydrogen fluoride, which are potentially dangerous acids; the ground state energy level of the electron in a hydrogen atom is −13.6 eV, equivalent to an ultraviolet photon of 91 nm wavelength. The energy levels of hydrogen can be calculated accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the Sun. However, the atomic electron and proton are held together by electromagnetic force, while planets and celestial objects are held by gravity; because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, therefore only certain allowed energies. A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation, Dirac equation or the Feynman path integral formulation to calculate the probability density of the electron around the proton.
The most complicated treatments allow for the small effects of special relativity and vacuum polarization. In the quantum mechanical treatment, the electron in a ground state hydrogen atom has no angular momentum at all—illustrating how the "planetary orbit" differs from electron motion. There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei. In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state with a molecular spin quantum number of 1. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form known as the "normal form"; the equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but because the ortho form is an excited state and has a higher energy
Titanium is a chemical element with symbol Ti and atomic number 22. It is a lustrous transition metal with a silver color, low density, high strength. Titanium is resistant to corrosion in sea water, aqua regia, chlorine. Titanium was discovered in Cornwall, Great Britain, by William Gregor in 1791, was named by Martin Heinrich Klaproth after the Titans of Greek mythology; the element occurs within a number of mineral deposits, principally rutile and ilmenite, which are distributed in the Earth's crust and lithosphere, it is found in all living things, water bodies and soils. The metal is extracted from its principal mineral ores by the Hunter processes; the most common compound, titanium dioxide, is a popular photocatalyst and is used in the manufacture of white pigments. Other compounds include a component of smoke screens and catalysts. Titanium can be alloyed with iron, aluminium and molybdenum, among other elements, to produce strong, lightweight alloys for aerospace, industrial processes, agri-food, medical prostheses, orthopedic implants and endodontic instruments and files, dental implants, sporting goods, mobile phones, other applications.
The two most useful properties of the metal are corrosion resistance and strength-to-density ratio, the highest of any metallic element. In its unalloyed condition, titanium is less dense. There are two allotropic forms and five occurring isotopes of this element, 46Ti through 50Ti, with 48Ti being the most abundant. Although they have the same number of valence electrons and are in the same group in the periodic table and zirconium differ in many chemical and physical properties; as a metal, titanium is recognized for its high strength-to-weight ratio. It is a strong metal with low density, quite ductile and metallic-white in color; the high melting point makes it useful as a refractory metal. It is paramagnetic and has low electrical and thermal conductivity. Commercially pure grades of titanium have ultimate tensile strength of about 434 MPa, equal to that of common, low-grade steel alloys, but are less dense. Titanium is 60% denser than aluminium, but more than twice as strong as the most used 6061-T6 aluminium alloy.
Certain titanium alloys achieve tensile strengths of over 1,400 MPa. However, titanium loses strength when heated above 430 °C. Titanium is not as hard as some grades of heat-treated steel. Machining requires precautions, because the material can gall unless sharp tools and proper cooling methods are used. Like steel structures, those made from titanium have a fatigue limit that guarantees longevity in some applications; the metal is a dimorphic allotrope of an hexagonal α form that changes into a body-centered cubic β form at 882 °C. The specific heat of the α form increases as it is heated to this transition temperature but falls and remains constant for the β form regardless of temperature. Like aluminium and magnesium, titanium metal and its alloys oxidize upon exposure to air. Titanium reacts with oxygen at 1,200 °C in air, at 610 °C in pure oxygen, forming titanium dioxide, it is, slow to react with water and air at ambient temperatures because it forms a passive oxide coating that protects the bulk metal from further oxidation.
When it first forms, this protective layer continues to grow slowly. Atmospheric passivation gives titanium excellent resistance to corrosion equivalent to platinum. Titanium is capable of withstanding attack by dilute sulfuric and hydrochloric acids, chloride solutions, most organic acids. However, titanium is corroded by concentrated acids; as indicated by its negative redox potential, titanium is thermodynamically a reactive metal that burns in normal atmosphere at lower temperatures than the melting point. Melting is possible only in a vacuum. At 550 °C, it combines with chlorine, it reacts with the other halogens and absorbs hydrogen. Titanium is one of the few elements that burns in pure nitrogen gas, reacting at 800 °C to form titanium nitride, which causes embrittlement; because of its high reactivity with oxygen and some other gases, titanium filaments are applied in titanium sublimation pumps as scavengers for these gases. Such pumps inexpensively and reliably produce low pressures in ultra-high vacuum systems.
Titanium is the ninth-most abundant element in the seventh-most abundant metal. It is present as oxides in most igneous rocks, in sediments derived from them, in living things, natural bodies of water. Of the 801 types of igneous rocks analyzed by the United States Geological Survey, 784 contained titanium, its proportion in soils is 0.5 to 1.5%. Common titanium-containing minerals are anatase, ilmenite, perovskite and titanite. Akaogiite is an rare mineral consisting of titanium dioxide. Of these minerals, only rutile and ilmenite have economic importance, yet they are difficult to find in high concentrations. About 6.0 and 0.7 million tonnes of those minerals were mined in 2011, respectively. Signi