Hydrogen
The hydrogen (in Greek, from ὕδωρ hýdōr, genitive ὑδρός hydrós, and γένος génos «which generates or produces water") is the chemical element with atomic number 1, represented by the symbol H. With an atomic mass of 1.00797, it is the lightest element on the periodic table of elements. It usually occurs in its molecular form, forming the diatomic gas H2 under normal conditions. This gas is flammable, colorless, odorless, non-metallic, and insoluble in water.
Due to its different and varied properties, hydrogen cannot be clearly placed in any group of the periodic table, although it is often placed in group 1 (or family 1A) because it has only one electron in the valence shell or top layer.
Hydrogen is the most abundant chemical element, constituting approximately 75% of the visible matter in the universe. In their main sequence, stars are composed primarily of hydrogen in the plasma state. Elemental hydrogen is relatively rare on Earth and is produced industrially from hydrocarbons such as methane. Most elemental hydrogen is obtained in situ, that is, where and when it is needed. The largest markets in the world enjoy the use of hydrogen for the upgrading of fossil fuels (in the hydrocracking process) and in the production of ammonia (mainly for the fertilizer market). Hydrogen can be obtained from water by an electrolysis process, but it is a much more expensive method than obtaining it from natural gas.
The most common isotope of hydrogen is protium, whose nucleus consists of a single proton and no neutrons. In ionic compounds, it can have a positive charge (becoming a cation called a hydron, H+, made up of only one proton, sometimes in the presence of 1 or 2 neutrons); or negatively charged (becoming an anion known as a hydride, H-). Other isotopes can also be formed, such as deuterium, with one neutron, and tritium, with two neutrons. In 2001, the isotope 4H was created in the laboratory and, starting in 2003, isotopes 5H through 7H were synthesized. Hydrogen forms compounds with most elements and is present in water and most organic compounds. It has a particularly important role in acid-base chemistry, in which many reactions involve the exchange of protons (hydrogen ions, H+) between soluble molecules. Since it is the only neutral atom for which the Schrödinger equation can be solved analytically, the study of the energy and bonding of the hydrogen atom has been fundamental to the point of having played a major role in the development of quantum mechanics..
The characteristics of this element and its solubility in various metals are very important in metallurgy, since many metals can become brittle in its presence, and in the development of safe ways to store it for use as fuel. highly soluble in various rare earth and transition metal compounds, and can be dissolved in both crystalline and amorphous metals. The solubility of hydrogen in metals is influenced by local distortions or impurities in the metal's crystal structure.
Etymology
The term hydrogen comes from the Latin hydrogenium, and this from the ancient Greek ὕδωρ (hydro): ' water' and γένος-ου(genos): 'generator'; that is, "producer of water." That was the name with which Antoine Lavoisier baptized it. The word can refer to both the hydrogen atom, described in this article, and the diatomic molecule (H2), which is found at trace level in the Earth's atmosphere. Chemists tend to refer to this molecule as dihydrogen, a hydrogen molecule, or diatomic hydrogen, to distinguish it from the atom of the element, which does not exist in isolation under ordinary conditions.
History
Discovery of hydrogen and use
Gas diatomic hydrogen, H2, was the first produced artificially and formally described by T. von Hohenheim (Paracelsus), who obtained it artificially by mixing metals with strong acids. Paracelsus was not aware that the flammable gas generated in these chemical reactions was composed of a new chemical element. In 1671, Robert Boyle rediscovered and described the reaction that occurred between iron filings and dilute acids, resulting in the production of hydrogen gas. In 1766, Henry Cavendish was the first to recognize gaseous hydrogen as a discrete substance, identifying the gas produced in the metal-acid reaction as "flammable air" and discovering further, in 1781, that the gas produces water when burned. He is generally given credit for its discovery as a chemical element. In 1783, Antoine Lavoisier gave the element the name hydrogen (from the Greek υδρώ (hydro), water and γένος-ου (genes) generate, i.e. "water producer") when he and Laplace reproduced Cavendish's discovery that water is produced when hydrogen is burned.
Lavoisier produced hydrogen for his experiments on conservation of mass by reacting a stream of steam with metallic iron through a fire-heated incandescent iron tube. The anaerobic oxidation of iron by the protons of water at high temperature can be represented schematically by the set of the following reactions:
- Fe + H2O → FeO + H2
- 2 Fe + 3 H2O → Fe2O3 + 3 H2
- 3 Fe + 4 H2O → Fe3O4 + 4 H2
Many metals, such as zirconium, undergo a similar reaction with water, leading to the production of hydrogen.
Hydrogen was first liquefied by James Dewar in 1898 using regenerative refrigeration, and his invention closely approximates what we know today as a thermos. He produced solid hydrogen the following year. Deuterium was discovered in December 1931 by Harold Urey, and tritium was prepared in 1934 by Ernest Rutherford, Marcus Oliphant, and Paul Harteck. Heavy water, which has deuterium instead of regular hydrogen in the water molecule, was discovered by the team at Urey in 1932.
François Isaac de Rivaz built the first internal combustion device powered by a mixture of hydrogen and oxygen in 1806. Edward Daniel Clarke invented the hydrogen gas blowout in 1819. The Döbereiner's lamp and the Drummond Luminaire were invented in 1823.
The filling of the first balloon with hydrogen gas was documented by Jacques Charles in 1783. Hydrogen provided ascent to the first reliable form of air travel after the invention of the first retired hydrogen airship in 1852 by Henri Giffard. German Count Ferdinand von Zeppelin promoted the idea of using hydrogen in rigid airships, later called zeppelins, the first of which had its maiden flight in 1900. Normal flight began in 1910, and by the start of the World War I, by August 1914, 35,000 passengers had been carried without any serious incident. Hydrogen-lifted airships are used as observation platforms and bombers during warfare.
The first non-stop transatlantic crossing was made by the British airship R34 in 1919. Starting in 1928, with the Graf Zeppelin LZ 127, regular passenger service continued until the middle of the decade of 1930 without any incident. With the discovery of reserves of another type of light gas in the United States, this project had to be modified, since the other element promised more security, but the United States Government refused to sell the gas for that purpose. Therefore, the H2 was used on the airship Hindenburg, which was destroyed in an in-flight incident over New Jersey on May 6, 1937. The incident was broadcast live on radio and filmed. The ignition of a hydrogen leak was attributed as the cause of the incident, but subsequent investigations pointed to ignition of the aluminized fabric lining by static electricity.
Role of hydrogen in quantum theory
Thanks to its relatively simple atomic structure, consisting of a single proton and a single electron for the most abundant isotope (protium), the hydrogen atom has an absorption spectrum that could be explained quantitatively, which was the central point Bohr's atomic model, which was a milestone in the development of the theory of atomic structure. In addition, the consequent simplicity of the diatomic hydrogen molecule and the corresponding dihydrogen cation, H2+, allowed a more complete understanding of the nature of chemical bonding, which was soon followed. then with the quantum mechanical treatment of the hydrogen atom, which had been developed in the mid-1920s by Erwin Schrödinger and Werner Heisenberg.
One of the first quantum effects that was explicitly noticed (but not understood at the time) was an observation by Maxwell involving hydrogen, half a century before quantum mechanical theory was fully established. Maxwell observed that the specific heat of H2 inexplicably deviated from that of a diatomic gas below room temperature and began to look more and more like that of a monatomic gas at very low temperatures.. According to quantum theory, this behavior results from the spacing of the rotational (quantized) energy levels, which are particularly far apart in H2 due to its small mass. These widely separated levels prevent the equal distribution of heat energy to generate rotational movement in hydrogen at low temperatures. Diatomic gases composed of heavy atoms do not have rotational energy levels as far apart and therefore do not have the same effect as hydrogen.
Abundance in nature
Hydrogen is the most abundant chemical element in the universe, accounting for more than 75% normal matter by mass and more than 90% by number of atoms. This element is found in abundance in stars and gas giant planets. H2 molecular clouds are associated with star formation. Hydrogen also plays a key role as fuel for stars through nuclear fusion reactions between hydrogen nuclei.
In the universe, hydrogen is found primarily in its atomic form and plasma state, whose properties are quite different from molecular hydrogen. As a plasma, the electron and proton of hydrogen are not bound, so it has high electrical conductivity and great emissivity (source of light emitted by the Sun and other stars). Charged particles are strongly influenced by electric and magnetic fields. For example, in solar winds the particles interact with the terrestrial magnetosphere generating Birkeland currents and the aurora phenomenon.
Under normal conditions of pressure and temperature, hydrogen exists as a diatomic gas, H2. However, gaseous hydrogen is extremely low in Earth's atmosphere (1 ppm by volume), due to its small mass which allows it to escape the influence of Earth's gravity more easily than other heavier gases. Although hydrogen atoms and diatomic hydrogen molecules abound in interstellar space, they are difficult to generate, concentrate, and purify on Earth. Hydrogen is the fifteenth most abundant element on the earth's surface. Most of the terrestrial hydrogen is found as part of chemical compounds such as hydrocarbons or water. Gaseous hydrogen is produced by some bacteria and algae, and is a component of flatulence.
Properties
Combustion
Hydrogen gas (dihydrogen) is highly flammable and burns in concentrations of 4% or more H2 in air. The enthalpy of combustion of hydrogen is −285.8 kJ/mol; burns according to the following balanced equation.
- 2 H2(g) + O2(g) → 2 H2O(l) + 572 kJ (285.8 kJ/mol)
When mixed with oxygen in a variety of proportions, hydrogen explodes on ignition. Hydrogen burns violently in air; ignition occurs automatically at a temperature of 560 °C. Pure hydrogen-oxygen flames burn in the ultraviolet color range and are nearly invisible to the naked eye, as evidenced by the dim flame of the shuttle's main engines space (unlike the easily visible flames of the solid's rocket booster). So a flame detector is needed to detect if a hydrogen leak is burning. The Hindenburg airship explosion was an infamous case of hydrogen combustion. The cause was debated, but combustible materials on the aircraft's deck were responsible for the color of the flames. Another characteristic of hydrogen fires is that the flames tend to rise rapidly with the gas in the air, as illustrated by the flames from the Hindenburg, causing less damage than hydrocarbon fires. Two-thirds of the passengers on the Hindenburg survived the fire, and many of the deaths that occurred were from falls or diesel fuel fires.
H2 reacts directly with other oxidizing elements. A spontaneous and violent reaction can occur at room temperature with chlorine and fluorine, forming the corresponding hydrogen halides: hydrogen chloride and hydrogen fluoride.
Unlike hydrocarbons, the combustion of hydrogen does not generate carbon oxides (monoxide and dioxide) but simply water in the form of steam, which is why it is considered an environmentally friendly fuel and helps to mitigate global warming.
Electronic energy levels
The energy level of the electronic ground state of a hydrogen atom is –13.6 eV, which is equivalent to an ultraviolet photon of approximately 92 nm wavelength.
The energy levels of hydrogen can be calculated fairly accurately using the Bohr atomic model, which assumes that the electron orbits the proton in a manner analogous to Earth's orbit around the Sun. However, the electromagnetic force causes the proton to and the electron attract each other, in the same way that planets and other celestial bodies attract each other by gravitational force. Due to the discrete (quantized) character of angular momentum postulated at the beginning of quantum mechanics by Bohr, the electron in Bohr's model can only orbit at certain allowed distances around the proton and, by extension, with certain allowed energy values. A more precise description of the hydrogen atom is given by a purely quantum mechanical treatment that uses the Schrödinger wave equation or the equivalent formulation of Feynman path integrals to calculate the probability density of the electron near the proton. treatment of the electron through the de Broglie hypothesis (wave-particle duality) reproduces chemical results (such as the configuration of the hydrogen atom) in a more natural way than the Bohr particle model, although the energy and spectral results they are the same. If the reduced mass of the nucleus and the electron are used in the construction of the model (as would be done in the two-body problem in Classical Mechanics), a better formulation is obtained for the spectra of hydrogen, and the correct spectral shifts for deuterium. and tritium. Small adjustments in the energy levels of the hydrogen atom, which correspond to real spectral effects, can be determined using the full quantum mechanical theory, which corrects for the effects of special relativity (see Dirac equation), and computing the quantum effects caused by the production of virtual particles in a vacuum and as a result of electric fields (see Quantum Electrodynamics).
In gaseous hydrogen, the energy level of the fundamental electronic state is divided into other levels of hyperfine structure, originated by the effect of the magnetic interactions produced between the spins of the electron and the proton. The energy of the atom when the spins of the proton and electron are aligned is greater than when the spins are not. The transition between those two states can take place by emitting a photon through a magnetic dipole transition. Radio telescopes can detect the radiation produced in this process, which is used to create maps of the distribution of hydrogen in the galaxy.
Elementary molecular forms
There are two distinct types of diatomic hydrogen molecules that differ in the relationship between the spins of their nuclei: While in the orthohydrogen form, the spins of the two protons are parallel and form a triplet state, in the form of para-hydrogen, the spins are antiparallel and form a singular. Under normal conditions of pressure and temperature, gaseous hydrogen contains approximately 25% of the para form and 75% of the ortho form, also known as " normal form". The equilibrium relationship between orthohydrogen and para-hydrogen depends on temperature, but since the ortho form is an excited state and therefore has a higher energy, it is also unstable and cannot be purified. At very low temperatures, the equilibrium state consists almost exclusively of the para form. The physical properties of pure para-hydrogen differ slightly from those of the normal (ortho) form. The distinction between ortho/para forms also it occurs in other molecules or functional groups that contain hydrogen, such as water or methylene.
The uncatalyzed interconversion between para-hydrogen and ortho-hydrogen increases with increasing temperature; for this reason, rapidly condensed H2 contains large amounts of the ortho form which transitions to the para form slowly. >ortho/para in the H2 condensate is something important to consider for the preparation and storage of liquid hydrogen: the conversion of the form ortho to the para form is exothermic and produces enough heat to evaporate the liquid hydrogen, causing the loss of liquefied material. Catalysts for ortho/para interconversion, such as iron compounds, are used in hydrogen refrigeration processes.
A molecular form called protonated molecular hydrogen, H3+, is found in the interstellar medium, where it is generated by the ionization of molecular hydrogen caused by cosmic rays. It has also been observed in the upper layers of Jupiter's atmosphere. This molecule is relatively stable in the middle of outer space due to the low temperatures and very low density. H3+ is one of the most abundant ions in the universe, and plays a notable role in the chemistry of the interstellar medium.
Metallic hydrogen
Although hydrogen is often classified as a nonmetal, at high temperatures and pressures it can behave like a metal. In March 1996, a group of scientists at the Lawrence Livermore National Laboratory reported that they had casually produced, within a microsecond and at temperatures of thousands of kelvins and pressures of more than a million atmospheres (>100 GPa), the first metallic hydrogen. identifiable.
Compounds
Covalent and organic compounds
Although H2 is not very reactive under normal conditions, it forms a multitude of compounds with most chemical elements. Millions of hydrocarbons are known, but they are not generated by the direct reaction of elemental hydrogen with carbon (although syngas production followed by the Fischer-Tropsch process to synthesize hydrocarbons seems to be an exception since it starts with carbon and elemental hydrogen generated internally). if you). Hydrogen can form compounds with more electronegative elements, such as halogens (fluorine, chlorine, bromine, iodine) or chalcogens (oxygen, sulfur, selenium); in these compounds, hydrogen acquires a partial positive charge due to the polarity of the covalent bond. When bound to fluorine, oxygen, or nitrogen, hydrogen can participate in a non-covalent form of bonding called "hydrogen bonding" or "hydrogen bonding", which is essential for the stability of many biological molecules. Hydrogen can also form compounds with less electronegative elements, such as metals or semimetals, in which it acquires a partial negative charge. These compounds are known as hydrides.
Hydrogen forms a huge variety of compounds with carbon. Due to their presence in living things, these compounds are called organic compounds; the study of its properties is the purpose of Organic Chemistry, and the study in the context of living organisms is known as Biochemistry. According to some definitions, "organic" they require the presence of carbon to be called that (there we have the classic example of urea) but not all carbon compounds are considered organic (this is the case of carbon monoxide, or metallic carbonates. Most organic compounds also they contain hydrogen and, since it is the carbon-hydrogen bond that gives these compounds many of their main characteristics, it is necessary to mention the carbon-hydrogen bond in some definitions of the word "organic" in Chemistry. (These recent definitions are not perfect, however, since an undoubtedly organic compound like urea could not be classified as such based on them.)
In inorganic chemistry, hydrides can also serve as bridging ligands that join two metal centers in a coordination complex. This function is particularly common in group 13 elements, especially boranes (boron hydrides) and aluminum complexes, as well as carborane clusters.
Some examples of important covalent compounds containing hydrogen are: ammonia (NH3), hydrazine (N2H4), water (H2O), hydrogen peroxide (H2O2), hydrogen sulfide (H2 S), etc.
Hydrides
Compounds of hydrogen are often called hydrides, a term used rather inaccurately. For chemists, the term "hydride" it generally implies that the hydrogen atom has acquired a partial negative charge or anionic character (denoted as H-). The existence of the hydride anion, proposed by G. N. Lewis in 1916 for group 1 (I) and 2 (II) ionic hydrides, was demonstrated by Moers in 1920 with the electrolysis of molten lithium hydride (LiH), which produced a quantity stoichiometric hydrogen at the anode. For metal hydrides of other groups, the term is quite misleading, considering the low electronegativity of hydrogen. An exception in group II hydrides is BeH2, which is polymeric. In lithium tetrahydride aluminate (III), the AlH4- anion has its hydric centers firmly bound to aluminum (III).
Although hydrides can be formed with almost all main group elements, the number and combination of possible compounds varies greatly; for example, there are more than 100 known boron binary hydrides, but only one of aluminum. The indium binary hydride has not yet been identified, although larger complexes exist.
“Protons” and acids
The oxidation of H2 formally gives rise to the proton, H+. This species is central to explaining the properties of acids, although the term "proton" is used imprecisely to refer to the cationic hydrogen or hydrogen ion, denoted H+. An isolated H+ proton cannot exist in solution because of its strong tendency to bind to atoms or molecules with electrons through a coordinate bond or dative bond. To avoid the convenient, though uncertain, idea of the isolated solvated proton in solution, in aqueous acid solutions the presence of the hydronium ion (H3O+) organized in clusters to form the species H9O4+. Other oxonium ions are present when water forms solutions with other solvents.
Although exotic on Earth, one of the most common ions in the universe is H3+, known as protonated molecular hydrogen or triatomic hydrogen cation.
Isotopes
The most common isotope of hydrogen does not have neutrons, there are two others, deuterium (D) with one and tritium (T), radioactive with two. Deuterium has a natural abundance between 0.0184 and 0.0082% (International Union of Pure and Applied Chemistry (IUPAC)). Hydrogen is the only chemical element that has different names and chemical symbols for its different isotopes.
Hydrogen also has other highly unstable isotopes (from 4H to 7H), which were synthesized in the laboratory, but never observed in nature.
- 1H, known as protio, is the most common isotope of hydrogen with an abundance of more than 99.98 %. Because the core of this isotope is formed by a single proton has been baptized as a protio, a name that despite being very descriptive, is unused.
- 2H, the other stable isotope of hydrogen, is known as deuterium and its core contains a proton and neutron. Deuterium represents 0.0026 % or 0.0184 % (according to molar fraction or atomic fraction) of the hydrogen present on Earth, with the lowest concentrations in gaseous hydrogen, and the largest (0.015 % or 150 ppm) in oceanic waters. Deuterium is not radioactive, and does not pose a significant risk of toxicity. Water enriched in molecules that include deuterium instead of hydrogen 1H (protio), is called heavy water. Deuterium and its compounds are used in non-radioactive marking in experiments and also in solvents used in spectroscopy 1H - RMN. Heavy water is used as a neutron and refrigerant moderator in nuclear reactors. Deuterium is also a fuel potential for commercial nuclear fusion.
- 3H is known as tritio and contains a proton and two neutrons in its core. It's radioactive, disintegrating in 32He+ through a beta emission. It has a period of semi-disintegration of 12,33 years. Small amounts of tritio are found in nature by effect of the interaction of cosmic rays with atmospheric gases. It has also been released tritio for testing nuclear weapons. The tritio is used in nuclear fusion reactions, such as a tracer in Isotopian geochemistry, and in self-feeding luminous devices. It was first common to use tritium as a radiomarker in chemical and biological experiments, but it is currently used less.
Hydrogen is the only element that has different common names for each of its (natural) isotopes. During the early days of radioactivity studies, some heavy radioactive isotopes were given names, but none are still in use. The symbols D and T (instead of ²H and ³H) are sometimes used to refer to deuterium and tritium, but the symbol P corresponds to phosphorus and therefore cannot be used to represent protium. The IUPAC declares that although the use of these symbols is common, it is not recommended.
Biological reactions
H2 is a product of some types of anaerobic metabolism and is produced by various microorganisms, usually through reactions catalyzed by iron- or nickel-containing enzymes called hydrogenases. These enzymes catalyze the reversible redox reaction between H2 and its components, two protons and two electrons. The creation of hydrogen gas occurs in the transfer of reduced equivalents produced during the fermentation of pyruvate to water.
Water splitting, in which water is broken down into its components, protons, electrons, and oxygen, occurs during the clear phase in all photosynthetic organisms. Some organisms—including the alga Chlamydomonas reinhardtii and cyanobacteria—evolved a further step in the dark phase in which protons and electrons are reduced to form H2 gas by specialized hydrogenases in the chloroplast. Efforts were made to genetically modify cyanobacterial hydrogenases to efficiently synthesize H2 gas even in the presence of oxygen. Efforts were also made with genetically modified algae in a bioreactor.
Production
H2 gas is produced in chemistry and biology laboratories, often as a by-product of the dehydrogenation of unsaturated substrates; and in nature as a means of expelling reducing equivalents in biochemical reactions.
Laboratory
In the laboratory, H2 gas is normally prepared by the reaction of acids with metals such as zinc, using the Kipp apparatus.
Zn + 2 H+ → Zn2+ + H2
Aluminum can also produce H2 after base treatment:
2 Al + 6 H2O + 2 OH- → 2 Al(OH)4- + 3 H2
Water electrolysis is a simple method of producing hydrogen. A low voltage electric current flows through the water, and oxygen gas is formed at the anode, while hydrogen gas is formed at the cathode. Typically, the cathode is made of platinum or another inert metal (usually platinum or graphite), when hydrogen is produced for storage. If, however, the gas is intended to be burned on site, it is desirable that there be oxygen to assist combustion, and then both electrodes may be made of inert metals (iron electrodes should be avoided as they consume oxygen by undergoing oxidation). Theoretical maximum efficiency (electricity used vs energy value of hydrogen produced) is between 80 and 94%.
2H2O(aq) → 2H2(g) + O2(g)
In 2007, it was discovered that a pelletized aluminum gallium alloy added to water could be used to generate hydrogen. The process also produces alumina, but gallium can be reused, which prevents the formation of an oxide film on the granules. This has major implications for the potential hydrogen-based economy, as it can be produced on site and does not have to be transported.
Industrial
Hydrogen can be prepared through various processes, but today the most important is the extraction of hydrogen from hydrocarbons. Most commercial hydrogen is produced by catalytic reforming of natural gas or liquid hydrocarbons. At high temperatures (700-1100 °C), water vapor is reacted with methane to produce carbon monoxide and H2:
- CH4 + H2O → CO + 3 H2
This reaction is thermodynamically favored by excess steam and low pressures but is normally practiced at high pressures (20 atm) for economic reasons. The mixture produced is known as "synthesis gas," as it is often used directly for the synthesis of methanol and other chemicals. Other hydrocarbons, in addition to methane, can be used to produce syngas with varying proportions of the products.
If the desired product is just hydrogen, carbon monoxide is reacted through the water vapor shift reaction, for example with an iron oxide catalyst. This reaction is also a common industrial source of carbon dioxide:
- CO + H2O → CO2 + H2
Other options for producing hydrogen from methane are pyrolysis, which results in the formation of solid carbon:
- CH4 → C + 2 H2
Or partial oxidation, which also applies to fuels like coal:
- 2 CH4 + O2 → 2 CO + 4 H2
Another process that produces hydrogen as a by-product is the electrolysis of brine to produce chlorine.
Solar thermochemicals
There are more than 200 thermochemical cycles that can be used for the separation of water, about a dozen of these cycles, such as the iron oxide cycle, the cerium(III) oxide-cerium(IV) oxide cycle, zinc-zinc oxide cycle, sulfur-iodine cycle, copper-chlorine cycle, hybrid sulfur cycle are under investigation and in the testing phase to produce hydrogen and oxygen from water and heat without using electricity. A number of Laboratories (including France, Germany, Greece, Japan, and the United States) are developing thermochemical methods to produce hydrogen from solar energy and water.
Anaerobic corrosion
Under anaerobic conditions, iron and steel alloys are slowly oxidized by the concomitant water protons reduced to molecular hydrogen (H2). Anaerobic corrosion of iron leads first to the formation of ferrous hydroxide (green rust) and can be described by the following reaction:
- Fe + 2 H2O → Fe(OH)2 + H2
In turn, under anaerobic conditions, ferrous hydroxide (Fe(OH)2) can be oxidized by water protons to form magnetite and molecular hydrogen. This process is described by the Schikorr reaction:
- 3 Fe(OH)2 → Fe3O4 + 2 H2O + H2
- ferrous hydroxide → magnetite + water + hydrogen
The magnetite thus crystallized (Fe3O4) is thermodynamically more stable than ferrous hydroxide (Fe(OH)2).
This process occurs during anaerobic corrosion of iron and steel in oxygen-depleted groundwater and reduced soils below the water table.
Geologic Occurrence: The Serpentinization Reaction
In the absence of atmospheric oxygen (O2), in deep geological conditions that prevail far from the Earth's atmosphere, hydrogen (H2) is produced during the process of serpentinization by the anaerobic oxidation of water protons (H+) of the ferrous silicate (Fe2+) present in the crystalline lattice of fayalite (Fe 2SiO4, the iron olivine). The corresponding reaction leading to the formation of magnetite (Fe3O4), quartz SiO2) and hydrogen (H2 ) is as follows:
- 3 Fe2Yes4 + 2 H2O → 2 Fe3O4 + 3 SiO2 + 3 H2
- fayalita + water → magnetite + quartz + hydrogen
This reaction closely resembles the Schikorr reaction observed in the anaerobic oxidation of ferrous hydroxide in contact with water.
Transformer training
Of all the fault gases formed in electrical transformers, hydrogen is the most common and is generated under most fault conditions, therefore, the formation of hydrogen is a first indication of serious problems in the cycle. life of the transformer.
Applications
Petrochemical Industry
Large amounts of H2 are needed in the petroleum and chemical industries. An additional application of H2 is in the treatment ("upgrading") of fossil fuels, and in the production of ammonia. The main consumers of H2 in a petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking. H2 is used as a hydrogenating agent, particularly in increasing the saturation level of unsaturated fats and oils (found in items such as margarine) and in the production of methanol. Similarly it is the source of hydrogen in the manufacture of hydrochloric acid. H2 is also used as a reducing agent for metallic minerals.
In addition to its use as a reagent, H2 has wide applications in physics and engineering. It is used as a shielding gas in welding methods such as atomic hydrogen welding. H2 is used as a generator cooler in power plants, because it has the highest thermal conductivity of all gases. Liquid H2 is used in cryogenic research, including superconductivity studies. Since H2 is lighter than air, having a little more than 1/15 of air density, it was widely used in the past as a lift gas in hot air balloons and dirigibles.
In newer applications, pure hydrogen or hydrogen mixed with nitrogen (sometimes called forming gas) is used as a tracer gas to detect leaks. Applications can be found in the automotive, chemical, power generation, aerospace, and telecommunications industries. Hydrogen is an authorized food additive (E 949) that allows for package leak testing, among other antioxidant properties.
The rarer isotopes of hydrogen also have specific applications for each one. Deuterium (hydrogen-2) is used in nuclear fission applications as a moderator for slow neutrons, and in nuclear fusion reactions. Deuterium compounds have applications in chemistry and biology in studies of isotope effects. Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs, as an isotopic marker in the biological sciences, as a source of radiation in glow paints.
The equilibrium temperature of the triple point of hydrogen is a fixed point defined on the ITS-90 temperature scale at 13.8033 Kelvin.
Energy Carrier
Hydrogen is not a source of energy, except in the hypothetical context of commercial nuclear fusion power plants using deuterium or tritium, a technology currently far from developed. The sun's energy comes from the nuclear fusion of hydrogen, but this process is difficult to achieve in a controllable way on Earth. Elemental hydrogen from solar, biological, or electrical sources requires more energy to create than is obtained by burning it, so hydrogen will do the trick in these cases. as a carrier of energy, like a battery. It can be obtained from fossil sources (such as methane), but these sources are unsustainable.
The energy density per unit volume of both liquid hydrogen and compressed hydrogen gas at any possible pressure is significantly less than that of traditional fuel sources, although the energy density per unit mass of fuel is higher. However, elemental hydrogen has been widely discussed in the context of energy, as a possible carrier of future large-scale energy in the economy. For example, CO2 sequestration followed by carbon capture and storage could take place at the point of production of H2 from fossil fuels. Hydrogen used in transportation would burn relatively clean, with some CO2 emissions. NOx, but no carbon emissions. However, the infrastructure costs associated with full conversion to a hydrogen economy could be substantial.
Semiconductor Industry
Hydrogen is used to saturate broken bonds of amorphous silicon and amorphous carbon which helps stabilize material properties. It is also a potential electron donor in different oxide materials, including ZnO, SnO2, CdO, MgO, ZrO2, HfO2, La2O3, Y2O3, TiO2, SrTiO3, LaAlO3, SiO2, Al2O3, ZrSiO4, HfSiO4, and SrZrO3.
Safety and precautions
Hydrogen generates various risks to human safety, potential detonations and fires when mixed with air as it is an asphyxiant in its pure, oxygen-free form. In addition, liquid hydrogen is cryogenic and presents dangers (such as frostbite) associated with very cold liquids. The element dissolves in some metals and, in addition to leakage, can have adverse effects on them, such as hydrogen embrittlement. Leakage of hydrogen gas into external air can ignite spontaneously. On the other hand, hydrogen fire, being extremely hot, is almost invisible, and therefore can lead to accidental burns.
Although even interpreting hydrogen data (including data for safety) is confounded by various phenomena. Many of the physical and chemical properties of hydrogen are dependent on the parahydrogen/orthohydrogen ratio (usually taking days or weeks at a given temperature to reach the equilibrium rate for which hydrogen detonation parameters results typically appear, like the critical casting pressure and temperature, are highly dependent on the geometry of the vessel.
Note
- ↑ However, most of the mass of the universe is not in the form of bars or chemical elements. See dark matter and dark energy.
Additional bibliography
- Cotton, F.A.; Wilkinson, G. Advanced Inorganic Chemistry: a comprehensive text, fourth edition. John Wiley & Sons. 1980. ISBN 0-471-02775-8
- Greenwood, N.N.; Earnshaw, A. Chemistry of the Elements, second edition. Butterworth - Heinemann. 1997 ISBN 0-7506-3365-4
- Gutierrez Rios, E. Inorganic chemistry. I reversed. 1994. ISBN 84-291-7215-7
- Shriver, D. F.; Atkins, P. W.; Langford, C. H. Inorganic chemistryVol. 1 Second edition. I reversed. 1997 ISBN 84-291-7004-9
- Quadular interactions in SrZrO3