Hydrogen

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Hydrogen,  ₁H

Purple glow in its plasma state
Spectral lines of hydrogen
General properties
Name, symbol	hydrogen, H
Pronunciation	/ˈhaɪdrədʒən/[1] HY-drə-jən
Appearance	colorless gas
Hydrogen in the periodic table
– ↑ H ↓ Li – ← hydrogen → helium
Atomic number	1
Standard atomic weight	1.008[2] (1.00784–1.00811)[3]
Element category	diatomic nonmetal, could be considered metalloid
Group, block	group 1, s-block
Period	period 1
Electron configuration	1s1
per shell	1
Physical properties
Color	colorless
Phase	gas
Melting point	13.99 K ​(−259.16 °C, ​−434.49 °F)
Boiling point	20.271 K ​(−252.879 °C, ​−423.182 °F)
Density at stp (0 °C and 101.325 kPa)	0.08988 g·L−1
when liquid, at m.p.	0.07 g·cm−3 (solid: 0.0763 g·cm−3)[4]
when liquid, at b.p.	0.07099 g·cm−3
Triple point	13.8033 K, ​7.041 kPa
Critical point	32.938 K, 1.2858 MPa
Heat of fusion	(H2) 0.117 kJ·mol−1
Heat of vaporization	(H2) 0.904 kJ·mol−1
Molar heat capacity	(H2) 28.836 J·mol−1·K−1
vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 15 20
Atomic properties
Oxidation states	1, −1 ​(an amphoteric oxide)
Electronegativity	Pauling scale: 2.20
Ionization energies	1st: 1312.0 kJ·mol−1
Covalent radius	31±5 pm
Van der Waals radius	120 pm
Miscellanea
Crystal structure	​hexagonal
Speed of sound	1310 m·s−1 (gas, 27 °C)
Thermal conductivity	0.1805 W·m−1·K−1
Magnetic ordering	diamagnetic[5]
CAS Registry Number	1333-74-0
History
Discovery	Henry Cavendish[6][7] (1766)
Named by	Antoine Lavoisier[8] (1783)
Most stable isotopes
Main article: Isotopes of hydrogen
iso NA half-life DM DE (MeV) DP 1H 99.9885% 1H is stable with 0 neutrons 2H 0.0115% 2H is stable with 1 neutron 3H trace 12.32 y β− 0.01861 3He

Hydrogen is a chemical element with chemical symbol H and atomic number 1. With an atomic weight of 1.00794 u, hydrogen is the lightest element on the periodic table. Its monatomic form (H) is the most abundant chemical substance in the universe, constituting roughly 75% of all baryonic mass.^{[9][note 1]} Non-remnant stars are mainly composed of hydrogen in its plasma state. The most common isotope of hydrogen, termed protium (name rarely used, symbol ¹H), has a single proton and zero neutrons.

The universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, odorless, tasteless, non-toxic, nonmetallic, highly combustible diatomic gas with the molecular formula H₂. Since hydrogen readily forms covalent compounds with most non-metallic elements, most of the hydrogen on Earth exists in molecular forms such as in the form of water or organic compounds. Hydrogen plays a particularly important role in acid–base reactions as many acid-base reactions involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a negative charge (i.e., anion) known as a hydride, or as a positively charged (i.e., cation) 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 species than that would suggest. 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, via the mixing of metals with acids. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance,^[10] and that it produces water when burned, a property which later gave it its name: in Greek, hydrogen means "water-former".

Industrial production is mainly from the steam reforming of natural gas, and less often from more energy-intensive hydrogen production methods like the electrolysis of water.^[11] Most hydrogen is employed near its production site, with the two largest uses being fossil fuel processing (e.g., hydrocracking) and ammonia production, mostly for the fertilizer market.

Hydrogen is a concern in metallurgy as it can embrittle many metals,^[12] complicating the design of pipelines and storage tanks.^[13]

Properties

Combustion

The Space Shuttle Main Engine burnt hydrogen with oxygen, producing a nearly invisible flame at full thrust.

Hydrogen gas (dihydrogen or molecular hydrogen)^[14] is highly flammable and will burn in air at a very wide range of concentrations between 4% and 75% by volume.^[15] The enthalpy of combustion for hydrogen is −286 kJ/mol:^[16]

2 H₂(g) + O₂(g) → 2 H₂O(l) + 572 kJ (286 kJ/mol)^{[note 2]}

Hydrogen gas forms explosive mixtures with air if it is 4–74% concentrated and with chlorine if it is 5–95% concentrated. The mixtures may be ignited by spark, heat or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C (932 °F).^[17] 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 highly visible plume of a Space Shuttle Solid Rocket Booster. The detection of a burning hydrogen leak may require a flame detector; such leaks can be very dangerous. Hydrogen flames in other conditions are blue, resembling blue natural gas flames.^[18] The destruction of the Hindenburg airship was an infamous example of hydrogen combustion; the cause is debated, but the visible orange flames 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 also potentially dangerous acids.^[19]

Electron energy levels

Depiction of a hydrogen atom with size of central proton shown, and the atomic diameter shown as about twice the Bohr model radius (image not to scale)

The ground state energy level of the electron in a hydrogen atom is −13.6 eV, which is equivalent to an ultraviolet photon of roughly 92 nm wavelength.^[20]

The energy levels of hydrogen can be calculated fairly 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 electromagnetic force attracts electrons and protons to one another, while planets and celestial objects are attracted to each other 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, and therefore only certain allowed energies.^[21]

A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation, Dirac equation or even the Feynman path integral formulation to calculate the probability density of the electron around the proton.^[22] 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—an illustration of how the "planetary orbit" conception of electron motion differs from reality.

Elemental molecular forms

First tracks observed in liquid hydrogen bubble chamber at the Bevatron

There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei.^[23] 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 (¹⁄₂+¹⁄₂); in the parahydrogen form the spins are antiparallel and form a singlet with a molecular spin quantum number of 0 (¹⁄₂–¹⁄₂). At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form, also known as the "normal form".^[24] The equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but because the ortho form is an excited state and has a higher energy than the para form, it is unstable and cannot be purified. At very low temperatures, the equilibrium state is composed almost exclusively of the para form. The liquid and gas phase thermal properties of pure parahydrogen differ significantly from those of the normal form because of differences in rotational heat capacities, as discussed more fully in spin isomers of hydrogen.^[25] The ortho/para distinction also occurs in other hydrogen-containing molecules or functional groups, such as water and methylene, but is of little significance for their thermal properties.^[26]

The uncatalyzed interconversion between para and ortho H₂ increases with increasing temperature; thus rapidly condensed H₂ contains large quantities of the high-energy ortho form that converts to the para form very slowly.^[27] The ortho/para ratio in condensed H₂ is an important consideration in the preparation and storage of liquid hydrogen: the conversion from ortho to para is exothermic and produces enough heat to evaporate some of the hydrogen liquid, leading to loss of liquefied material. Catalysts for the ortho-para interconversion, such as ferric oxide, activated carbon, platinized asbestos, rare earth metals, uranium compounds, chromic oxide, or some nickel^[28] compounds, are used during hydrogen cooling.^[29]

Phases

• Compressed hydrogen
• Liquid hydrogen
• Slush hydrogen
• Solid hydrogen
• Metallic hydrogen

Compounds

Covalent and organic compounds

While H₂ is not very reactive under standard conditions, it does form compounds with most elements. Hydrogen can form compounds with elements that are more electronegative, such as halogens (e.g., F, Cl, Br, I), or oxygen; in these compounds hydrogen takes on a partial positive charge.^[30] When bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a form of medium-strength noncovalent bonding called hydrogen bonding, which is critical to the stability of many biological molecules.^[31][32] Hydrogen also forms compounds with less electronegative elements, such as the metals and metalloids, in which it takes on a partial negative charge. These compounds are often known as hydrides.^[33]

Hydrogen forms a vast array of compounds with carbon called the hydrocarbons, and an even vaster array with heteroatoms that, because of their general association with living things, are called organic compounds.^[34] The study of their properties is known as organic chemistry^[35] and their study in the context of living organisms is known as biochemistry.^[36] By some definitions, "organic" compounds are only required to contain carbon. However, most of them also contain hydrogen, and because it is the carbon-hydrogen bond which gives this class of compounds most of its particular chemical characteristics, carbon-hydrogen bonds are required in some definitions of the word "organic" in chemistry.^[34] Millions of hydrocarbons are known, and they are usually formed by complicated synthetic pathways, which seldom involve elementary hydrogen.

Hydrides

Compounds of hydrogen are often called hydrides, a term that is used fairly loosely. The term "hydride" suggests that the H atom has acquired a negative or anionic character, denoted H⁻, and is used when hydrogen forms a compound with a more electropositive element. The existence of the hydride anion, suggested by Gilbert N. Lewis in 1916 for group I and II salt-like hydrides, was demonstrated by Moers in 1920 by the electrolysis of molten lithium hydride (LiH), producing a stoichiometry quantity of hydrogen at the anode.^[37] For hydrides other than group I and II metals, the term is quite misleading, considering the low electronegativity of hydrogen. An exception in group II hydrides is BeH
2, which is polymeric. In lithium aluminium hydride, the AlH−
4 anion carries hydridic centers firmly attached to the Al(III).

Although hydrides can be formed with almost all main-group elements, the number and combination of possible compounds varies widely; for example, there are over 100 binary borane hydrides known, but only one binary aluminium hydride.^[38] Binary indium hydride has not yet been identified, although larger complexes exist.^[39]

In inorganic chemistry, hydrides can also serve as bridging ligands that link two metal centers in a coordination complex. This function is particularly common in group 13 elements, especially in boranes (boron hydrides) and aluminium complexes, as well as in clustered carboranes.^[40]

Protons and acids

Oxidation of hydrogen removes its electron and gives H⁺, which contains no electrons and a nucleus which is usually composed of one proton. That is why H⁺ is often called a proton. This species is central to discussion of acids. Under the Bronsted-Lowry theory, acids are proton donors, while bases are proton acceptors.A bare proton, H+, cannot exist in solution or in ionic crystals, because of its unstoppable attraction to other atoms or molecules with electrons. Except at the high temperatures associated with plasmas, such protons cannot be removed from the electron clouds of atoms and molecules, and will remain attached to them. However, the term 'proton' is sometimes used loosely and metaphorically to refer to positively charged or cationic hydrogen attached to other species in this fashion, and as such is denoted "H+" without any implication that any single protons exist freely as a species.
To avoid the implication of the naked "solvated proton" in solution, acidic aqueous solutions are sometimes considered to contain a less unlikely fictitious species, termed the "hydronium ion" (H
3O+). However, even in this case, such solvated hydrogen cations are more realistically conceived as being organized into clusters that form species closer to H
9O+
4.^[41] Other oxonium ions are found when water is in acidic solution with other solvents.^[42]

Although exotic on Earth, one of the most common ions in the universe is the H+
3 ion, known as protonated molecular hydrogen or the trihydrogen cation.^[43]

Isotopes

Hydrogen discharge (spectrum) tube

Deuterium discharge (spectrum) tube

Protium, the most common isotope of hydrogen, has one proton and one electron. Unique among all stable isotopes, it has no neutrons (see diproton for a discussion of why others do not exist).

Hydrogen has three naturally occurring isotopes, denoted 1H, 2H and 3H. Other, highly unstable nuclei (4H to 7H) have been synthesized in the laboratory but not observed in nature.^[44][45]

• 1H is the most common hydrogen isotope with an abundance of more than 99.98%. Because the nucleus of this isotope consists of only a single proton, it is given the descriptive but rarely used formal name protium.^[46]
• 2H, the other stable hydrogen isotope, is known as deuterium and contains one proton and one neutron in its nucleus. Essentially all deuterium in the universe is thought to have been produced at the time of the Big Bang, and has endured since that time. Deuterium is not radioactive, and does not represent a significant toxicity hazard. Water enriched in molecules that include deuterium instead of normal hydrogen is called heavy water. Deuterium and its compounds are used as a non-radioactive label in chemical experiments and in solvents for 1H-NMR spectroscopy.^[47] Heavy water is used as a neutron moderator and coolant for nuclear reactors. Deuterium is also a potential fuel for commercial nuclear fusion.^[48]
• 3H is known as tritium and contains one proton and two neutrons in its nucleus. It is radioactive, decaying into helium-3 through beta decay with a half-life of 12.32 years.^[40] It is so radioactive that it can be used in luminous paint, making it useful in such things as watches. The glass prevents the small amount of radiation from getting out.^[49] Small amounts of tritium occur naturally because of the interaction of cosmic rays with atmospheric gases; tritium has also been released during nuclear weapons tests.^[50] It is used in nuclear fusion reactions,^[51] as a tracer in isotope geochemistry,^[52] and specialized in self-powered lighting devices.^[53] Tritium has also been used in chemical and biological labeling experiments as a radiolabel.^[54]

Hydrogen is the only element that has different names for its isotopes in common use today. During the early study of radioactivity, various heavy radioactive isotopes were given their own names, but such names are no longer used, except for deuterium and tritium. The symbols D and T (instead of 2H and 3H) are sometimes used for deuterium and tritium, but the corresponding symbol for protium, P, is already in use for phosphorus and thus is not available for protium.^[55] In its nomenclatural guidelines, the International Union of Pure and Applied Chemistry allows any of D, T, 2H, and 3H to be used, although 2H and 3H are preferred.^[56]

History

Discovery and use

In 1671, Robert Boyle discovered and described the reaction between iron filings and dilute acids, which results in the production of hydrogen gas.^[57][58] In 1766, Henry Cavendish was the first to recognize hydrogen gas as a discrete substance, by naming the gas from a metal-acid reaction "flammable air". He speculated that "flammable air" was in fact identical to the hypothetical substance called "phlogiston"^[59][60] and further finding in 1781 that the gas produces water when burned. He is usually given credit for its discovery as an element.^[6][7] In 1783, Antoine Lavoisier gave the element the name hydrogen (from the Greek ὑδρο- hydro meaning "water" and -γενής genes meaning "creator")^[8] when he and Laplace reproduced Cavendish's finding that water is produced when hydrogen is burned.^[7]

Antoine-Laurent de Lavoisier

Lavoisier produced hydrogen for his experiments on mass conservation by reacting a flux of steam with metallic iron through an incandescent iron tube heated in a fire. Anaerobic oxidation of iron by the protons of water at high temperature can be schematically represented by the set of 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 liquefied for the first time by James Dewar in 1898 by using regenerative cooling and his invention, the vacuum flask.^[7] He produced solid hydrogen the next year.^[7] Deuterium was discovered in December 1931 by Harold Urey, and tritium was prepared in 1934 by Ernest Rutherford, Mark Oliphant, and Paul Harteck.^[6] Heavy water, which consists of deuterium in the place of regular hydrogen, was discovered by Urey's group in 1932.^[7] François Isaac de Rivaz built the first de Rivaz engine, an internal combustion engine powered by a mixture of hydrogen and oxygen in 1806. Edward Daniel Clarke invented the hydrogen gas blowpipe in 1819. The Döbereiner's lamp and limelight were invented in 1823.^[7]

The first hydrogen-filled balloon was invented by Jacques Charles in 1783.^[7] Hydrogen provided the lift for the first reliable form of air-travel following the 1852 invention of the first hydrogen-lifted airship by Henri Giffard.^[7] German count Ferdinand von Zeppelin promoted the idea of rigid airships lifted by hydrogen that later were called Zeppelins; the first of which had its maiden flight in 1900.^[7] Regularly scheduled flights started in 1910 and by the outbreak of World War I in August 1914, they had carried 35,000 passengers without a serious incident. Hydrogen-lifted airships were used as observation platforms and bombers during the war.

The first non-stop transatlantic crossing was made by the British airship R34 in 1919. Regular passenger service resumed in the 1920s and the discovery of helium reserves in the United States promised increased safety, but the U.S. government refused to sell the gas for this purpose. Therefore, H₂ was used in the Hindenburg airship, which was destroyed in a midair fire over New Jersey on 6 May 1937.^[7] The incident was broadcast live on radio and filmed. Ignition of leaking hydrogen is widely assumed to be the cause, but later investigations pointed to the ignition of the aluminized fabric coating by static electricity. But the damage to hydrogen's reputation as a lifting gas was already done.

In the same year the first hydrogen-cooled turbogenerator went into service with gaseous hydrogen as a coolant in the rotor and the stator in 1937 at Dayton, Ohio, by the Dayton Power & Light Co,^[61] because of the thermal conductivity of hydrogen gas this is the most common type in its field today.

The nickel hydrogen battery was used for the first time in 1977 aboard the U.S. Navy's Navigation technology satellite-2 (NTS-2).^[62] For example, the ISS,^[63] Mars Odyssey^[64] and the Mars Global Surveyor^[65] are equipped with nickel-hydrogen batteries. In the dark part of its orbit, the Hubble Space Telescope is also powered by nickel-hydrogen batteries, which were finally replaced in May 2009,^[66] more than 19 years after launch, and 13 years over their design life.^[67]

Role in quantum theory

Hydrogen emission spectrum lines in the visible range. These are the four visible lines of the Balmer series

Because of its relatively simple atomic structure, consisting only of a proton and an electron, the hydrogen atom, together with the spectrum of light produced from it or absorbed by it, has been central to the development of the theory of atomic structure.^[68] Furthermore, the corresponding simplicity of the hydrogen molecule and the corresponding cation H+
2 allowed fuller understanding of the nature of the chemical bond, which followed shortly after the quantum mechanical treatment of the hydrogen atom had been developed in the mid-1920s.

One of the first quantum effects to be explicitly noticed (but not understood at the time) was a Maxwell observation involving hydrogen, half a century before full quantum mechanical theory arrived. Maxwell observed that the specific heat capacity of H₂ unaccountably departs from that of a diatomic gas below room temperature and begins to increasingly resemble that of a monatomic gas at cryogenic temperatures. According to quantum theory, this behavior arises from the spacing of the (quantized) rotational energy levels, which are particularly wide-spaced in H₂ because of its low mass. These widely spaced levels inhibit equal partition of heat energy into rotational motion in hydrogen at low temperatures. Diatomic gases composed of heavier atoms do not have such widely spaced levels and do not exhibit the same effect.^[69]

Natural occurrence

NGC 604, a giant region of ionized hydrogen in the Triangulum Galaxy

Hydrogen, as atomic H, is the most abundant chemical element in the universe, making up 75% of normal matter by mass and over 90% by number of atoms (most of the mass of the universe, however, is not in the form of chemical-element type matter, but rather is postulated to occur as yet-undetected forms of mass such as dark matter and dark energy).^[70] This element is found in great abundance in stars and gas giant planets. Molecular clouds of H₂ are associated with star formation. Hydrogen plays a vital role in powering stars through the proton-proton reaction and the CNO cycle nuclear fusion.^[71]

Throughout the universe, hydrogen is mostly found in the atomic and plasma states whose properties are quite different from molecular hydrogen. As a plasma, hydrogen's electron and proton are not bound together, resulting in very high electrical conductivity and high emissivity (producing the light from the Sun and other stars). The charged particles are highly influenced by magnetic and electric fields. For example, in the solar wind they interact with the Earth's magnetosphere giving rise to Birkeland currents and the aurora. Hydrogen is found in the neutral atomic state in the interstellar medium. The large amount of neutral hydrogen found in the damped Lyman-alpha systems is thought to dominate the cosmological baryonic density of the Universe up to redshift z=4.^[72]

Under ordinary conditions on Earth, elemental hydrogen exists as the diatomic gas, H₂. However, hydrogen gas is very rare in the Earth's atmosphere (1 ppm by volume) because of its light weight, which enables it to escape from Earth's gravity more easily than heavier gases. However, hydrogen is the third most abundant element on the Earth's surface,^[73] mostly in the form of chemical compounds such as hydrocarbons and water.^[40] Hydrogen gas is produced by some bacteria and algae and is a natural component of flatus, as is methane, itself a hydrogen source of increasing importance.^[74]

A molecular form called protonated molecular hydrogen (H+
3) is found in the interstellar medium, where it is generated by ionization of molecular hydrogen from cosmic rays. This charged ion has also been observed in the upper atmosphere of the planet Jupiter. The ion is relatively stable in the environment of outer space due to the low temperature and density. H+
3 is one of the most abundant ions in the Universe, and it plays a notable role in the chemistry of the interstellar medium.^[75] Neutral triatomic hydrogen H₃ can only exist in an excited form and is unstable.^[76] By contrast, the positive hydrogen molecular ion (H+
2) is a rare molecule in the universe.

Production

H2 is produced in chemistry and biology laboratories, often as a by-product of other reactions; in industry for the hydrogenation of unsaturated substrates; and in nature as a means of expelling reducing equivalents in biochemical reactions.

Metal-acid

In the laboratory, H
2 is usually prepared by the reaction of dilute non-oxidizing acids on some reactive metals such as zinc with Kipp's apparatus.

Zn + 2 H+ → Zn2+ + H
2

Aluminium can also produce H
2 upon treatment with bases:

2 Al + 6 H
2O + 2 OH− → 2 Al(OH)−
4 + 3 H
2

The electrolysis of water is a simple method of producing hydrogen. A low voltage current is run through the water, and gaseous oxygen forms at the anode while gaseous hydrogen forms at the cathode. Typically the cathode is made from platinum or another inert metal when producing hydrogen for storage. If, however, the gas is to be burnt on site, oxygen is desirable to assist the combustion, and so both electrodes would be made from inert metals. (Iron, for instance, would oxidize, and thus decrease the amount of oxygen given off.) The theoretical maximum efficiency (electricity used vs. energetic value of hydrogen produced) is in the range 80–94%.^[77]

2 H
2O(l) → 2 H
2(g) + O
2(g)

In 2007, it was discovered that an alloy of aluminium and gallium in pellet form added to water could be used to generate hydrogen. The process also creates alumina, but the expensive gallium, which prevents the formation of an oxide skin on the pellets, can be re-used. This has important potential implications for a hydrogen economy, as hydrogen can be produced on-site and does not need to be transported.^[78]

Steam reforming

Hydrogen can be prepared in several different ways, but economically the most important processes involve removal of hydrogen from hydrocarbons. Commercial bulk hydrogen is usually produced by the steam reforming of natural gas.^[79] At high temperatures (1000–1400 K, 700–1100 °C or 1300–2000 °F), steam (water vapor) reacts with methane to yield carbon monoxide and H
2.

CH
4 + H
2O → CO + 3 H
2

This reaction is favored at low pressures but is nonetheless conducted at high pressures (2.0  MPa, 20 atm or 600 inHg). This is because high-pressure H
2 is the most marketable product and Pressure Swing Adsorption (PSA) purification systems work better at higher pressures. The product mixture is known as "synthesis gas" because it is often used directly for the production of methanol and related compounds. Hydrocarbons other than methane can be used to produce synthesis gas with varying product ratios. One of the many complications to this highly optimized technology is the formation of coke or carbon:

CH
4 → C + 2 H
2

Consequently, steam reforming typically employs an excess of H
2O. Additional hydrogen can be recovered from the steam by use of carbon monoxide through the water gas shift reaction, especially with an iron oxide catalyst. This reaction is also a common industrial source of carbon dioxide:^[79]

CO + H
2O → CO
2 + H
2

Other important methods for H
2 production include partial oxidation of hydrocarbons:^[80]

2 CH
4 + O
2 → 2 CO + 4 H
2

and the coal reaction, which can serve as a prelude to the shift reaction above:^[79]

C + H
2O → CO + H
2

Hydrogen is sometimes produced and consumed in the same industrial process, without being separated. In the Haber process for the production of ammonia, hydrogen is generated from natural gas.^[81] Electrolysis of brine to yield chlorine also produces hydrogen as a co-product.^[82]

Thermochemical

There are more than 200 thermochemical cycles which can be used for water splitting, around a dozen of these cycles such as the iron oxide cycle, cerium(IV) oxide–cerium(III) oxide cycle, zinc zinc-oxide cycle, sulfur-iodine cycle, copper-chlorine cycle and hybrid sulfur cycle are under research and in testing phase to produce hydrogen and oxygen from water and heat without using electricity.^[83] A number of laboratories (including in France, Germany, Greece, Japan, and the USA) are developing thermochemical methods to produce hydrogen from solar energy and water.^[84]

Anaerobic corrosion

Under anaerobic conditions, iron and steel alloys are slowly oxidized by the protons of water concomitantly reduced in molecular hydrogen (H
2). The anaerobic corrosion of iron leads first to the formation of ferrous hydroxide (green rust) and can be described by the following reaction:

Fe + 2 H
2O → Fe(OH)
2 + H
2

In its turn, under anaerobic conditions, the ferrous hydroxide (Fe(OH)
2 ) can be oxidized by the protons of water to form magnetite and molecular hydrogen. This process is described by the Schikorr reaction:

3 Fe(OH)
2 → Fe
3O
4 + 2 H
2O + H
2

ferrous hydroxide → magnetite + water + hydrogen

The well crystallized magnetite (Fe
3O
4) is thermodynamically more stable than the ferrous hydroxide (Fe(OH)
2 ).

This process occurs during the anaerobic corrosion of iron and steel in oxygen-free groundwater and in reducing soils below the water table.

Geological occurrence: the serpentinization reaction

In the absence of atmospheric oxygen (O
2), in deep geological conditions prevailing far away from Earth atmosphere, hydrogen (H
2) is produced during the process of serpentinization by the anaerobic oxidation by the water protons (H⁺) of the ferrous (Fe²⁺) silicate present in the crystal lattice of the fayalite (Fe
2SiO
4, the olivine iron-endmember). The corresponding reaction leading to the formation of magnetite (Fe
3O
4), quartz (SiO
2) and hydrogen (H
2) is the following:

3Fe
2SiO
4 + 2 H
2O → 2 Fe
3O
4 + 3 SiO
2 + 3 H
2

fayalite + water → magnetite + quartz + hydrogen

This reaction closely resembles the Schikorr reaction observed in the anaerobic oxidation of the ferrous hydroxide in contact with water.

Formation in transformers

From all the fault gases formed in power transformers, hydrogen is the most common and is generated under most fault conditions; thus, formation of hydrogen is an early indication of serious problems in the transformer's life cycle.^[85]

Xylose

In 2014 a low-temperature 50 °C (122 °F), atmospheric-pressure enzyme-driven process to convert xylose into hydrogen with nearly 100% of the theoretical yield was announced. The process employs 13 enzymes, including a novel polyphosphate xylulokinase (XK).^[86][87]

Applications

Consumption in processes

Large quantities of H
2 are needed in the petroleum and chemical industries. The largest application of H
2 is for the processing ("upgrading") of fossil fuels, and in the production of ammonia. The key consumers of H
2 in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking. H
2 has several other important uses. H
2 is used as a hydrogenating agent, particularly in increasing the level of saturation of unsaturated fats and oils (found in items such as margarine), and in the production of methanol. It is similarly the source of hydrogen in the manufacture of hydrochloric acid. H
2 is also used as a reducing agent of metallic ores.^[88]

Hydrogen is highly soluble in many rare earth and transition metals^[89] and is soluble in both nanocrystalline and amorphous metals.^[90] Hydrogen solubility in metals is influenced by local distortions or impurities in the crystal lattice.^[91] These properties may be useful when hydrogen is purified by passage through hot palladium disks, but the gas's high solubility is a metallurgical problem, contributing to the embrittlement of many metals,^[12] complicating the design of pipelines and storage tanks.^[13]

Apart from its use as a reactant, H
2 has wide applications in physics and engineering. It is used as a shielding gas in welding methods such as atomic hydrogen welding.^[92][93] H₂ is used as the rotor coolant in electrical generators at power stations, because it has the highest thermal conductivity of any gas. Liquid H₂ is used in cryogenic research, including superconductivity studies.^[94] Because H
2 is lighter than air, having a little more than ¹⁄₁₄ of the density of air, it was once widely used as a lifting gas in balloons and airships.^[95]

In more recent applications, hydrogen is used pure or mixed with nitrogen (sometimes called forming gas) as a tracer gas for minute leak detection. Applications can be found in the automotive, chemical, power generation, aerospace, and telecommunications industries.^[96] Hydrogen is an authorized food additive (E 949) that allows food package leak testing among other anti-oxidizing properties.^[97]

Hydrogen's rarer isotopes also each have specific applications. Deuterium (hydrogen-2) is used in nuclear fission applications as a moderator to slow neutrons, and in nuclear fusion reactions.^[7] Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects.^[98] Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs,^[99] as an isotopic label in the biosciences,^[54] and as a radiation source in luminous paints.^[100]

The triple point temperature of equilibrium hydrogen is a defining fixed point on the ITS-90 temperature scale at 13.8033 kelvins.^[101]

Coolant

Hydrogen is commonly used in power stations as a coolant in generators due to a number of favorable properties that are a direct result of its light diatomic molecules. These include low density, low viscosity, and the highest specific heat and thermal conductivity of all gases.

Energy carrier

Hydrogen is not an energy resource,^[102] except in the hypothetical context of commercial nuclear fusion power plants using deuterium or tritium, a technology presently far from development.^[103] The Sun's energy comes from nuclear fusion of hydrogen, but this process is difficult to achieve controllably on Earth.^[104] Elemental hydrogen from solar, biological, or electrical sources require more energy to make it than is obtained by burning it, so in these cases hydrogen functions as an energy carrier, like a battery. Hydrogen may be obtained from fossil sources (such as methane), but these sources are unsustainable.^[102]
The energy density per unit volume of both liquid hydrogen and compressed hydrogen gas at any practicable pressure is significantly less than that of traditional fuel sources, although the energy density per unit fuel mass is higher.^[102] Nevertheless, elemental hydrogen has been widely discussed in the context of energy, as a possible future carrier of energy on an economy-wide scale.^[105] For example, CO
2 sequestration followed by carbon capture and storage could be conducted at the point of H
2 production from fossil fuels.^[106] Hydrogen used in transportation would burn relatively cleanly, with some NO_x emissions,^[107] but without carbon emissions.^[106] However, the infrastructure costs associated with full conversion to a hydrogen economy would be substantial.^[108]

Semiconductor industry

Hydrogen is employed to saturate broken ("dangling") bonds of amorphous silicon and amorphous carbon that helps stabilizing material properties.^[109] It is also a potential electron donor in various oxide materials, including ZnO,^[110][111] SnO₂, CdO, MgO,^[112] ZrO₂, HfO₂, La₂O₃, Y₂O₃, TiO₂, SrTiO₃, LaAlO₃, SiO₂, Al₂O₃, ZrSiO₄, HfSiO₄, and SrZrO₃.^[113]

Biological reactions

H₂ is a product of some types of anaerobic metabolism and is produced by several microorganisms, usually via reactions catalyzed by iron- or nickel-containing enzymes called hydrogenases. These enzymes catalyze the reversible redox reaction between H₂ and its component two protons and two electrons. Creation of hydrogen gas occurs in the transfer of reducing equivalents produced during pyruvate fermentation to water.^[114]
Water splitting, in which water is decomposed into its component protons, electrons, and oxygen, occurs in the light reactions in all photosynthetic organisms. Some such organisms, including the alga Chlamydomonas reinhardtii and cyanobacteria, have evolved a second step in the dark reactions in which protons and electrons are reduced to form H₂ gas by specialized hydrogenases in the chloroplast.^[115] Efforts have been undertaken to genetically modify cyanobacterial hydrogenases to efficiently synthesize H₂ gas even in the presence of oxygen.^[116] Efforts have also been undertaken with genetically modified alga in a bioreactor.^[117]

Safety and precautions

Hydrogen poses a number of hazards to human safety, from potential detonations and fires when mixed with air to being an asphyxiant in its pure, oxygen-free form.^[118] In addition, liquid hydrogen is a cryogen and presents dangers (such as frostbite) associated with very cold liquids.^[119] Hydrogen dissolves in many metals, and, in addition to leaking out, may have adverse effects on them, such as hydrogen embrittlement,^[120] leading to cracks and explosions.^[121] Hydrogen gas leaking into external air may spontaneously ignite. Moreover, hydrogen fire, while being extremely hot, is almost invisible, and thus can lead to accidental burns.^[122]
Even interpreting the hydrogen data (including safety data) is confounded by a number of phenomena. Many physical and chemical properties of hydrogen depend on the parahydrogen/orthohydrogen ratio (it often takes days or weeks at a given temperature to reach the equilibrium ratio, for which the data is usually given). Hydrogen detonation parameters, such as critical detonation pressure and temperature, strongly depend on the container geometry.^[118]

Sunspot

From Wikipedia, the free encyclopedia

Sunspot region 2192,^[1] during the partial solar eclipse of 23 October 2014

Detailed view, 13 December 2006

Heliophysics
Phenomena
Coronal mass ejection Geomagnetic storm Solar EMP Solar flare Solar prominence Solar proton event Solar superstorm Sunspot

Sunspots are temporary phenomena on the photosphere of the Sun that appear visibly as dark spots compared to surrounding regions. They correspond to concentrations of magnetic field that inhibit convection and result in reduced surface temperature compared to the surrounding photosphere.

Sunspots usually appear as pairs, with each spot having the opposite magnetic polarity of the other.^[2]

Although they are at temperatures of roughly 3,000–4,500 K (2,700–4,200 °C), the contrast with the surrounding material at about 5,780 K (5,500 °C) leaves them clearly visible as dark spots, as the luminous intensity of a heated black body (closely approximated by the photosphere) is proportional to the fourth power of its temperature. If the sunspot were isolated from the surrounding photosphere it would be brighter than the Moon.^[3] Sunspots expand and contract as they move across the surface of the Sun and can be as small as 16 kilometers (10 mi)^[4] and as large as 160,000 kilometers (100,000 mi)^[5] in diameter, making the larger ones visible from Earth without the aid of a telescope.^[6] They may also travel at relative speeds ("proper motions") of a few hundred meters per second when they first emerge onto the solar photosphere.

Manifesting intense magnetic activity, sunspots host secondary phenomena such as coronal loops (prominences) and reconnection events. Most solar flares and coronal mass ejections originate in magnetically active regions around visible sunspot groupings. Similar phenomena indirectly observed on stars other than the sun are commonly called starspots and both light and dark spots have been measured.^[7]

History

Prehistoric evidence

Studies of stratigraphic data have suggested that the solar cycles have been active for hundreds of millions of years, if not longer; measuring varves in precambrian sedimentary rock has revealed repeating peaks in layer thickness, with a pattern repeating approximately every eleven years. It is possible that the early atmosphere on Earth was more sensitive to changes in solar radiation than today, so that greater glacial melting (and thicker sediment deposits) could have occurred during years with greater sunspot activity.^[8][9] This would presume annual layering; however, alternate explanations (diurnal) have also been proposed.^[10]

Analysis of tree rings has revealed a detailed picture of past solar cycles: Dendrochronologically dated radiocarbon concentrations have allowed for a reconstruction of sunspot activity dating back 11,400 years, far beyond the four centuries of available, reliable records from direct solar observation.^[11]

Early observations

A drawing of a sunspot in the Chronicles of John of Worcester

The earliest surviving record of sunspot observation dates from 364 BC, based on comments by Chinese astronomer Gan De in a star catalogue.^[12] By 28 BC, Chinese astronomers were regularly recording sunspot observations in official imperial records.^[13]

The first clear mention of a sunspot in Western literature, around 300 BC, was by the ancient Greek scholar Theophrastus, student of Plato and Aristotle and successor to the latter.^[14] A more recent sunspot observation was made on 17 March 807 AD by the Benedictine monk Adelmus, who observed a large sunspot that was visible for eight days; however, Adelmus incorrectly concluded he was observing a transit of Mercury.^[15] A large sunspot was also seen at the time of Charlemagne's death in 813 AD.^[16] Sunspot activity in 1129 was described by John of Worcester, and Averroes provided a description of sunspots later in the 12th century;^[17] however, these observations were also misinterpreted as planetary transits, until Galileo gave the correct explanation in 1612.^[18]

17th and 18th centuries

Sunspots in 1794 Samuel Dunn Map

Sunspots were first observed telescopically in late 1610 by the English astronomer Thomas Harriot and Frisian astronomers Johannes and David Fabricius, who published a description in June 1611. At the latter time, Galileo had been showing sunspots to astronomers in Rome, and Christoph Scheiner had probably been observing the spots for two or three months using an improved helioscope of his own design. The ensuing priority dispute between Galileo and Scheiner, neither of whom knew of the Fabricius' work, was thus as pointless as it was bitter.

Sunspots had some importance in the debate over the nature of the Solar System. They showed that the Sun rotated, and their comings and goings showed that the Sun changed, contrary to Aristotle (who taught that all celestial bodies were perfect, unchanging spheres).

Rudolf Wolf studied the historical record in an attempt to establish a database on past cyclic variations. His database extended only to 1700, although the technology and techniques for careful solar observations were first available in 1610. Gustav Spörer later suggested a 70-year period before 1716 in which sunspots were rarely observed as the reason for Wolf's inability to extend the cycles into the 17th century.

Sunspots were rarely recorded during the second part of 17th century. Later analysis revealed the problem not to be a lack of observational data but included references to negative observations. Building upon Spörer's earlier work, Edward Maunder suggested that the Sun had changed from a period in which sunspots all but disappeared from the solar surface to a renewal of sunspot cycles starting in about 1700. Adding to this understanding of the absence of solar cycles were observations of aurorae, which were absent at the same time. Even the lack of a solar corona during solar eclipses was noted prior to 1715. The period of low sunspot activity from 1645 to 1717 is known as the "Maunder Minimum".

19th century

The cyclic variation of the number of sunspots was first observed by Heinrich Schwabe between 1826 and 1843 and led Wolf to make systematic observations starting in 1848. The Wolf number is a measure of individual spots and spot groupings, which correlates to a number of solar observables. Also in 1848, Joseph Henry projected an image of the Sun onto a screen and determined that sunspots were cooler than the surrounding surface.^[19]

After the resumption of sunspot activity, Heinrich Schwabe in 1844 in Astronomische Nachrichten (Astronomical News) reported a periodic change in the number of sunspots.

The Sun emitted an extremely powerful flare on its visible hemisphere on 1 September 1859, leading to what is known as the Carrington Event. It interrupted electrical telegraph service and caused visible aurorae as far south as Havana, Hawaii, and Rome with similar activity in the southern hemisphere.

20th century

The American solar astronomer George Ellery Hale, as an undergraduate at MIT, invented the spectroheliograph, with which he made the discovery of solar vortices. In 1908, Hale used a modified spectroheliograph to show that the spectra of hydrogen exhibited the Zeeman effect whenever the area of view passed over a sunspot on the solar disc. This was the first indication that sunspots were basically magnetic phenomena, which appeared in pairs that corresponded with two magnetic poles of opposite polarity.^[20] Subsequent work by Hale demonstrated a strong tendency for east-west alignment of magnetic polarities in sunspots, with mirror symmetry across the solar equator; and that the magnetic polarity for sunspots in each hemisphere switched orientation, from one sunspot cycle to the next.^[21] This systematic property of sunspot magnetic fields is now commonly referred to as the "Hale–Nicholson law",^[22] or in many cases simply "Hale's law".

21st century

The most powerful flare observed by satellite instrumentation began on 4 November 2003 at 19:29 UTC, and saturated instruments for 11 minutes. Region 486 has been estimated to have produced an X-ray flux of X28. Holographic and visual observations indicate significant activity continued on the far side of the Sun.

Measurements made in the latter part of the 2000s (decade) and based also on observation of infrared spectral lines, have suggested that sunspot activity may again be disappearing, possibly leading to a new minimum.^[23] From 2007 to 2009, sunspot levels were far below average. In 2008, the Sun was spot-free 73 percent of the time, extreme even for a solar minimum. Only 1913 was more pronounced, with 85 percent of that year clear. The Sun continued to languish through mid-December 2009, when the largest group of sunspots to emerge for several years appeared. Even then, sunspot levels remained well below normal.^[24]

Nasa's 2006 prediction. At 2010/2011, the sunspot count was expected to be at its maximum, but in reality in 2010 it was still at its minimum.

In 2006, NASA made a prediction for the next sunspot maximum, being between 150 and 200 around the year 2011 (30–50% stronger than cycle 23), followed by a weak maximum at around 2022.^[25][26] The prediction did not come true. Instead, the sunspot cycle in 2010 was still at its minimum, where it should have been near its maximum, which shows the Sun's unusually low current activity.^[27]

Due to a missing jet stream, fading spots, and slower activity near the poles, independent scientists of the National Solar Observatory (NSO) and the Air Force Research Laboratory (AFRL) predicted in 2011 that the next 11-year solar sunspot cycle, Cycle 25, would be greatly reduced or might not happen at all.^[28]

Cycle 24 is now well underway (as of March 2013^[update]). Measurements indicate that the minimum occurred around December 2008 and the next maximum was predicted to reach a sunspot number of 90 around May 2013.^[29] The monthly mean sunspot number in the northern solar hemisphere peaked in November 2011 but that in the southern hemisphere appears to have peaked in February 2014, giving a peak total monthly mean of 102 in that month. Subsequent months have seen a decline to around 70 by June 2014.^[30] In October of 2014, the sunspot known as AR 12192 was shown to be the largest observed since 1990.^[31] The flare that erupted from this sunspot was classified as an X3.1-class solar storm.^[32]

Physics

A sunspot viewed close-up in ultraviolet light, taken by the TRACE spacecraft

Although the details of sunspot generation are still a matter of research, it appears that sunspots are the visible counterparts of magnetic flux tubes in the Sun's convective zone that get "wound up" by differential rotation. If the stress on the tubes reaches a certain limit, they curl up like a rubber band and puncture the Sun's surface. Convection is inhibited at the puncture points; the energy flux from the Sun's interior decreases; and with it surface temperature.

The Wilson effect tells us that sunspots are actually depressions on the Sun's surface. Observations using the Zeeman effect show that prototypical sunspots come in pairs with opposite magnetic polarity. From cycle to cycle, the polarities of leading and trailing (with respect to the solar rotation) sunspots change from north/south to south/north and back. Sunspots usually appear in groups.

The sunspot itself can be divided into two parts:

• The central umbra, which is the darkest part, where the magnetic field is approximately vertical (normal to the Sun's surface).
• The surrounding penumbra, which is lighter, where the magnetic field is more inclined.

Magnetic pressure should tend to remove field concentrations, causing the sunspots to disperse, but sunspot lifetimes are measured in days or even weeks. In 2001, observations from the Solar and Heliospheric Observatory (SOHO) using sound waves traveling below the Sun's photosphere (local helioseismology) were used to develop a three-dimensional image of the internal structure below sunspots; these observations show that there is a powerful downdraft underneath each sunspot, forming a rotating vortex that concentrates the magnetic field.^[33] Sunspots can thus be thought of as self-perpetuating storms, analogous in some ways to terrestrial hurricanes.

Butterfly diagram showing paired Spörer's law behavior

Sunspot activity cycles about every eleven years. The point of highest sunspot activity during this cycle is known as Solar Maximum, and the point of lowest activity is Solar Minimum. Early in the cycle, sunspots appear in the higher latitudes and then move towards the equator as the cycle approaches maximum: this is called Spörer's law.

Wolf number sunspot index displays various periods, the most prominent of which is at about 11 years in the mean. This period is also observed in most other expressions of solar activity and is deeply linked to a variation in the solar magnetic field that changes polarity with this period, too.

The modern understanding of sunspots starts with George Ellery Hale, who first linked magnetic fields and sunspots in 1908.^[20] Hale suggested that the sunspot cycle period is 22 years, covering two polar reversals of the solar magnetic dipole field. Horace W. Babcock later proposed a qualitative model for the dynamics of the solar outer layers. The Babcock Model explains that magnetic fields cause the behavior described by Spörer's law, as well as other effects, which are twisted by the Sun's rotation.

Variation

400 year sunspot history

11,000 year sunspot reconstruction

Sunspot populations quickly rise and more slowly fall on an irregular cycle of 11 years, although significant variations in the number of sunspots attending the 11-year period are known over longer spans of time. For example, from 1900 to the 1960s, the solar maxima trend of sunspot count has been upward; from the 1960s to the present, it has diminished somewhat.^[34] Over the last decades the Sun has had a markedly high average level of sunspot activity; it was last similarly active over 8,000 years ago.^[11]

Sunspots are caused by solar magnetic fields in the photosphere and the associated centennial variations in magnetic fields in the corona and heliosphere have also been deduced using carbon-14 and beryllium-10 cosmogenic isotopes stored in terrestrial reservoirs such as ice sheets and tree rings^[35] and by using historic observations of geomagnetic storm activity, which bridge the time gap between the end of the usable cosmogenic isotope data and the start of modern spacecraft data.^[36] These variations have been successfully reproduced using models which employ magnetic flux continuity equations and observed sunspot numbers to quantify the emergence of magnetic flux from the top of the solar atmosphere and into the heliosphere,^[37] showing that sunspot observations, geomagnetic activity and cosmogenic isotopes are giving a coherent understanding of solar activity variations.

The number of sunspots correlates with the intensity of solar radiation over the period since 1979, when satellite measurements of absolute radiative flux became available. Since sunspots are darker than the surrounding photosphere it might be expected that more sunspots would lead to less solar radiation and a decreased solar constant. However, the surrounding margins of sunspots are brighter than the average, and so are hotter; overall, more sunspots increase the Sun's solar constant or brightness. The variation caused by the sunspot cycle to solar output is relatively small, on the order of 0.1% of the solar constant (a peak-to-trough range of 1.3 W·m⁻² compared to 1366 W·m⁻² for the average solar constant).^[38][39] Sunspots were rarely observed during the Maunder Minimum in the second part of the 17th century (approximately from 1645 to 1715).

The 11-year solar cycles are numbered sequentially, starting with the observations made in the 1750s.^[40]

Observation

The Swedish 1-m Solar Telescope at Roque de los Muchachos Observatory

Sunspots are observed with land-based and Earth-orbiting solar telescopes. These telescopes use filtration and projection techniques for direct observation, in addition to various types of filtered cameras. Specialized tools such as spectroscopes and spectrohelioscopes are used to examine sunspots and sunspot areas. Artificial eclipses allow viewing of the circumference of the Sun as sunspots rotate through the horizon.

Since looking directly at the Sun with the naked eye permanently damages vision, amateur observation of sunspots is generally conducted indirectly using projected images, or directly through protective filters. Small sections of very dark filter glass, such as a #14 welder's glass are effective. A telescope eyepiece can project the image, without filtration, onto a white screen where it can be viewed indirectly, and even traced, to follow sunspot evolution. Special purpose hydrogen-alpha narrow bandpass filters as well as aluminum coated glass attenuation filters (which have the appearance of mirrors due to their extremely high optical density) on the front of a telescope provide safe observation through the eyepiece.

Application

Detail of a sunspot in 2005. The granulation of the Sun's surface can be seen clearly

Due to its link to other kinds of solar activity, sunspot occurrence can be used to help predict space weather, the state of the ionosphere, and hence the conditions of short-wave radio propagation or satellite communications. Solar activity (and the sunspot cycle) are frequently discussed in the context of global warming; Jack Eddy noted the apparent correlation between the Maunder Minimum of sunspot occurrence and the Little Ice Age in European winter climate.^[41] Sunspots themselves, in terms of the magnitude of their radiant-energy deficit, have only a weak effect on the terrestrial climate^[42] in a direct sense. On longer time scales, such as the solar cycle, other magnetic phenomena (faculae and the chromospheric network) do correlate with sunspot occurrence. It is these other features that make the solar constant increase slightly at sunspot maxima, when naively one might expect that sunspots would make it decrease.^[43]

British economist William Stanley Jevons suggested in the 1870s that there is a relationship between sunspots and business cycle crises. Jevons reasoned that sunspots affect Earth's weather, which, in turn, influences crops and, therefore, the economy.^[44]

Spots on other stars

In 1947, G. E. Kron proposed that starspots were the reason for periodic changes in brightness on red dwarfs.^[7] Since the mid-1990s, starspot observations have been made using increasingly powerful techniques yielding more and more detail: photometry showed starspot growth and decay and showed cyclic behavior similar to the Sun's; spectroscopy examined the structure of starspot regions by analyzing variations in spectral line splitting due to the Zeeman Effect; Doppler imaging showed differential rotation of spots for several stars and distributions different from the Sun's; spectral line analysis measured the temperature range of spots and the stellar surfaces. For example, in 1999, Strassmeier reported the largest cool starspot ever seen rotating the giant K0 star XX Triangulum (HD 12545) with a temperature of 3,500 K (3,230 °C), together with a warm spot of 4,800 K (4,530 °C).^[7][45]

Gallery

Play media

Sunspot NOAA 875.

Play media

A flare from sunspot NOAA 875.

Play media

This visualization tracks the emergence and evolution of a sunspot group as seen starting in early February 2011 and continuing for two weeks. Images are sampled one hour apart. The camera tracks the movement of the solar rotation. At this scale, a 'shimmer' of the solar surface is visible, created by the turnover of convection cells.

Play media

Groups of sunspots grow and die over a matter of days. This is a movie built from images taken by the SDO/HMI instrument over the course of 13 days during the rise of solar cycle 24.

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  Sunspots, September 2011.
• 

  A view of the coronal structure above a different sunspot seen in October 2010.
• 

  Sunspot 923 at sunset and in solar scope.
• 

  Sunset Superior Mirage of sunspot #930.
• 

  Sunset in Bangladesh, January 2004.