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Hydrogen

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Hydrogen

Atomic Mass	1.00794
Electron Configuration	1s1
Oxidation States	+1, -1
Year Discovered	1766

Hydrogen

Atomic Mass	1.00794
Electron Configuration	1s1
Oxidation States	+1, -1
Year Discovered	1766

Identifiers

Element Name	Hydrogen
Element Symbol	H
InChI	InChI=1S/H
InChIKey	YZCKVEUIGOORGS-UHFFFAOYSA-N

Properties

Atomic Weight	[1.007 84, 1.008 11] 1.00794 1.008 [1.00784,1.00811]
Electron Configuration	1s¹
Atomic Radius	Van der Waals Atomic Radius :120 pm (Van der Waals) Empirical Atomic Radius :25pm (Empirical) Covalent Atomic Radius :31(5) pm (Covalent)
Oxidation States	+1, -1 -1, +1 (an amphoteric oxide)
Ground Level	²S_1/2
Ionization Energy	13.598 eV 13.598434599702 ± 0.000000000012 eV (Theoretical value.)
Electronegativity	Pauling Scale Electronegativity :2.2(Pauling Scale) Allen Scale Electronegativity :2.3(Allen Scale)
Electron Affinity	0.754eV 0.77eV
Atomic Spectra	Lines Holdings Levels Holdings
Physical Description	Gas
Element Classification	Non-metal
Element Period Number	1
Element Group Number	1
Density	0.00008988 grams per cubic centimeter
Melting Point	13.81 K (-259.34°C or -434.81°F) -259.16°C
Boiling Point	20.28 K (-252.87°C or -423.17°F)
Estimated Crustal Abundance	1.40×10³ milligrams per kilogram
Estimated Oceanic Abundance	1.08×10⁵ milligrams per liter

History

The name derives from the Greek hydro for "water" and genes for "forming" because it burned in air to form water. Hydrogen was discovered by the English physicist Henry Cavendish in 1766.

Scientists had been producing hydrogen for years before it was recognized as an element. Written records indicate that Robert Boyle produced hydrogen gas as early as 1671 while experimenting with iron and acids. Hydrogen was first recognized as a distinct element by Henry Cavendish in 1766. Composed of a single proton and a single electron, hydrogen is the simplest and most abundant element in the universe. It is estimated that 90% of the visible universe is composed of hydrogen.

Hydrogen is the raw fuel that most stars 'burn' to produce energy. The same process, known as fusion, is being studied as a possible power source for use on earth. The sun's supply of hydrogen is expected to last another 5 billion years.

From the Greek word hydro (water), and genes (forming). Hydrogen was recognized as a distinct substance by Henry Cavendish in 1776. Diagram of a simple hydrogen atom.

Hydrogen is the most abundant of all elements in the universe. The heavier elements were originally made from hydrogen atoms or from other elements that were originally made from hydrogen atoms.

Historical Atomic Weights

Year	Atomic Weight (uncertainty) [u]	Reference
2009	[1.007 84, 1.008 11]	https://doi.org/10.1351/PAC-REP-10-09-14
1981	1.007 94(7)	https://doi.org/10.1351/pac198355071101
1971	1.0079(1)	https://doi.org/10.1351/pac197230030637
1969	1.0080(3)	https://doi.org/10.1351/pac197021010091
1961	1.007 97(1)	https://doi.org/10.1021/ja00881a001
1940	1.0080	https://doi.org/10.1039/JR9400000475
1938	1.0081	https://doi.org/10.1039/JR9380001101
1931	1.0078	https://doi.org/10.1039/JR9310001617
1903	1.008	https://doi.org/10.1021/ja02003a001
1902	1.01	https://doi.org/10.1007/BF01370337

Historical Isotopic Abundances

Year	Isotope	Abundance (uncertainty)	Reference
2013	¹H	[0.999 72, 0.999 99]	https://doi.org/10.1515/pac-2015-0503
2013	²H	[0.000 01, 0.000 28]	https://doi.org/10.1515/pac-2015-0503
1997	¹H	0.999 885(70)	https://doi.org/10.1351/pac199870010217
1997	²H	0.000 115(70)	https://doi.org/10.1351/pac199870010217
1975	¹H	0.999 85	https://doi.org/10.1351/pac197647010075
1975	²H	0.000 15	https://doi.org/10.1351/pac197647010075

Users

Hydrogen is a commercially important element. Large amounts of hydrogen are combined with nitrogen from the air to produce ammonia (NH3) through a process called the Haber process. Hydrogen is also added to fats and oils, such as peanut oil, through a process called hydrogenation. Liquid hydrogen is used in the study of superconductors and, when combined with liquid oxygen, makes an excellent rocket fuel.

Hydrogen combines with other elements to form numerous compounds. Some of the common ones are: water (H2O), ammonia (NH3), methane (CH4), table sugar (C12H22O11), hydrogen peroxide (H2O2) and hydrochloric acid (HCl).

Hydrogen has three common isotopes. The simplest isotope, called protium, is just ordinary hydrogen. The second, a stable isotope called deuterium, was discovered in 1932. The third isotope, tritium, was discovered in 1934.

Great quantities of hydrogen are required commercially for nitrogen fixation using the Haber ammonia process, and for the hydrogenation of fats and oils. It is also used in large quantities in methanol production, in hydrodealkylation, hydrocracking, and hydrodesulfurization. Other uses include rocket fuel, welding, producing hydrochloric acid, reducing metallic ores, and filling balloons.

The lifting power of 1 cubic foot of hydrogen gas is about 0.07 lb at °C, 760 mm pressure.

The hydrogen fuel cell is a developing technology that will allow great amounts of electrical power to be obtained using a source of hydrogen gas.

Consideration is being given to an entire economy based on solar- and nuclear-generated hydrogen. Public acceptance, high capital investment, and the high cost of hydrogen with respect to today's fuels are but a few of the problems facing such an economy. Located in remote regions, power plants would electrolyze seawater; the hydrogen produced would travel to distant cities by pipelines. Pollution-free hydrogen could replace natural gas, gasoline, etc., and could serve as a reducing agent in metallurgy, chemical processing, refining, etc. It could also be used to convert trash into methane and ethylene.

Sources

Hydrogen is estimated to make up more than 90% of all the atoms three quarters of the mass of the universe! This element is found in the stars, and plays an important part in powering the universe through both the proton-proton reaction and carbon-nitrogen cycle. Stellar hydrogen fusion processes release massive amounts of energy by combining hydrogens to form helium.

Production of hydrogen in the U.S. alone amounts to about 3 billion cubic feet per year. Hydrogen is prepared by

▸ steam on heated carbon,

▸ decomposition of certain hydrocarbons with heat,

▸ reaction of sodium or potassium hydroxide on aluminum

▸ electrolysis of water, or

▸ displacement from acids by certain metals.

Liquid hydrogen is important in cryogenics and in the study of superconductivity, as its melting point is only 20 degrees above absolute zero.

Tritium is readily produced in nuclear reactors and is used in the production of the hydrogen bomb.

Hydrogen is the primary component of Jupiter and the other gas giant planets. At some depth in the planet's interior the pressure is so great that solid molecular hydrogen is converted to solid metallic hydrogen.

In 1973, a group of Russian experimenters may have produced metallic hydrogen at a pressure of 2.8 Mbar. At the transition the density changed from 1.08 to 1.3 g/cm3. Earlier, in 1972, at Livermore, California, a group also reported on a similar experiment in which they observed a pressure-volume point centered at 2 Mbar. Predictions say that metallic hydrogen may be metastable; others have predicted it would be a superconductor at room temperature.

Compounds

Although pure hydrogen is a gas, we find very little of it in our atmosphere. Hydrogen gas is so light that, uncombined, hydrogen will gain enough velocity from collisions with other gases that they will quickly be ejected from the atmosphere. On earth, hydrogen occurs chiefly in combination with oxygen in water, but it is also present in organic matter such as living plants, petroleum, coal, etc. It is present as the free element in the atmosphere, but only less than 1 ppm by volume. The lightest of all gases, hydrogen combines with other elements sometimes explosively to form compounds.

Quite apart from isotopes, it has been shown that under ordinary conditions hydrogen gas is a mixture of two kinds of molecules, known as ortho- and para-hydrogen, which differ from one another by the spins of their electrons and nuclei.

Normal hydrogen at room temperature contains 25% of the para form and 75% of the ortho form. The ortho form cannot be prepared in the pure state. Since the two forms differ in energy, the physical properties also differ. The melting and boiling points of parahydrogen are about 0.1°C lower than those of normal hydrogen.

See more information at the Hydrogen compound page.

Element Forms

CID	Name	Formula	SMILES	Molecular Weight
1038	hydron	H+	[H+]	1.0080
5362549	hydrogen	H	[H]	1.0080
166653	hydride	H-	[H-]	1.0080
5460653	proton	H+	[1H+]	1.0078250319
5460634	deuteron	H+	[2H+]	2.0141017778
5460632	triton	H+	[3H+]	3.0160492813
5460633	tritium	H	[3H]	3.0160492813
5460635	deuterium	H	[2H]	2.0141017778
6857427	tritide	H-	[3H-]	3.0160492813
6857428	deuteride	H-	[2H-]	2.0141017778
6857429	protide	H-	[1H-]	1.0078250319
5460654	protium	H	[1H]	1.0078250319

Isotopes

Stable Isotope Count	2
Summary	The ordinary isotope of hydrogen, H, is known as Protium, the other two isotopes are Deuterium (a proton and a neutron) and Tritium (a protron and two neutrons). Hydrogen is the only element whose isotopes have been given different names. Deuterium and Tritium are both used as fuel in nuclear fusion reactors. One atom of Deuterium is found in about 6000 ordinary hydrogen atoms.

Isotopes in Earth/Planetary Science

Molecules, atoms, and ions of the stable isotopes of hydrogen possess slightly different physical and chemical properties and they are commonly fractionated during physical, chemical, and biological processes, giving rise to variations in isotopic abundances and in atomic weights (Fig. IUPAC.1.1). Hydrogen has the largest relative mass difference among its isotopes and consequently exhibits the largest variation in isotopic composition of any element that does not have radioactive or radiogenic isotopes. Ranges in the stable isotopic composition of naturally occurring hydrogen-bearing materials are shown in Fig. IUPAC.1.1. These variations enable hydrogen isotopes to be used as tracers in environmental studies [13].

A primary use of stable hydrogen isotopes is in isotope hydrology. Although the evolution of the stable hydrogen and oxygen isotopic composition of precipitation begins with the evaporation of water from the oceans, their local and global relationship arises primarily from equilibrium isotopic fractionation of heavier (2H and 18O) and lighter (1H and 16O) isotopes of hydrogen and oxygen during condensation as a tropospheric vapor mass follows a trajectory to higher latitudes and over continents [14], [15]. As a consequence, the hydrogen isotopic composition of precipitation, rivers, and tap waters varies with elevation, season, and distance from the ocean-continent boundary. Figure 4.1.2 shows the variation in the atomic weight of hydrogen in water from rivers across the United States. These variations in the hydrogen isotopic composition of environmental water are often combined with stable oxygen isotopic compositions and have been used to identify the origin of water samples and to investigate the interaction between groundwater and surface water (e.g. lakes, streams, and rivers) [16].

Fig. IUPAC.1.1: Variation in atomic weight with isotopic composition of selected hydrogen-bearing materials (modified from [13], [17]).

Fig. IUPAC.1.2: Variation in atomic weight of hydrogen in river waters across the continental United States (modified from [16]). Blue color indicates waters most depleted in ²H (resulting in lower atomic weight of hydrogen) and brown color indicates those most enriched in ²H (resulting in higher atomic weight of hydrogen).

[13] M. W. Wieser, T. B. Coplen. Pure Appl Chem.83, 359 (2011).

[14] W. Dansgaard. Tellus16, 436 (1964).

[15] I. D. Clark, P. Fritz. Environmental Isotopes in Hydrogeology, p. 328, Lewis Publishers, New York (1997).

[16] C. Kendall, T. B. Coplen. Hydrol. Processes.15, 1363 (2011).

[17] T. B. Coplen, J. A. Hopple, J. K. Böhlke, H. S. Peiser, S. E. Rieder, H. R. Krouse, K. J. R. Rosman, T. Ding, R. D. Vocke, K. Revesz, A. Lamberty, P. D. P. Taylor, P. D. Bièvre. United States Geological Survey Water-Resources Investigations Report, 01-4222, (2002).

Isotopes in Forensic Science and Anthropology

Measurements of relative 2H abundances are used to determine the breeding grounds of many species of migrant songbirds. These species of songbirds only grow their feathers before migration, and they grow them on or close to their breeding grounds. Therefore, the isotopic composition of a bird’s feathers correlates to the isotopic composition of the growing season’s precipitation [18], [19], [20].

Measurements of relative 2H abundances of human hair samples collected at archeological sites are used to determine the geographic region in which a subject lived based on the hydrogen isotopic composition of the water they drank. This is possible because hair stores a daily record of the hydrogen isotopic composition of intake water, which correlates to local meteoric water [18], [21].

[18] Z. D. Sharp, V. Atudorei, H. O. Panarello, J. Fernández, C. Douthitt. J. Archaeolog. Sci.30, 1709 (2003).

[19] K. A. Hobson. Oecologia120, 314 (1999).

[20] K. A. Hobson, L. I. Wassenaar. Oecologia.109, 142 (1996).

[21] T. B. Coplen, H. Qi. Forensic Sci. Int.266, 222 (2016).

Isotopes in Geochronology

3H (tritium), with a half-life of 12.31 years, decays to 3He. The relative variations in n(3He)/n(3H) ratios can be interpreted in terms of elapsed time for dating purposes. The dates of groundwater recharge (water moving downward from the surface), where large amounts of 3H were received from precipitation following thermonuclear bomb test periods, come from the elapsed time since a water mass became isolated from the atmosphere in the time range from the mid-1950s to the present [15].

[15] I. D. Clark, P. Fritz. Environmental Isotopes in Hydrogeology, p. 328, Lewis Publishers, New York (1997).

Isotopes in Industry

3H is used for self-luminous exit signs in aircraft and commercial buildings. It is found in luminous dials, gauges, wristwatches, and luminous paints [22]. 2H, in the form of heavy water, is used in CANDU (CANada Deuterium Uranium) nuclear reactors as a moderator and coolant [23].

[22] United States Nuclear Regulatory Commission. NRC: Fact Sheet on Tritium Exit Signs, (2013), November 12; http://www.nrc.gov/reading-rm/doc-collections/fact-sheets/fs-tritium.html.

[23] CANDU Owners Group Inc. CANDU Reactors, CANDU Owners Group Inc (2013), July 30; http://www.candu.org/candu_reactors.html.

Isotopes in Medicine

2H is used for the isotopic labeling of drugs and nutrients to trace their uptake and metabolism in the human body [24], [25]. 2H, in the form of heavy water, is used to study human metabolism. For example, 2H is used in combination with 18O (double labeled water) to measure energy expenditure [26].

[24] S. P. O’Grady, A. R. Wende, C. H. Remien, L. O. Valenzuela, L. E. Enright, L. A. Chesson, E. D. Abel, T. E. Cerling, J. R. Ehleringer. PLoS One5, e11699 (2010). https://doi.org/10.1371/journal.pone.0011699.

[25] M. van Lieshout, C. E. West, R. B. van Breemen. Am. J. Clin. Nutr.77, 12 (2003).

[26] D. A. Schoeller, E. Ravussin, Y. Schutz, K. J. Acheson, P. Baertschi, E. Jequier. Am. J. Physiol. Regul. Integr. Comp. Physiol.250, R823 (1986).

Isotope Mass and Abundance

Isotope	Atomic Mass (uncertainty) [u]	Abundance (uncertainty)
^1H	1.007 825 0322(6)	[0.999 72, 0.999 99]
^2H	2.014 101 7781(8)	[0.000 01, 0.000 28]

Isotope	Atomic Mass (uncertainty) [u]	Abundance (uncertainty)
¹H	1.00782503223(9)	0.999885(70)
²D	2.01410177812(12)	0.000115(70)
³T	3.0160492779(24)

Atomic Mass, Half Life, and Decay

Nuclide	Atomic Mass and Uncertainty [u]	Half Life and Uncertainty	Discovery Year	Decay Modes, Intensities and Uncertainties [%]
¹H	1.007825031898 ± 0.000000000014	Stable	1920	IS=99.9855±7.8%
²H	2.014101777844 ± 0.000000000015	Stable	1932	IS=0.0145±7.8%
³H	3.01604928132 ± 0.00000000008	12.32 y ± 0.02	1934	β^-=100%
⁴H	4.026431867 ± 0.000107354	139 ys ± 10	1981	n=100%
⁵H	5.035311492 ± 0.00009602	86 ys ± 6	1987	2n=100%
⁶H	6.044955437 ± 0.000272816	294 ys ± 67	1984	n ?; 3n ?
⁷H	7.052749 ± 0.001078 [Estimated]	652 ys ± 558	2003	2n ?

Information Sources

1. PubChem
https://pubchem.ncbi.nlm.nih.gov
2. Atomic Mass Data Center (AMDC), International Atomic Energy Agency (IAEA)
LICENSE
https://www-nds.iaea.org/disclaimer.html

H
https://www-nds.iaea.org/amdc/
3. IUPAC Commission on Isotopic Abundances and Atomic Weights (CIAAW)
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http://www.ciaaw.org/bylaws.htm

Hydrogen
https://www.ciaaw.org/hydrogen.htm
4. Jefferson Lab, U.S. Department of Energy
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Please see citation and linking information https https://www.jlab.org/privacy-and-security-notice

Hydrogen
https://education.jlab.org/itselemental/ele001.html
5. Los Alamos National Laboratory, U.S. Department of Energy
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https://about.lanl.gov/lanl-resources/web-policies/

Hydrogen
https://periodic.lanl.gov/1.shtml
6. NIST Physical Measurement Laboratory
LICENSE
https://www.nist.gov/pml/database-disclaimer

Hydrogen
https://physics.nist.gov/cgi-bin/Elements/elInfo.pl?element=1
7. IUPAC Periodic Table of the Elements and Isotopes (IPTEI)
LICENSE
Copyright (c) 2020 International Union of Pure and Applied Chemistry. The International Union of Pure and Applied Chemistry (IUPAC) contribution within Pubchem is provided under a CC-BY-NC-ND 4.0 license, unless otherwise stated.
https://creativecommons.org/licenses/by-nc-nd/4.0/
https://doi.org/10.1515/pac-2015-0703
8. PubChem Elements
LICENSE
https://www.nlm.nih.gov/copyright.html

Hydrogen

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