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Hydrogen

Hydrogen is a chemical element at the start of the periodic table. It has the symbol H and atomic number 1. It also has a standard atomic weight of 1.008, which makes it the lightest element in the periodic table.

Hydrogen,  1H
Template:Infobox element/symbol-to-top-image/alt
Purple glow in its plasma state
General properties
Appearancecolorless gas
Standard atomic weight (Ar, standard)[1.00784, 1.00811][1]
Hydrogen in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
–
↑
H
↓
Li
– ← hydrogen → helium
Atomic number (Z)1
Groupgroup 1
Periodperiod 1
Blocks-block
Element category  reactive nonmetal
Electron configuration1s1
Electrons per shell
1
Physical properties
Phase at STPH: Gas
Melting point13.99 K ​(−259.16 °C, ​−434.49 °F)
Boiling point20.271 K ​(−252.879 °C, ​−423.182 °F)
Density (at STP)0.08988 g/L
when liquid (at m.p.)0.07 g/cm3 (solid: 0.0763 g/cm3)[2]
when liquid (at b.p.)0.07099 g/cm3
Triple point13.8033 K, ​7.041 kPa
Critical point32.938 K, 1.2858 MPa
Heat of fusion(H2) 0.117 kJ/mol
Heat of vaporization(H2) 0.904 kJ/mol
Molar heat capacity(H2) 28.836 J/(mol·K)
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)
ElectronegativityPauling scale: 2.20
Ionization energies
  • 1st: 1312.0 kJ/mol
Covalent radius31±5 pm
Van der Waals radius120 pm
Color lines in a spectral range
Spectral lines of hydrogen
Other properties
Natural occurrenceH: Primordial
Crystal structure ​hexagonal
Hexagonal crystal structure for hydrogen
Speed of sound1310 m/s (gas, 27 °C)
Thermal conductivity0.1805 W/(m·K)
Magnetic orderingdiamagnetic[3]
Magnetic susceptibility−3.98·10−6 cm3/mol (298 K)[4]
CAS Number12385-13-6
1333-74-0 (H2)
History
DiscoveryHenry Cavendish[5][6] (1766)
Named byAntoine Lavoisier[7] (1783)
Main isotopes of hydrogen
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct
1H 99.98% stable
2H 0.02% stable
3H trace 12.32 y β− 3He
| references

Hydrogen is the most common chemical element in the Universe, making up 75% of all normal (baryonic) matter (by mass). Most stars are made of mostly hydrogen. Hydrogen's most common isotope has one proton with one electron orbiting around it.

Properties

Hydrogen is grouped as a reactive nonmetal, unlike the other elements found in the first column of the periodic table, which are grouped as alkali metals. The solid form of hydrogen should behave like a metal, though.

When by itself, hydrogen normally binds with itself to make dihydrogen (H2) which is very stable, because of its high bond-dissociation energy of 435.7 kJ/mol.[8] At normal temperature and pressure, this hydrogen gas (H2) has no color, smell, or taste and is not poisonous as it is a nonmetal and burns very easily.

Combustion

Molecular hydrogen is flammable and reacts with oxygen:

2 H2(g) + O2(g) → 2 H2O(l) + 572 kJ (286 kJ/mol)

At temperatures higher than 500 °C, hydrogen suddenly burns in air.

Compounds

While hydrogen gas in its natural form is not reactive, it does form compounds with many elements, especially halogens, which are very electronegative, meaning they want an electron very badly. Hydrogen also forms massive arrays with carbon atoms, forming hydrocarbons. The study of the properties of hydrocarbons is known as organic chemistry.

The H- anion (negatively charged atom) is named a hydride, though the word is not commonly used. An example of a hydride is lithium hydride (LiH), which is used as a "spark plug" in nuclear weapons.

Acids

Acids dissolved in water normally contain high levels of hydrogen ions, in other words, free protons. Their level is generally used to determine its pH, that is, the content of hydrogen ions in a volume. For example, hydrochloric acid, found in people's stomachs, can dissociate into a chloride anion and a free proton, and the property of the free proton is how it can digest food by corroding it.

Though uncommon on Earth, the H3+ cation is one of the most common ions in the universe.

Isotopes

Main page: Isotopes of hydrogen

Hydrogen has 7 known isotopes, two of which are stable (1H and 2H), which are commonly named protium and deuterium. The isotope 3H is known as tritium, has a half-life of 12.33 years, and is produced in small amounts by cosmic rays. The 4 isotopes left have half-lives on the scale of yoctoseconds.

Hydrogen in nature

In its natural form on Earth, hydrogen is generally a gas. Hydrogen is also one of the parts that make up a water molecule. Hydrogen is important because it is the fuel that powers the Sun and other stars. Hydrogen makes up about 74% of the complete universe.[9]

Natural hydrogen is normally made of two hydrogen atoms connected together. Scientists name these diatomic molecules. Hydrogen will have a chemical reaction when mixed with most other elements, though it has no color or smell.

Natural hydrogen is very uncommon in the Earth's atmosphere, because nearly all hydrogen would have escaped into space because of its weight. In nature, it is generally in water. Hydrogen is also in all living things, as a part of the organic compounds that living things are made of. In addition, hydrogen atoms can join with carbon atoms to form hydrocarbons. Petroleum and other fossil fuels are made of these hydrocarbons and commonly used to make energy.

Some other facts about hydrogen:

History of Hydrogen

Hydrogen was first separated in 1671 by Robert Boyle.[12] In 1776, Henry Cavendish identified it as its own element and named it "inflammable air". He saw in 1781 that burning it made water.[13][5][6]

Hydrogen Spectrum Test

Antoine Lavoisier gave Hydrogen its name, from the Greek word for water, 'υδορ (read /HEEW-dor/) and gennen meaning to "produce"[14] as it forms water in a chemical reaction with oxygen.[15]

Uses of Hydrogen

The most common uses are in the petroleum industry and in making ammonia by the Haber process. Some is used in other places in the chemical industry. A little of it is used as fuel, for example in rockets for spacecraft. Most of the hydrogen that people use comes from a chemical reaction between natural gas and steam.

Nuclear fusion

Nuclear fusion is a very powerful source of energy. It depends on forcing atoms together to make helium and energy, as in a star like the Sun, or in a hydrogen bomb. This needs a large amount of energy to get started, and is not easy to do currently. A big advantage over nuclear fission, which is used in today's nuclear power stations, is that it makes less nuclear waste and does not use a poisonous and uncommon fuel like uranium. More than 600 million tons of hydrogen undergo fusion every second on the Sun.[16][17]

Using hydrogen

Hydrogen is mostly used in the petroleum industry, to change heavy petroleum parts into lighter, more useful ones. It is also used to make ammonia. Smaller amounts are burned as fuel. Most hydrogen is made by a reaction between natural gas and steam.

The electrolysis of water breaks water into hydrogen and oxygen, using electricity. Burning hydrogen joins with oxygen molecules to make steam (natural water vapor). A fuel cell joins hydrogen with an oxygen molecule, releasing an electron as electricity. For these reasons, many people believe hydrogen power will replace other synthetic fuels in the future.

Hydrogen can also be burned to make heat for steam turbines or internal combustion engines. Like other synthetic fuels, hydrogen can be made from natural fuels such as coal or natural gas, or from electricity, and therefore represents a valuable addition to the power grid; in the same role as natural gas. Such a grid and infrastructure with fuel cell vehicles is now planned by a number of countries, such as Japan, Korea and many European countries. This lets these countries buy less petroleum, which is an economic advantage. The other advantage is that, used in a fuel cell or burned in a combustion engine as in a hydrogen car, the engine does not make pollution. Only water, and a small amount of nitrogen oxides, forms.

Hydrogen Media

  • Combustion of hydrogen with the oxygen in the air. When the bottom cap is removed, allowing air to enter at the bottom, the hydrogen in the container rises out of top and burns as it mixes with the air.

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

  • 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)

  • Hydrogen gas is colorless and transparent, here contained in a glass ampoule.

  • Phase diagram of hydrogen. The temperature and pressure scales are logarithmic, so one unit corresponds to a 10× change. The left edge corresponds to 105 Pa, which is about atmospheric pressure.Template:Image reference needed

  • A sample of sodium hydride

  • Hydrogen Isotopes. See a related animation of this medical topic.

  • Hydrogen discharge (spectrum) tube

  • Deuterium discharge (spectrum) tube

  • Antoine-Laurent de Lavoisier

References

  1. ↑ Meija, J.; Coplen, T. B.; Berglund, M.; Brand, W.A.; De Bièvre, P.; Gröning, M.; Holden, N.E.; Irrgeher, J.; Loss, R.D.; Walczyk, T.; Prohaska, T. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure and Applied Chemistry. 88 (3): 265–91. doi:10.1515/pac-2015-0305. {{cite journal}}: Unknown parameter |displayauthors= ignored (|display-authors= suggested) (help)
  2. ↑ Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. p. 240. ISBN 978-0123526519.
  3. ↑ Lide, D. R., ed. (2005). "Magnetic susceptibility of the elements and inorganic compounds". CRC Handbook of Chemistry and Physics (PDF) (86th ed.). Boca Raton (FL): CRC Press. ISBN 978-0-8493-0486-6.
  4. ↑ Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 978-0-8493-0464-4.
  5. ↑ 5.0 5.1 "Hydrogen". Van Nostrand's Encyclopedia of Chemistry. (2005). Wylie-Interscience. 797–799. 
  6. ↑ 6.0 6.1 Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 183–191. ISBN 978-0-19-850341-5.
  7. ↑ Stwertka, Albert (1996). A Guide to the Elements. Oxford University Press. pp. 16–21. ISBN 978-0-19-508083-4.
  8. ↑ Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, FL: CRC Press. ISBN 0-8493-0487-3.
  9. ↑ Cain, Fraser, Universe Today (November 7, 2016). "When was the first light in the universe?". Phys.org. Archived from the original on 18 December 2016. Retrieved 29 November 2016.
  10. ↑ "EIA.doe.gov - What is Hydrogen?". Archived from the original on 2009-02-05. Retrieved 2008-12-22.
  11. ↑ "The magic of syngas". chemrec.se. 2012. Archived from the original on 20 April 2012. Retrieved 7 March 2012.
  12. ↑ Boyle, R. (1672). "Tracts written by the Honourable Robert Boyle containing new experiments, touching the relation betwixt flame and air..." London.
  13. ↑ Boyle, R. (1672). "Tracts written by the Honourable Robert Boyle containing new experiments, touching the relation betwixt flame and air..." London.
  14. ↑ Stwertka, Albert (1996). A Guide to the Elements. Oxford University Press. pp. 16–21. ISBN 978-0-19-508083-4.
  15. ↑ Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 183–191. ISBN 978-0-19-850341-5.
  16. ↑ "What is Fusion?". iter.org. ITER Organization. 2012. Archived from the original on 6 March 2012. Retrieved 7 March 2012.
  17. ↑ "NASA's Cosmicopia". NASA. Archived from the original on 6 November 2011. Retrieved 28 February 2013.

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