Hydrogen
Appearance colorless gas
Standard atomic weight Ar, std(H) [1.00784, 1.00811] conventional: 1.008[1]
Atomic number (Z) 1
Group group 1: hydrogen and alkali metals
Period period 1
Block s-block
Electron configuration 1s1
Electrons per shell 1
Physical properties
Phase at STP gas
Melting point (H2) 13.99 K (−259.16 °C, −434.49 °F)
Boiling point (H2) 20.271 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 point 13.8033 K, 7.041 kPa
Critical point 32.938 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)
Electronegativity Pauling scale: 2.20
Ionization energies
1st: 1312.0 kJ/mol
Covalent radius 31±5 pm
Van der Waals radius 120 pm
Color lines in a spectral range
Spectral lines of hydrogen
Other properties
Natural occurrence primordial
Crystal structure hexagonalHexagonal crystal structure for hydrogen
Speed of sound 1310 m/s (gas, 27 °C)
Thermal conductivity 0.1805 W/(m⋅K)
Magnetic ordering diamagnetic
Molar magnetic susceptibility −3.98×10−6 cm3/mol (298 K)[4]
CAS Number 12385-13-6
1333-74-0 (H2)
History
Discovery Henry Cavendish (1766)
Named by Antoine Lavoisier (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
Hydrogen is the chemical element with the symbol H and atomic number 1. Hydrogen is the lightest element. At standard conditions hydrogen is a gas of diatomic molecules having the formula H2. It is colorless, odorless, tasteless, non-toxic, and highly combustible. Hydrogen is the most abundant chemical substance in the universe, constituting roughly 75% of all normal matter.] Stars such as the Sun are mainly composed of hydrogen in the plasma state. Most of the hydrogen on Earth exists in molecular forms such as water and organic compounds. For the most common isotope of hydrogen (symbol 1H) each atom has one proton, one electron, and no neutrons.
In the early universe, the formation of protons, the nuclei of hydrogen, occurred during the first second after the Big Bang. The emergence of neutral hydrogen atoms throughout the universe occurred about 370,000 years later during the recombination epoch, when the plasma had cooled enough for electrons to remain bound to protons.
The Space Shuttle Main Engine burnt hydrogen with oxygen, producing a nearly invisible flame at full thrust.
Hydrogen is nonmetallic, except at extremely high pressures, and readily forms a single covalent bond with most nonmetallic elements, forming compounds such as water and nearly all organic compounds. Hydrogen plays a particularly important role in acid–base reactions because these reactions usually involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a negative charge (i.e., anion) where it is known as a hydride, or as a positively charged (i.e., cation) species denoted by the symbol H+. The H+ cation is simply a proton (symbol p) but its behavior in aqueous solutions and in ionic compounds involves screening of its electric charge by nearby polar molecules or anions. Because hydrogen is the only neutral atom for which the Schrödinger equation can be solved analytically, the study of its energetics and chemical bonding has played a key role in the development of quantum mechanics.
Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. In 1766–1781, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, and that it produces water when burned, the property for which it was later named: in Greek, hydrogen means "water-former".
Industrial production is mainly from steam reforming natural gas, and less often from more energy-intensive methods such as the electrolysis of water.Most hydrogen is used near the site of its production, the two largest uses being fossil fuel processing (e.g., hydrocracking) and ammonia production, mostly for the fertilizer market. Hydrogen is problematic in metallurgy because it can embrittle many metals,complicating the design of pipelines and storage tanks.
Combustion
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.
A black cup-like object hanging by its bottom with blue glow coming out of its opening.
The Space Shuttle Main Engine burnt hydrogen with oxygen, producing a nearly invisible flame at full thrust.
Hydrogen gas (dihydrogen or molecular hydrogen)is highly flammable:
The enthalpy of combustion is −286 kJ/mol.
Hydrogen gas forms explosive mixtures with air in concentrations from 4–74% and with chlorine at 5–95%. The explosive reactions may be triggered by spark, heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C (932 °F).
Flame
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, which uses an ammonium perchlorate composite. 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. The destruction of the Hindenburg airship was a notorious example of hydrogen combustion and the cause is still debated. The visible flames in the photographs were the result of carbon compounds in the airship skin burning.
Reactants
H2 is unreactive compared to diatomic elements such as halogens or oxygen. The thermodynamic basis of this low reactivity is the very strong H-H bond, with a bond dissociation energy of 435.7 kJ/mol. The kinetic basis of the low reactivity is the nonpolar nature of H2 and its weak polarizability. It spontaneously reacts with chlorine and fluorine to form hydrogen chloride and hydrogen fluoride, respectively. The reactivity of H2 is strongly affected by the presence of metal catalysts. Thus, while mixtures of H2 with O2 or air combust readily when heated to at least 500 C by a spark or flame, they do not react at room temperature in the absence of a catalyst.
Hydrogen Compounds
Hydrogen compounds are the most important oxidants for many chemicals in the atmosphere and are involved in the cycles of many chemical families. Hydrogen compounds include atomic hydrogen (H), which is very short-lived because it combines quickly with O2 to form the hydroperoxyl radical (HO2); molecular hydrogen (H2), which, next to CH4, is the most abundant reactive trace gas in the troposphere; the hydroxyl radical OH; the hydroperoxyl radical HO2; hydrogen peroxide (H2O2), which is formed by the reaction of HO2 radicals and is an important oxidant for SO2 in cloud droplets; and H2O, which plays multivarious roles in atmospheric chemistry, including its reaction with excited atomic oxygen to form OH.The primary source of odd hydrogen species (HOx), in the form of OH, in the lower troposphere where water vapor is abundant. Another source of HOx is the photolysis of formaldehyde followed by and both of which produce HO2. The simplest loss mechanisms for HOx are of the form, and HOx and NOx can react to produce O3. Consequently, there is considerable interest in HOx chemistry in the upper troposphere where NOx is emitted in significant quantities by jet aircraft. Cycling between OH and HO2 occurs on a timescale of a few seconds and is controlled by reactions. These reactions account for ∼80% of the measured concentrations of HOx in the upper troposphere. NOx controls the cycling within HOx that leads to O3 production. NOx also regulates the loss of HOx through reactions of OH with HO2, NO2, and HO2NO2.
