The name hydride is used for the negative ion of hydrogen, H−, and for compounds of hydrogen with other elements. Every element of the periodic table (except some noble gases) forms one or more compounds with hydrogen, and these compounds (especially those with elements in groups 1–15 of the periodic table) may be referred to as hydrides. Thus, the term "hydride" can be used very broadly. These compounds may be classified into three main types: saline (ionic) hydrides, covalent hydrides, and interstitial hydrides. The hydrides are called binary if they involve only two elements including hydrogen.
Hydrides are useful for a wide diversity of applications. For instance, sodium hydride is a strong base used in organic chemistry; lithium aluminum hydride and sodium borohydride are reducing agents in chemical reactions; nickel hydride is found in nickel metal hydride batteries; and hydrides with arsenic and antimony (arsine and stibine) are used in the semiconductor industry. In addition, silane is used for the manufacture of composite materials, and diborane is a rocket fuel, semiconductor dopant, and reducing agent. Also, various metal hydrides are being studied for possible hydrogen storage in fuel cell-powered electric cars and batteries.
Aside from electrides, the hydride ion is the simplest possible anion, consisting of two electrons and a proton. However, the free hydride ion is so unstable that it exists only under exceptional conditions.
Hydrogen has a relatively low electron affinity, 72.77 kJ/mol, thus hydride is so basic that it is unknown in solution. The reactivity of the hypothetical hydride ion is dominated by its exothermic protonation to give dihydrogen:
- H− + H+ → H2; ΔH = −1675 kJ/mol
As a result, the hydride ion is one of the strongest bases known. It would extract protons from almost any hydrogen-containing species. The low electron affinity of hydrogen and the strength of the H–H bond (436 kJ/mol) means that the hydride ion would also be a strong reducing agent:
- H2 + 2e− ⇌ 2H−; E
o= −2.25 V
- H2 + 2e− ⇌ 2H−; E
Compounds known as "hydrides"
The compounds known as "hydrides" are classified according to the predominant nature of their bonding:
- Saline (ionic) hydrides, which have significant ionic character;
- Covalent hydrides, which include the hydrocarbons and many other compounds; and
- Interstitial hydrides, which may be described as having metallic bonding.
Saline (ionic) hydrides
Saline (or ionic) hydrides are ionic compounds, and therefore salt-like. They are solids with high melting points. In these cases, hydrogen is in the form of the anion (H−), which is combined with a highly electropositive element, usually one of the alkali metals or some of the alkaline earth metals (calcium, strontium, barium). Examples are sodium hydride (NaH) and calcium hydride (CaH2).
In each ionic hydride, the hydrogen atom behaves as a halogen atom, obtaining an electron from the metal atom to form a hydride ion (H−). The hydrogen atom thereby fills its 1s-orbital and attaining the stable electron configuration of helium.
Ionic hydrides are commonly encountered as basic reagents in organic synthesis:
- C6H5C(O)CH3 + KH → C6H5C(O)CH2K + H2
Such reactions are heterogeneous, the KH does not dissolve. Typical solvents for such reactions are ethers.
Alkali metal hydrides react with metal halides. For example, lithium aluminum hydride (often abbreviated as LAH) arises from reactions with aluminum chloride.
- 4 LiH + AlCl3 → LiAlH4 + 3 LiCl
In covalent hydrides, hydrogen is covalently bonded to an element in the p-block of the periodic table (boron, aluminum, and elements in groups 14-17), as well as beryllium. The hydrocarbons and ammonia could be considered hydrides of carbon and nitrogen, respectively.
Charge-neutral covalent hydrides that are made up of small molecules are often volatile at room temperature and atmospheric pressure. Some covalent hydrides are not volatile because they are polymeric (i.e., nonmolecular), such as the binary hydrides of aluminum and beryllium. Replacing some hydrogen atoms in such compounds with larger ligands, one obtains molecular derivatives. For example, diisobutylaluminum hydride (DIBAL) consists of two aluminum centers bridged by hydride ligands.
Interstitial hydrides of transition metals
Transition metals form binary hydrides in which hydrogen atoms are bonded to the metal atoms, but the exact nature of those bonds is not clear. In addition, the ratio of hydrogen atoms to metal atoms in a number of these hydrides is not fixed. The lattice of metal atoms contains a variable number of hydrogen atoms that can migrate through it. In materials engineering, the phenomenon of hydrogen embrittlement is a consequence of interstitial hydrides.
For example, palladium absorbs up to 900 times its own volume of hydrogen at room temperature, forming palladium hydride, which was once thought of as a means to carry hydrogen for vehicular fuel cells. Hydrogen gas is liberated proportional to the applied temperature and pressure but not to the chemical composition.
Interstitial hydrides show some promise as a way for safe hydrogen storage. During the last 25 years, many interstitial hydrides were developed that readily absorb and discharge hydrogen at room temperature and atmospheric pressure. They are usually based on intermetallic compounds and solid-solution alloys. However, their application is still limited, as they are capable of storing only about 2 percent (by weight) of hydrogen, which is not enough for automotive applications.
Transition metal hydride (or hydrido) complexes
Most transition metal complexes form molecular compounds described as hydrides. Usually, such compounds are discussed in the context of organometallic chemistry. Transition metal hydrides are intermediates in many industrial processes that rely on metal catalysts, such as hydroformylation, hydrogenation, and hydrodesulfurization. Two famous examples, HCo(CO)4 and H2Fe(CO)4, are acidic, thus demonstrating that the term hydride is used very broadly.
When a dihydrogen complex loses a proton, a metal hydride is produced. The anion [ReH9]2- (nonahydridorhenate) is an example of a molecular metal hydride.
The following list gives the nomenclature for hydrides of main group elements:
- alkali and alkaline earth metals: metal hydride
- boron: borane and rest of the group as metal hydride
- carbon: alkanes, alkenes, alkynes, and all hydrocarbons
- silicon: silane
- germanium: germane
- tin: stannane
- lead: plumbane
- nitrogen: ammonia ('azane' when substituted), hydrazine
- phosphorus: phosphine ('phosphane' when substituted)
- arsenic: arsine ('arsane' when substituted)
- antimony: stibine ('stibane' when substituted)
- bismuth: bismuthine ('bismuthane' when substituted)
According to the convention used above, the following elements form "hydrogen compounds" and not "hydrides":
- oxygen: water ('oxidane' when substituted), hydrogen peroxide
- sulfur: hydrogen sulfide ('sulfane' when substituted)
- selenium: hydrogen selenide ('selane' when substituted)
- tellurium: hydrogen telluride ('tellane' when substituted)
- halogens: hydrogen halides
Isotopes of hydride
According to IUPAC convention, by precedence (stylized electronegativity), hydrogen falls between group 15 and group 16 elements. Therefore we have NH3, 'nitrogen hydride' (ammonia), versus H2O, 'hydrogen oxide' (water).
Various metal hydrides are currently being studied for use as a means of hydrogen storage in fuel cell-powered electric cars and batteries. They also have important uses in organic chemistry as powerful reducing agents, and many promising uses in the proposed hydrogen economy.
The names and uses of some specific hydrides are given below:
- nickel hydride: used in NiMH batteries
- palladium hydride: catalyst in organic reactions; electrodes in cold fusion experiments
- lithium aluminum hydride: a powerful reducing agent used in organic chemistry
- sodium borohydride: selective specialty reducing agent, hydrogen storage in fuel cells
- sodium hydride: a powerful base used in organic chemistry
- diborane: reducing agent, rocket fuel, semiconductor dopant, catalyst, used in organic synthesis; also borane, pentaborane and decaborane
- arsine: used for doping semiconductors
- stibine: used in semiconductor industry
- phosphine: used for fumigation
- silane: many industrial uses, e.g. manufacture of composite materials and water repellents
- ammonia: coolant, fertilizer, many other industrial uses
ReferencesISBN links support NWE through referral fees
- Brown Jr., Theodore L., H. Eugene LeMay, Bruce Edward Bursten, and Julia R. Burdge. 2002. Chemistry: The Central Science. 9th ed. Upper Saddle River, NJ: Prentice Hall. ISBN 0130669970.
- Chang, Raymond. 2006. Chemistry, 9th ed. New York: McGraw-Hill Science/Engineering/Math. ISBN 0073221031.
- Cotton, F. Albert, and Geoffrey Wilkinson. 1980. Advanced Inorganic Chemistry, 4th ed. New York: Wiley. ISBN 0471027758.
- McMurry, J., and R.C. Fay. 2004. Chemistry, 4th ed. Upper Saddle River, NJ: Prentice Hall. ISBN 0131402080.
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