Binary compounds of hydrogen are binary chemical compounds containing just hydrogen and one other chemical element. By convention all binary hydrogen compounds are called hydrides even when the hydrogen atom in it is not an anion.[1][2][3][4] These hydrogen compounds can be grouped into several types.

Overview

Binary hydrogen compounds in group 1 are the ionic hydrides (also called saline hydrides) wherein hydrogen is bound electrostatically. Because hydrogen is located somewhat centrally in an electronegative sense, it is necessary for the counterion to be exceptionally electropositive for the hydride to possibly be accurately described as truly behaving ionic. Therefore, this category of hydrides contains only a few members.

Hydrides in group 2 are polymeric covalent hydrides. In these, hydrogen forms bridging covalent bonds, usually possessing mediocre degrees of ionic character, which make them difficult to be accurately described as either covalent or ionic. The one exception is beryllium hydride, which has definitively covalent properties.

Hydrides in the transition metals and lanthanides are also typically polymeric covalent hydrides. However, they usually possess only weak degrees of ionic character. Usually, these hydrides rapidly decompose into their component elements at ambient conditions. The results consist of metallic matrices with dissolved, often stoichiometric or near so, concentrations of hydrogen, ranging from negligible to substantial. Such a solid can be thought of as a solid solution and is alternately termed a metallic- or interstitial hydride. These decomposed solids are identifiable by their ability to conduct electricity and their magnetic properties (the presence of hydrogen is coupled with the delocalisation of the valence electrons of the metal), and their lowered density compared to the metal. Both the saline hydrides and the polymeric covalent hydrides typically react strongly with water and air.

It is possible to produce a metallic hydride without requiring decomposition as a necessary step. If a sample of bulk metal is subjected to any one of numerous hydrogen absorption techniques, the characteristics, such as luster and hardness of the metal is often retained to a large degree. Bulk actinoid hydrides are only known in this form. The affinity for hydrogen for most of the d-block elements are low. Therefore, elements in this block do not form hydrides (the hydride gap) under standard temperature and pressure with the notable exception of palladium.[5] Palladium can absorb up to 900 times its own volume of hydrogen and is therefore actively researched in the field hydrogen storage.

Elements in group 13 to 17 (p-block) form covalent hydrides (or nonmetal hydrides). In group 12 zinc hydride is a common chemical reagent but cadmium hydride and mercury hydride are very unstable and esoteric. In group 13 boron hydrides exist as a highly reactive monomer BH3, as an adduct for example ammonia borane or as dimeric diborane and as a whole group of BH cluster compounds. Alane (AlH3) is a polymer. Gallium exists as the dimer digallane. Indium hydride is only stable below −90 °C (−130 °F). Not much is known about the final group 13 hydride, thallium hydride.

Due to the total number of possible binary saturated compounds with carbon of the type CnH2n+2 being very large, there are many group 14 hydrides. Going down the group the number of binary silicon compounds (silanes) is small (straight or branched but rarely cyclic) for example disilane and trisilane. For germanium only 5 linear chain binary compounds are known as gases or volatile liquids. Examples are n-pentagermane, isopentagermane and neopentagermane. Of tin only the distannane is known. Plumbane is an unstable gas.

The hydrogen halides, hydrogen chalcogenides and pnictogen hydrides also form compounds with hydrogen, whose lightest members show many anomalous properties due to hydrogen bonding.

Non-classical hydrides are those in which extra hydrogen molecules are coordinated as a ligand on the central atoms. These are very unstable but some have been shown to exist.

Polyhydrides or superhydrides are compounds in which the number of hydrogen atoms exceed the valency of the combining atom. These may only be stable under extreme pressure, but may be high temperature superconductors, such as H3S, superconducting at up to 203 K. Polyhydrides are actively studied with the hope of discovering a room temperature superconductor.

The periodic table of the stable binary hydrides

The relative stability of binary hydrogen compounds and alloys at standard temperature and pressure can be inferred from their standard enthalpy of formation values.[6]

H2 0 He
LiH −91 BeH2 negative BH3 41 CH4 −74.8 NH3 −46.8 H2O −243 HF −272 Ne
NaH −57 MgH2 −75 AlH3 −46 SiH4 31 PH3 5.4 H2S −20.7 HCl −93 Ar
KH −58 CaH2 −174 ScH2 TiH2 VH CrH Mn FeH, FeH2 Co Ni CuH ZnH2 GaH3 GeH4 92 AsH3 67 H2Se 30 HBr −36.5 Kr
RbH −47 SrH2 −177 YH2 ZrH2 NbH Mo Tc Ru Rh PdH Ag CdH2 InH3 SnH4 163 SbH3 146 H2Te 100 HI 26.6 Xe
CsH −50 BaH2 −172 LuH2 HfH2 TaH W Re Os Ir Pt Au Hg Tl PbH4 252 BiH3 247 H2Po 167 HAt positive Rn
Fr Ra Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
↓
LaH2 CeH2 PrH2 NdH2 PmH2 SmH2 EuH2 GdH2 TbH2 DyH2 HoH2 ErH2 TmH2 YbH2
Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No
Binary compounds of hydrogen
Covalent hydrides metallic hydrides
Ionic hydrides Intermediate hydrides
Do not exist Not assessed

Molecular hydrides

The isolation of monomeric molecular hydrides usually require extremely mild conditions, which are partial pressure and cryogenic temperature. The reason for this is threefold - firstly, most molecular hydrides are thermodynamically unstable toward decomposition into their elements; secondly, many molecular hydrides are also thermodynamically unstable toward polymerisation; and thirdly, most molecular hydrides are also kinetically unstable toward these types of reactions due to low activation energy barriers.

Instability toward decomposition is generally attributable to poor contribution from the orbitals of the heavier elements to the molecular bonding orbitals. Instability toward polymerisation is a consequence of the electron-deficiency of the monomers relative to the polymers. Relativistic effects play an important role in determining the energy levels of molecular orbitals formed by the heavier elements. As a consequence, these molecular hydrides are commonly less electron-deficient than otherwise expected. For example, based on its position in the 12th column of the periodic table alone, mercury(II) hydride would be expected to be rather deficient. However, it is in fact satiated, with the monomeric form being much more energetically favourable than any oligomeric form.

The table below shows the monomeric hydride for each element that is closest to, but not surpassing its heuristic valence. A heuristic valence is the valence of an element that strictly obeys the octet, duodectet, and sexdectet valence rules. Elements may be prevented from reaching their heuristic valence by various steric and electronic effects. In the case of chromium, for example, stearic hindrance ensures that both the octahedral and trigonal prismatic molecular geometries for CrH
6 are thermodynamically unstable to rearranging to a Kubas complex structural isomer.

Where available, both the enthalpy of formation for each monomer and the enthalpy of formation for the hydride in its standard state is shown (in brackets) to give a rough indication of which monomers tend to undergo aggregation to lower enthalpic states. For example, monomeric lithium hydride has an enthalpy of formation of 139 kJ mol−1, whereas solid lithium hydride has an enthalpy of −91 kJ mol−1. This means that it is energetically favourable for a mole of monomeric LiH to aggregate into the ionic solid, losing 230 kJ as a consequence. Aggregation can occur as a chemical association, such as polymerisation, or it can occur as an electrostatic association, such as the formation of hydrogen-bonding in water.

Classical hydrides

Classical hydrides
1 2 3 4 5 6 5 4 3 2 1 2 3 4 3 2 1
H
2 0
LiH[7] 139
(−91) 		BeH
2[8] 123 		BH
3[9] 107
(41) 		CH
4 −75 		NH
3 −46 		H
2O −242
(−286) 		HF −273
NaH[10] 140
(−56) 		MgH
2 142
(−76) 		AlH
3[11] 123
(−46) 		SiH
4 34 		PH
3 5 		H
2S −21 		HCl −92
KH 132
(−58) 		CaH
2 192
(−174) 		ScH
3 		TiH
4 		VH
2[12] 		CrH
2[13] 		MnH
2[14] (−12) 		FeH
2[15] 324 		CoH
2[16] 		NiH
2[17] 168 		CuH[18] 278
(28) 		ZnH
2[19] 162 		GaH
3[20] 151 		GeH
4 92 		AsH
3 67 		H
2Se 30 		HBr −36
RbH 132
(−47) 		SrH
2 201
(−177) 		YH
3 		ZrH
4 		NbH
4[12] 		MoH
6[21] 		Tc RuH
2[15] 		RhH
2[22] 		PdH[23] 361 AgH[18] 288 CdH
2[19] 183 		InH
3[24] 222 		SnH
4 163 		SbH
3 146 		H
2Te 100 		HI 27
CsH 119
(−50) 		BaH
2 213
(−177) 		LuH
3 		HfH
4 		TaH
4[12] 		WH
6[25] 586 		ReH
4[14] 		Os Ir PtH
2[26] 		AuH[18] 295 HgH
2[27] 101 		TlH
3[28] 293 		PbH
4 252 		BiH
3 247 		H
2Po 167 		HAt 88
Fr Ra Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts
↓
3 4 5 6 7 8 7 6 5 4 3 2 1 2
LaH
3 		CeH
4 		PrH
3 		NdH
4 		Pm SmH
4 		EuH
3[29] 		GdH
3 		TbH
3 		DyH
4 		HoH
3 		ErH
2 		TmH YbH
2
Ac ThH
4[30] 		Pa UH
4[31] 		Np Pu Am Cm Bk Cf Es Fm Md No
Legend
Monomeric covalent Oligomeric covalent
Polymeric covalent Ionic
Polymeric delocalised covalent
Unknown solid structure Not assessed

This table includes the thermally unstable dihydrogen complexes for the sake of completeness. As with the above table, only the complexes with the most complete valence is shown, to the negligence of the most stable complex.

Non-classical covalent hydrides

A molecular hydride may be able to bind to hydrogen molecules acting as a ligand. The complexes are termed non-classical covalent hydrides. These complexes contain more hydrogen than the classical covalent hydrides, but are only stable at very low temperatures. They may be isolated in inert gas matrix, or as a cryogenic gas. Others have only been predicted using computational chemistry.

Non-classical covalent hydrides
8 18 8
LiH(H
2)
2[7] 		Be BH
3(H
2)
Na MgH
2(H
2)
n[32] 		AlH
3(H
2)
K Ca[33] ScH
3(H
2)
6[34][35] 		TiH
2(H
2)[36] 		VH
2(H
2)[12] 		CrH2(H2)2[37] Mn FeH
2(H
2)
3[38] 		CoH(H
2) 		Ni(H
2)
4 		CuH(H2) ZnH
2(H
2) 		GaH
3(H
2)
Rb Sr[33] YH
2(H
2)
3 		Zr NbH
4(H
2)
4[39] 		Mo Tc RuH
2(H
2)
4[40] 		RhH3(H2) Pd(H
2)
3 		AgH(H2) CdH(H
2) 		InH
3(H
2)[41]
Cs Ba[33] Lu Hf TaH
4(H
2)
4[12] 		WH
4(H
2)
4[42] 		Re Os Ir PtH(H
2) 		AuH
3(H
2) 		Hg Tl
Fr Ra Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Nh
↓
32 18
LaH
2(H
2)
2 		CeH
2(H
2) 		PrH
2(H
2)
2 		Nd Pm Sm Eu GdH
2(H
2) 		Tb Dy Ho Er Tm Yb
Ac ThH4(H2)4[43] Pa UH
4(H
2)
6[31] 		Np Pu Am Cm Bk Cf Es Fm Md No
Legend
Assessed[by whom?] Not assessed

Hydrogen solutions

Hydrogen has a highly variable solubility in the elements. When the continuous phase of the solution is a metal, it is called a metallic hydride or interstitial hydride, on account of the position of the hydrogen within the crystal structure of the metal. In solution, hydrogen can occur in either the atomic or molecular form. For some elements, when hydrogen content exceeds its solubility, the excess precipitates out as a stoichiometric compound. The table below shows the solubility of hydrogen in each element as a molar ratio at 25 °C (77 °F) and 100 kPa.

He
LiH
<1×10−4
[nb 1][44] 		Be B C N O F Ne
NaH
<8×10−7
[nb 2][45] 		MgH
<0.010
[nb 3][46] 		AlH
<2.5×10−8
[nb 4][47] 		Si P S Cl Ar
KH
<<0.01
[nb 5][48] 		CaH
<<0.01
[nb 6][49] 		ScH
≥1.86
[nb 7][50] 		TiH
2.00
[nb 8][51] 		VH
1.00
[nb 9][52] 		Cr MnH
<5×10−6
[nb 10][53] 		FeH
3×10−8
[54] 		Co NiH
3×10−5
[55] 		CuH
<1×10−7
[nb 11][56] 		ZnH
<3×10−7
[nb 12][57] 		Ga Ge As Se Br Kr
RbH
<<0.01
[nb 13][58] 		Sr YH
≥2.85
[nb 14][59] 		ZrH
≥1.70
[nb 15][60] 		NbH
1.1
[nb 16][61] 		Mo Tc Ru Rh PdH
0.724
[62] 		AgH
3.84×10−14
[63] 		Cd In Sn Sb Te I Xe
CsH
<<0.01
[nb 17][64] 		Ba Lu Hf TaH
0.79
[nb 18][65] 		W Re Os Ir Pt AuH
3.06×10−9
[62] 		HgH
5×10−7
[66] 		Tl Pb Bi Po At Rn
Fr Ra Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
↓
LaH
≥2.03
[nb 19][67] 		CeH
≥2.5
[nb 20][68] 		Pr Nd Pm SmH
3.00
[69] 		Eu Gd Tb Dy Ho Er Tm Yb
Ac Th Pa UH
≥3.00
[nb 21][70] 		Np Pu Am Cm Bk Cf Es FM Md No
Legend
Miscible Undetermined

Notes

  1. ^ Upper limit imposed by phase diagram, taken at 454 K.
  2. ^ Upper limit imposed by phase diagram, taken at 383 K.
  3. ^ Upper limit imposed by phase diagram, taken at 650 K and 25 MPa.
  4. ^ Upper limit imposed by phase diagram, taken at 556 K.
  5. ^ Upper limit imposed by phase diagram.
  6. ^ Upper limit imposed by phase diagram, taken at 500 K.
  7. ^ Lower limit imposed by phase diagram.
  8. ^ Limit imposed by phase diagram.
  9. ^ Limit imposed by phase diagram.
  10. ^ Upper limit imposed by phase diagram, taken at 500 K.
  11. ^ Upper limit imposed by phase diagram, taken at 1000 K.
  12. ^ Upper limit at 500 K.
  13. ^ Upper limit imposed by phase diagram.
  14. ^ Lower limit imposed by phase diagram.
  15. ^ Lower limit imposed by phase diagram.
  16. ^ Limit imposed by phase diagram.
  17. ^ Upper limit imposed by phase diagram.
  18. ^ Limit imposed by phase diagram.
  19. ^ Lower limit imposed by phase diagram.
  20. ^ Lower limit imposed by phase diagram.
  21. ^ Lower limit imposed by phase diagram.

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