HYDRIDES
Binary compounds of the elements with hydrogen are called hydrides. The type of hydride, which an element forms depends on its electro negativity, and hence on the type of bond formed. While there is not a sharp division between ionic, covalent and metallic bonding, it is convenient to consider hydrides in three classes:
1. ionic or salt-like hydrides
2. covalent or molecular hydrides
3. metallic or interstitial hydrides
Ionic or salt-like hydrides
At high temperature the metals of group 1 (alkali metals) and the heavier Group 2 metal (alkaline earth metals) Ca, Sr and Ba form ionic hydrides such as NaH and CaH2 . These compounds are solids with high melting points, and are classified as ionic (salt-like) hydrides. The evidence that they are ionic is:
1. Molten LiH (m.p. 6910C) conducts electricity, and H2 is liberated at the anode, thus confirming the presence of the hydride ion H–
2. The other ionic hydrides decompose before melting, but they may be dissolved in melts of alkali halides (e.g., CaH2 dissolves in a eutectic mixture of LiCl/KCl), and when the melt is electrolyzed then H2 is evolved at the anode.
3. The crystal structures of these hydrides are known, and they show evidence of directional bonding.
Lithium is more polarizing and hence more likely to form covalent compounds than the other metals. Thus if LiH is largely ionic, the others must be ionic, and thus contains the hydride ion H
The density of these hydrides is greater than that of the metal form which they were formed. This is explained by H– ions occupying holes in the lattice of the metal, without distorting the metal lattice. Ionic hydrides have high heats of formation, and are always stoichiometric.
Group 1 hydrides are more reactive than the corresponding
Group 2 hydrides, and reactivity increases down the group.
Except for LiH, ionic hydrides decompose elements on strong heating (400 – 5000C).
The hydrides ion H– is not very common, and it is unstable in water.
Ionic hydrides all react with water and liberate hydrogen.
LiH + H2O → LiOH + H2
CaH2 + 2H2O → Ca(OH)2 + 2H2
They are powerful reducing agents, especially at high temperatures, thought their reactivity towards water limits their usefulness.
2CO + NaH → H.COONa + C
SiCl4 + 4NaH → SiH4 + 4NaCl
PbSO4 + 2CaH2 → PbS + 2Ca(OH)2
NaH is used to produce important hydrides, particularly lithium aluminium hydride Li[AlH4] and sodium borohydride Na[BH4] ; Which have important uses as reducing agents in both organic and inorganic syntheses.
4LiH + AlCl3 → Li[AlH4] + 3LiCl
4NaH + B(OCH3)3 → Na[BH4] + 3NaOCH3
Covalent hydrides
Hydrides of the p-block elements are covalent. This would be expected since there is only a small electronegativity difference between these atoms and hydrogen. The compounds usually consists of discrete covalent molecules, with only weak Van der Waal’s forces holding the molecules together, and so they are usually volatile, and have low melting and boiling points. They do not conduct electricity. The formula of these hydrides is XHn .
Preparation of hydrides:-
1. By direct action.
| Group | 13 | 14 | 15 | 16 | 17 |
| B | C | N | O | F | |
| Al | Si | P | S | Cl | |
| Ga | Ge | As | Se | Br | |
| In | Sn | Sb | Te | I | |
| Pb | Bi | Po |
Elements which form covalent hydrides by direct action.
3H2 + N2 → 2NH3 (high temperature and pressure + catalyst, Haber process)
2H2 + O2 → 2H2O (Spark – explosive)
H2 + Cl2 → 2HCl (Burn – preparation of pure HCl)
2. Reaction of a halide with Li[AlH4] in a dry solvent such as Ether.
4BCl3 + 3Li[AlH4] → 2B2H6 + 3AlCl3 + 3LiCl
SiCl4 + Li[AlH4] → SiH4 + AlCl3 + LiCl
3. Treating the appropriate binary compound with acid.
2Mg3B2 + 4H3PO4 → B4H10 + 2Mg3(PO4)2 + H2
Al4C3 + 12HCl → 3CH4 + 4AlCl3
FeS + H2SO4 → H2S + FeSO4
Ca3P2 + 3H2SO4 → 2PH3 + 3CaSO4
4. Reaction of an oxoacid with Na[BH4] in aqueous solution.
4H3AsO3 + 3Na[BH4] → 4AsH3 + 3NaOH
5. Converting one hydride into another by hydrolysis (heating).
B4H10 → B2H6 + other products
6. A silent electric discharge or microwave discharge may produce long chains form simple hydrides.
GeH4 → Ge2H6 → Ge3H8 → upto Ge9H20
Metallic (or interstitial) hydrides
Many of the elements in the d-block and the lanthanide and actinide elements in the f-block, react with H2 and form metallic hydrides. However, the elements in the middle of the d-block do not form hydrides. The absence of hydrides in this part of the periodic table is sometimes called the hydrogen gap.
Metallic hydrides are usually prepared by heating the metal with hydrogen under high pressure. (if heated to higher temperature the hydrides decompose, and this may be used as a convenient method of making very pure hydrogen.)
These hydrides generally have properties similar to those of the parent metals: they are hard, have a metallic luster, conduct electricity and have magnetic properties. The hydrides are less than the parent metal, because the crystal lattice has expanded through the inclusion of hydrogen.
This distortion of the crystal lattice may make the hydride brittle. Thus when the hydride is formed a solid piece of metal turns into finely powdered hydride. If the finely powered hydrides are heated they decompose giving hydrogen and very finely divided metal. These finely divided metals may be used as catalysts. They are also used in metallurgy in powder fabrication and zirconium hydride has been used as a moderator in nuclear reactors.
The Pd/H2 system is both extraordinary and interesting. When red hot Pd is cooled in H2 it may absorb or occlude up to 935 times its own volume of H2 gas. This may be used to separate H2 or deuterium D2 from He or other gases. The hydrogen is given off when the metal is heated and this provides an easy method of weighing H2. The limiting formula is PdH0.7, but neither the structure nor the nature of the interaction between Pd and H2 are understood. As hydrogen is absorbed, the metallic conductivity decreases, and the material eventually becomes a semiconductor. The hydrogen is mobile and diffuses throughout the metal.
