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8.4.1: Hydrogen's Chemical Properties

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  • Chemical Properties

    As described in Section, hydrogen and helium are distinguished from all other elements in that their valence shell only consists of the 1s orbital. In the case of neutral atomic hydrogen this orbital is occupied by one electron. Consequently, the chemistry of hydrogen is distinguished by stable bonding arrangements in which the 1s orbital is "filled" by either

    • loss of an electron to give hydrogen ion, H+. In condensed phases these H+ ions are typically stabilized as Lewis base adducts (e.g. in species like H3O+ and NH4+)
    • gain of an electron to give hydride ion, H-. This type of bonding adequately explains the behavior of many metal hydrides.
    • pairwise sharing of electrons to give covalent E-H bonds that can adequately be described by Lewis theory.
    • multicenter sharing of electrons in multicenter covalent bonds, such as those in hydrides that bridge two or more atoms.
    • contributing electrons and orbitals to the band structure of a solid state lattice. This is common in interstitial/metallic hydrides.

    The first four of these possibilities are summarized in Scheme \(\sf{\PageIndex{I}}\).

    Scheme \(\sf{\PageIndex{I}}\). Common bonding arrangements for hydrogen. The E-H and E---E bonds in the bridging hydride represent sharing of two or more electrons among the three atoms.


    In forming H+ and E-H and H- like species the chemical properties of hydrogen are similar to both the alkali metals and the halogens. For this reason it is considered to have an ambiguous position in the periodic table (Figure \(\sf{\PageIndex{1}}\)).


    Figure \(\sf{\PageIndex{1}}\). Hydrogen has an ambiguous position in the Periodic table. It is usually placed above the alkali metal group although it could also be placed above the halogens.

    Elemental Hydrogen

    At room temperature and pressure elemental hydrogen exists in the form of dihydrogen, H2. Dihydrogen is a colorless odorless gas that finds wide industrial application.


    In the laboratory H2 may be prepared by the electrolytic or chemical reductions of water involving the half reaction.

    \[\sf{2H_2O + 2e^- \rightarrow H_2 + OH^-}\]

    Such reductions are commonly carried out on acidic solutions since the potentials required are lower, as may be seen from in the Pourbaix diagram for hydrogen given in Figure \(\sf{\PageIndex{2}}\).

    Figure \(\sf{\PageIndex{2}}\). Pourbaix diagram for hydrogen in water. By tem5psu -, CC BY-SA 3.0,

    A common method is to add Zn to a solution of hydrochloric acid.

    \[ \sf{2~H^+(aq)~~+~~Zn(s)~\rightarrow H_2(g)~~+~~Zn^{2+}(aq) }\]

    In electrolysis the electrons come from oxidation of water at the anode so that hydrogen production involves water splitting

    \[ \sf{2~H_2O(l)~~ \rightarrow~~2~H_2(g)~~+~~O_2(g)} \]

    Industrially it is more common to produce hydrogen via steam reforming of methane and other hydrocarbons.

    \[ \sf{CH_4(g)~~+~~H_2O(g)~\overset{Ni}{\longrightarrow}~CO(g)~~+~~3~H_2(g)~~~~(steam~reforming)} \]

    In this process the formally C4- of methane acts as the reductant. The product of steam reforming is a mixtures of CO and H2.. Similar mixtures can also be produced by the anaerobic thermal decomposition of organics and in goal gasification reactions. In all cases they are called syngas (i.e. synthesis gas) since they can be used in other industrial syntheses. Its s CO component is capable of acting as a reductant so additional hydrogen can be produced from it via the water-gas shift reaction.

    \[ \sf{CO(g)~~+~~H_2O(g)~\rightarrow~CO_2(g)~~+~~H_2(g)~~~~(water~gas~shift)} \]

    The "cracking" of hydrocarbons also serves as a source of industrially-produced hydrogen, although the alkenes so produced are perhaps even more important as the source of a majority of commodity organic chemicals.


    Much of the hydrogen produced industrially is consumed in the Haber-Bosch synthesis of ammonia. The quantities involved are such that hydrogen production for this purpose has been variously estimated to account for 1-2% of global energy consumption.


    Dihydrogen has also been considered for use as a fuel since its combustion is both highly exothermic and green, giving rise only to water.


    One of the major obstacles to the implementation of hydrogen as a fuel is that its production by steam reforming and the water-gas shift reaction generates CO & CO2 and costs more energy than is gained from its combustion. For this reason photocatalytic water splitting is considered one of the hold grails of energy research.


    Compounds of hydrogen are called hydrides whether or not they contain hydride anion. There are three main types of hydrides - ionic, covalent, and interstitial hydrides. As shown in the periodic table of hydrides given in Figure \(\sf{\PageIndex{3}}\), interstitial or metallic hydrides are formed by some transition metals while ionic hydrides are mainly formed by more electropositive metals and covalent hydrides by the nonmetals. The hydrides of Be, some metalloids, and some post transition metals are said to be intermediate hydrides since they form network covalent structures (sometimes in addition to molecular ones) and tend to function as bases and hydride donors like the ionic hydrides. Not all of the transition metals are known to form hydrides. No hydrides are known for the transition metals of groups 7-9, which are said to constitute the hydride gap.


    Figure \(\sf{\PageIndex{3}}\). Distribution of the different types of element hydrides across the periodic table. The figure is adapted from the Periodic Table at

    Ionic hydrides (a.k.a. saline hydrides)

    Ionic hydrides are metal salts of the hydride anion, H-. These are formed by the alkali and all the alkaline earth metals except Be. These are typically prepared by direct reaction of the metal and hydrogen.



    As salts of H- ionic hydrides form ionic lattices (the NaCl structure is common for MH, the Rutile and PbI2 for MH2).

    Chemically ionic hydrides act as

    • reducing agents towards metal oxides. For example

    \[\sf{2CaH_2(s)~~+~~TiO_2 (l)~\rightarrow~2CaO(s)~~+~~Ti(s)~~+~~2~H_2(aq)}\]

    • strong bases towards protic E-H bonds. All react exothermically with water to liberate hydrogen gas.


    For this reason CaH2 is widely used as a drying agent for organic solvents.

    Reactive metal hydrides may also be used to deprotonate reactive C-H bonds

    \[\sf{NaH(s)~~+~~CH_3C \equiv C-H(g)~\rightarrow~CH_3C \equiv C:^-Na^+~~+~~H_2(g)}\]

    The H:- ion in a ionic hydrides may in principle act as a nucleophile. However, in practice this application is limited to the less reactive and consequently more selective hydrides of aluminum and boron, both of which are usually classified as intermediate hydrides on account of the covalent character of their E-H bonds.

    Covalent and intermediate hydrides

    Covalent molecular hydrides are formed by the nonmetals, metalloids, and many post-transition metals. The chemical and physical properties they possess varies across the main group and depends somewhat on the row and whether the element hydride is electron deficient, electron rich, or electron precise. Specifically,

    Electron deficient hydrides are those of Be and the group 13 elements (B, Al, Ga, In, and Tl) for which the neutral monomeric element hydride (BeH2, BH3, AlH3, GaH3, InH3, and TlH3 does not possess enough electrons to satisfy the octet rule. Thus these hydrides commonly form dimers (B, Al, Ga, In, Tl) or polymers (Be) held together by bridging E-H-E bonds (Scheme \(\sf{\PageIndex{IIA}}\)). These E-H-E bonds are explained as three-center two-electron bonds in valence bond theory Scheme \(\sf{\PageIndex{IIB}}\) but may also be described in terms of molecular orbitals Note \(\sf{\PageIndex{1}}\).

    Scheme \(\sf{\PageIndex{II}}\). (A) Bridging E-H-E bonds in Al2H6 and (B) their valence bond description in terms of overlap between the H 1s and Al sp3 orbitals.


    Electron precise and electron rich hydrides are formed by C, N, O, F, and their heavier cogeners. These E-H bonds in these may be described as classical two center two electron E-H bonds of Lewis Theory. The electron precise and electron rich hydrides are distinguished in that the electron rich hydrides possess lone pair electrons while electron precise hydrides do not. In other words

    • electron precise hydrides are those of the group 14 elements, and include the alkanes, alkenes, and alkynes of carbon along with SiH4, GeH4, SnH4, and PbH4, of which the hydride adducts of group 13 EH3 compounds like BH4- and AlH4- are analogues.
    • electron rich hydrides are NH3, H2O,, HF and their heavier analogues (PH3, H2S, HCl, etc.). Regardless of the hydride's classification, the stability of element hydrides decreases down a group. For example, among the group 14 elements it follows the order CH4 > SiH4 > GeH4 > SnH4 > PbH4. The same is true of compounds possessing E-E bonds so that while a vast number of alkanes are known there are relatively few silanes, fewer germanes, and only the organic analougues of stannane are known (such as (CH3)3Sn-Sn(CH3)3).

    The distinction between Electron precise and electron rich hydrides is important mainly in thinking about the Lewis-acid base properties of the element hydrides. As illustrated in Scheme \(\sf{\PageIndex{III}}\), electron deficient hydrides tend to function as Lewis acids and electron rich hydrides as Lewis bases and Brønsted acids.

    Scheme \(\sf{\PageIndex{III}}\). In the absence of extremely strong acids or bases (A) electron-deficient hydrides like BH3 tend to act as Lewis acids in forming dducts with bases like THF while electron rich hydrides as can act as (B) Lewis bases through their lone pairs, as water with Cu2+ when anhydrous CuSO4 is dissolved in water, or (C) as Brønsted acids, as water does when it is used to quench the alkoxide product of a nucleophilic addition reaction.


    The reactivity of electron precise hydrides depends on the characteristics of E. For example, while most alkanes do not act as Lewis acids at carbon, row 3 and heavier electron precise hydrides can form trigonal bipyramidal adducts.

    Moreover, all element hydrides - whether electron deficient, precise, or rich - can function as a weak Brønsted acids or hydride donors depending on the polarity of the E-H bond. Hydrides in which hydrogen is bound to an electron rich and electornegative elements tend to act as Brønsted acids while those involving more electropositive elements tend to function as hydride donors, as illustrated in Scheme \(\sf{\PageIndex{IV}}\).

    Scheme \(\sf{\PageIndex{IV}}\). (A) The electron rich hydride of chlorine acts as Brønsted acid in forming a pyridinium complex while (B) the relatively electropositive hydride in tetrahydroaluminate is widely used as a hydride donor in organic chemistry, as illustrated by the use of Lithium aluminum hydride to form alcohols from ketones.


    Whether a given E-H bond is polarized so as to favor negative a partial positive or negative charge on hydrogen depends on the electronegativity of the element, as shown in in Figure \(\sf{\PageIndex{4}}\). As can be seen in Figure \(\sf{\PageIndex{4}}\), the metals give hydrides of \(\sf{E^{\delta-}-H^{\delta+}}\) bonds, the metalloids (including B) weakly polarized \(\sf{E^{\delta-}-H^{\delta+}}\) bonds and most nonmetals \(\sf{E^{\delta-}-H^{\delta+}}\) bonds.


    Figure \(\sf{\PageIndex{4}}\). Difference between element and hydrogen Pauling electronegativities. More positive values correspond to positively polarized hydrogen while more negative ones a larger partial negative charge on hydrogen.

    The ability of a given element hydride to function as an acid or hydride donor may be modified by a number of factors. One of these is the solvation energy of the species formed. According to Figure \(\sf{\PageIndex{4}}\), Germane (GeH4) might be expected to function as a hydride donor. However, it can be deprotonated in liquid ammonia, likely because of the large solvation energy of the resulting H+ ion.


    Electronic factors that affect the stability of the element hydride's conjugate acid or base forms also play a role. For example, carbon-hydrogen bonds are normally very weakly acidic but can (A) act as strong Brønsted acids when the resulting anion is highly stabilized or (B) This is illustrated by the well-known ability of C-H bonds to function as Brønsted acids, hydride donors, or neither depending on the electron-richness of the carbon center and the stability of the resulting structure (Scheme \(\sf{\PageIndex{V}}\)).

    Scheme \(\sf{\PageIndex{V}}\). (A) enolate chemistry such as that used in the formation of acetylacetate (acac) ligands is based on the ability of C-H bonds \(\sf{\alpha}\) to a carbonyl to act as acids and (B) a C-H bond of NADH functions as a hydride donor in biochemical systems. Notice the similarity in reactivity of the C-H hydride in NADH and the Al-H in LiAlH4 shown in in Scheme \(\sf{\PageIndex{IV}}\).


    Additional aspects of the acid-base chemistry of the element hydrides is described in 6: Acid-Base and Donor-Acceptor Chemistry, as is hydrogen's ability to form hydrogen bonds.

    Interstitial hydrides (a.k.a. metallic hydrides)

    In interstitial or metallic hydrides hydrogen dissolves in a metal to form nonstoichiometrix compounds (solid solutions) of formula MHn. They are called metallic hydrides since they possess the typical metallic properties of luster, hardness, and conductivity and are called interstitial hydrides because the H occupies interstices in a FCC, HCP, or BCC metal lattice, as illustrated schematically in Figure \(\sf{\PageIndex{5}}\).

    10.7D: Molecular Hydrides and Complexes Derived from them ...

    Figure \(\sf{\PageIndex{5}}\). Schematic illustration of the process of formation of an interstitial hydride showing breakup of dihydrogen and its uptake into the metal lattice. The metal lattice expands 10-20% during this process.1 Taken from

    The process of interstitial hydride formation is reversible and metals can dissolve varying amounts of hydrogen depending on the number of interstices avaialble. Because of this interstitial hydrides have been considered as storage materials for hydrogen.

    Note. \(\sf{\PageIndex{1}}\). A Qualitative Molecular Orbital Description of the bonding in Diborane

    The B-H-B bonds in diborane may also be explained using a molecular orbital description, as illustrated by the qualitative MO diagram in Figure \(\sf{PageIndex{6}}\).


    Figure \(\sf{\PageIndex{6}}\). Qualitative MO diagram for diborane. The symmetry adapted bonding and antibonding MOs contributing involved in bridge bonding are shown in violet. The derivation of this diagram is presented in 8.5.1. Properties of the Group 13 Elements and Boron Chemistry.


    1. Møller, K. T.; Jensen, T. R.; Akiba, E.; Li, H.-w., Hydrogen - A sustainable energy carrier. Progress in Natural Science: Materials International 2017, 27 (1), 34-40.

    Contributors and Attributions

    Stephen Contakes, Westmont College