Part A: The Essentials
Hydrogen is the most abundant element in the universe and the tenth most abundant by mass on Earth, where it is found in the oceans, minerals, and in all forms of life. The partial depletion of elemental hydrogen from Earth reflects its volatility during formation of the planet.
The stable form of elemental hydrogen under normal conditions is dihydrogen, H2, which occurs at trace levels in the Earth's lower atmosphere (0.5 ppm) and is essentially the only component of the extremely thin outer atmosphere.
Major Uses of Hydrogen
10.1 The Element
The hydrogen atom, with ground-state configuration 1s1, has only one electron so it might be thought that the element's chemical properties will be limited, but this is far from the case. Hydrogen has richly varied chemical properties and forms compounds with nearly every other element.
Hydrogen ranges in character from being a strong Lewis base (the hydride ion, H−) to being a strong Lewis acid (as the hydrogen cation, H+, the proton).
(a) The Atom and Its Ions
The Three Isotopes of Hydrogen
The free hydrogen cation (H+, the proton) has a very high charge-to-radius ratio and is a very strong Lewis acid. In the gas phase it readily attaches to other molecules and atoms; it even attaches to He to form HeH+. In the condensed phase, H+ is always found in combination with a Lewis base.
(b) Properties and Reactions
One of the strongest single bonds known
Very short bond
Weak intermolecular forces
Dissociation of H2
Combustion Reaction
10.2 Simple Compounds
The nature of the bonding in binary compounds of hydrogen (EHn) is largely rationalized by noting that an H atom has:
- High ionization energy: 1310 kJ mol−1
- Low but positive electron affinity: 77 kJ mol−1
- Intermediate electronegativity: 2.2 (Pauling scale)
(a) Classification of Binary Compounds
Discrete molecules, covalent E–H bonds
Ionic, with electropositive elements
(b) Thermodynamic Considerations
Average Bond Energies for Binary Molecular Hydrides (kJ mol−1)
(c) Reactions of Binary Compounds
| Condition | Reaction Type | Product |
|---|---|---|
| χ(E) ≈ χ(H) | Homolytic cleavage | E• + H• (radicals) |
| χ(E) > χ(H) | Heterolytic → Protonic | E− + H+ (Brønsted acid) |
| χ(E) < χ(H) | Heterolytic → Hydridic | E+ + H− (hydride donor) |
Part B: The Detail
10.3 Nuclear Properties
Effect of Deuteration on Physical Properties
| Property | H2 | D2 | H2O | D2O |
|---|---|---|---|---|
| Boiling point (°C) | −252.8 | −249.7 | 100.0 | 101.4 |
| Bond enthalpy (kJ mol−1) | 436.0 | 443.3 | 463.5 | 470.9 |
Nuclear Fusion Reaction
10.4 Production of Dihydrogen
In 2012, world production of H2 exceeded 65 Mt. Most H2 is used for ammonia synthesis (Haber process), hydrogenation of fats, hydrocracking, and organic chemical manufacture.
(a) Small-Scale Preparation
(b) Production from Fossil Sources
Steam Reforming
Coal Gasification
Water Gas Shift Reaction
(c) Production from Renewable Sources
The Earth receives about 100,000 TW from the Sun — approximately 7000 times greater than global energy consumption (15 TW). Technologies under development include high-temperature solar H2 production and photoelectrochemical "artificial photosynthesis".
10.5 Reactions of Dihydrogen
Initiation: Br2 → Br• + Br•
Propagation: Br• + H2 → HBr + H• ; H• + Br2 → HBr + Br•
Termination: H• + H• → H2 ; Br• + Br• → Br2
10.6 Compounds of Hydrogen
(a) Molecular Hydrides
| Type | Description | Examples |
|---|---|---|
| Electron-precise | All valence electrons in bonds | CH4, SiH4, GeH4 |
| Electron-rich | Lone pairs on central atom | NH3, H2O, HF |
| Electron-deficient | 3c,2e bonds used | B2H6, AlH3 |
An E–H bond between an electronegative element E and hydrogen is highly polar. The partially positive H can interact with a lone pair on E of another molecule, forming a hydrogen bond.
| Hydrogen Bond | kJ mol−1 | Covalent Bond | kJ mol−1 |
|---|---|---|---|
| HO–H···OH2 | 22 | O–H | 464 |
| F–H···FH | 29 | F–H | 567 |
| F···H···F− | >155 | F–H | 567 |
(b) Saline Hydrides
Ionic solids containing discrete H− ions. Ionic radius varies from 126 pm in LiH to 154 pm in CsH. Group 1 hydrides adopt rock salt structure.
(c) Metallic Hydrides
Nonstoichiometric compounds (e.g., ZrH1.30 to ZrH1.75) with metallic lustre and conductivity. The "hydride gap" (Groups 7-9) contains no stable binary hydrides.
MgH2: 8% | LiBH4: 20% | LiNH2: 10% | LaNi5H6: 2% (higher H density than liquid H2)
(d) Hydrido and Dihydrogen Complexes
The H atom is usually regarded as H− (hydrido) ligand. H2 can also coordinate intact using σ-donation and π-backbonding. If the metal is electron-rich, oxidative addition occurs.
[IrCl(CO)(PPh3)2] + H2 → [IrCl(CO)(H)2(PPh3)2]
Ir(I) → Ir(III), two H− ligands formed.
10.7 General Methods for Synthesis
Summary: Key Properties
| Property | Value |
|---|---|
| Atomic number | 1 |
| Configuration | 1s1 |
| Electronegativity | 2.2 |
| Ionization energy | 1310 kJ mol−1 |
| H–H bond enthalpy | 436 kJ mol−1 |
| H–H bond length | 74 pm |
Special Topics
H2 is cycled by microbial organisms using metalloenzymes. Fermentative bacteria produce H2 as waste; methanogens use it to produce CH4; Desulfovibrio produces H2S. High breath H2 levels (>70 ppm) can diagnose lactose intolerance.
| Fuel | Specific Enthalpy (MJ/kg) | Energy Density (MJ/dm³) |
|---|---|---|
| Liquid H2 | 120 | 8.5 |
| Petrol | 46 | 34.2 |
| Li-ion battery | 2.0 | 6.1 |
Ni metal-hydride batteries: M–H bond enthalpy ideally 25–50 kJ mol−1. Too low → H2 evolved; too high → not reversible. Uses alloys with Li, Mg, Al, V, Mn, Zr, etc.