Hydrogen is the lightest, simplest and most abundant element in the universe, yet it occupies an unusual position in the periodic table. It has one electron and one proton, and its position in the periodic table has been debated because it resembles both the alkali metals (group 1) and the halogens (group 17). Its chemistry is unique because its single electron allows only one bond, and its small size gives its compounds exceptional properties.
Hydrogen occurs as three isotopes: protium (H), deuterium (D) and tritium (T). Tritium is radioactive and used as a tracer, while deuterium finds applications in nuclear energy. In the Earth's crust and oceans, hydrogen is abundant in combination with other elements, mainly in water and hydrocarbons, and it is present in all living matter.
Industrially, hydrogen is a crucial chemical feedstock. It is used in the Haber process to manufacture ammonia, in the hydrogenation of oils to produce fats, in the production of hydrochloric acid, and as a fuel in rockets and fuel cells. Because its compounds range from the essential electron-pair donor in acids to the metal hydrides used for hydrogen storage, this chapter covers hydrogen's position, preparation, properties and the full range of its chemistry.
Hydrogen's position in the periodic table is unique because it resembles both the alkali metals and the halogens. Like the alkali metals, it has one electron in its valence shell and forms a +1 oxidation state, giving the cation H+. Like the halogens, it can gain one electron to form the hydride ion H- and can form diatomic molecules such as H2 and Cl2.
Hydrogen differs from both groups, however. It has a much higher ionisation enthalpy than the alkali metals and does not form a metallic lattice. It is a non-metal with a very high ionisation enthalpy of 1312 kJ mol^-1, roughly comparable to the halogens. Because of these anomalies, hydrogen is often placed separately above group 1, with its unique properties acknowledged.
Some chemists place hydrogen at the top of group 17 because of similarities with the halogens, while others place it at the top of group 1. The most widely accepted treatment recognises hydrogen as a unique element, and this is the convention adopted in the modern periodic table.
Hydrogen has three isotopes: protium (1H), deuterium (2H or D) and tritium (3H or T). Protium has one proton and no neutrons and makes up about 99.985% of natural hydrogen. Deuterium has one proton and one neutron, making up about 0.015%, and tritium has one proton and two neutrons.
The isotopes differ in atomic mass, which leads to differences in their physical properties. Heavy water (D2O) has a higher boiling point and higher density than ordinary water. Because of its higher mass, deuterium forms bonds that are slightly stronger than protium bonds, and the rates of reactions involving D2O differ from those of H2O.
Tritium is radioactive, decaying by beta emission with a half-life of 12.3 years. It is used as a tracer in biological and chemical studies and is produced artificially. The isotopes are separated using methods based on their different physical properties, such as electrolysis of water and fractional distillation of liquid hydrogen.
In the laboratory, hydrogen is prepared by the reaction of zinc with dilute hydrochloric acid or sulphuric acid:
$$\text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2$$
Industrially, hydrogen is produced by the electrolysis of water, the reaction of steam with hydrocarbons (steam reforming), and the water-gas shift reaction. The water-gas reaction produces a mixture called water gas:
$$\text{C} + \text{H}_2\text{O} \rightarrow \text{CO} + \text{H}_2$$
Dihydrogen is a colourless, odourless, tasteless gas that is the lightest substance known. It is sparingly soluble in water, easily diffuses, and has very low density, making it useful for filling balloons. It is highly inflammable and burns in air with a pale blue flame, forming water.
Dihydrogen reacts with many elements and compounds. It combines with the halogens to form hydrogen halides, with oxygen to form water, and with nitrogen in the Haber process to form ammonia. It is a reducing agent, reducing metal oxides to metals, as in the reduction of copper oxide to copper.
Hydrogen forms compounds with most elements, called hydrides. These are classified into three types based on the nature of the bond. Ionic or saline hydrides are formed by the alkali and alkaline earth metals (except Be and Mg), in which hydrogen exists as the hydride ion H-. These are white crystalline solids that conduct electricity when molten and react vigorously with water to release hydrogen.
Covalent or molecular hydrides are formed by the non-metals and metalloids, such as CH4, NH3, H2O and HCl. These have low melting and boiling points and are generally volatile. Metallic or interstitial hydrides are formed by the transition metals and lanthanoids, in which hydrogen occupies the interstitial sites of the metal lattice. These are non-stoichiometric, hard, and often behave as alloys.
The properties of covalent hydrides vary with the electronegativity of the element bonded to hydrogen. The electron-rich hydrides, like NH3 and H2O, act as Lewis bases and form hydrogen bonds. This classification is important for predicting the properties and reactions of hydrides across the periodic table.
Water is the most abundant hydrogen compound and is essential for life. It has a bent structure with a bond angle of 104.5 degrees and is a highly polar molecule because of its two lone pairs and the high electronegativity of oxygen. This polarity, combined with intermolecular hydrogen bonding, gives water its exceptional properties.
Water has anomalously high boiling and melting points, high surface tension, high heat capacity and high latent heat of vapourisation, all resulting from hydrogen bonding. It expands on freezing, so ice is less dense than liquid water and floats. Water is an excellent solvent, dissolving ionic and polar compounds through its high dielectric constant and hydrogen-bonding ability.
Chemically, water reacts with metals, oxides and salts. It reacts with alkali and alkaline earth metals to release hydrogen, with metal oxides to form bases, and with non-metal oxides to form acids. It is amphoteric, acting both as an acid and as a base, and it participates in hydrolysis reactions of salts and esters.
Heavy water (D2O) is the oxide of deuterium. It is used as a moderator in nuclear reactors to slow down neutrons, as a tracer in the study of chemical and biological reactions, and in the preparation of deuterium compounds. Its physical properties differ from ordinary water, with higher density, boiling point and melting point.
Hydrogen peroxide (H2O2) is an important compound of hydrogen and oxygen, prepared by electrolysis of sulphuric acid or by the anthraquinone process. It is a colourless liquid that is unstable and decomposes slowly to water and oxygen:
$$2\text{H}_2\text{O}_2 \rightarrow 2\text{H}_2\text{O} + \text{O}_2$$
Hydrogen peroxide is a powerful oxidising agent, and in the presence of strong oxidising agents it can also act as a reducing agent. It bleaches fabrics and hair, is used as an antiseptic and disinfectant, and finds applications in the recovery of metals and in rocket propulsion. It is stored in wax-lined glass or plastic bottles with stabilisers because it decomposes in the presence of light and metals.
Dihydrogen has a wide range of applications. The largest use is in the Haber process for the synthesis of ammonia, which is then used to make fertilisers, explosives and plastics. It is also used in the hydrogenation of vegetable oils to produce vanaspati ghee, in the manufacture of hydrochloric acid and methanol, and in the reduction of metal oxides in metallurgy.
Liquid hydrogen is used as rocket fuel, combining with liquid oxygen to produce enormous thrust. Hydrogen fuel cells convert the chemical energy of hydrogen and oxygen directly into electricity, producing only water as a by-product, making them an environmentally clean energy source. Hydrogen is also used in the petroleum industry for the cracking and desulphurisation of crude oil.
The growing interest in a hydrogen economy stems from its clean combustion product and its high energy density per unit mass. Hydrogen can be produced from water using renewable electricity, stored and transported, and used in fuel cells for vehicles and power generation, offering a sustainable path to reduce carbon emissions.
| Isotope | Symbol | Protons | Neutrons | Abundance | Special use |
|---|---|---|---|---|---|
| Protium | 1H | 1 | 0 | 99.985% | Ordinary hydrogen |
| Deuterium | 2H or D | 1 | 1 | 0.015% | Nuclear moderator (D2O) |
| Tritium | 3H or T | 1 | 2 | Trace (radioactive) | Tracer studies |
| Type | Elements | Nature | Example |
|---|---|---|---|
| Ionic (saline) | s-block (except Be, Mg) | White solid, contains H- | NaH, CaH2 |
| Covalent (molecular) | Non-metals and metalloids | Volatile, low melting | CH4, NH3, HCl |
| Metallic (interstitial) | Transition metals | Non-stoichiometric, hard | LaH2.87, PdHx |
| Property | H2O | D2O |
|---|---|---|
| Molar mass | 18.02 | 20.03 |
| Boiling point | 100 degrees Celsius | 101.42 degrees Celsius |
| Density (25 degrees Celsius) | 0.997 g/mL | 1.106 g/mL |
| Heat of fusion | 6.01 kJ/mol | 6.28 kJ/mol |
Hydrogen is an element of unique character whose chemistry spans the entire periodic table. Its intermediate position reflects its resemblance to both alkali metals and halogens, and its three isotopes provide useful probes in nuclear and tracer science. Dihydrogen is a versatile fuel and reducing agent, prepared in the laboratory and produced industrially in huge quantities for ammonia synthesis. Its hydrides illustrate the spectrum of bonding from ionic to covalent to metallic, while water and hydrogen peroxide display the extraordinary consequences of hydrogen bonding and the dual oxidising-reducing nature of oxygen compounds. As the search for clean energy intensifies, hydrogen's role as a zero-carbon fuel makes its chemistry increasingly important for the future of technology and sustainability.