The English chemist Henry Cavendish (1731–1810) was the first to isolate hydrogen. Because the element produces water when burned in air, the French chemist Antoine Lavoisier  (Figure 3.1) gave it the name hydrogen, which means “water producer” (Greek: hydro, water; gennao, to produce).

Hydrogen is the most abundant element in the universe. It is the nuclear fuel consumed by our Sun and other stars to produce energy.  (Section 21.8) Although about 75% of the known mass of the universe is hydrogen, it constitutes only 0.87% of Earth's mass. Most of the hydrogen on our planet is found associated with oxygen. Water, which is 11% hydrogen by mass, is the most abundant hydrogen compound.

Isotopes of Hydrogen

The most common isotope of hydrogen, , has a nucleus consisting of a single proton. This isotope, sometimes referred to as protium,* makes up 99.9844% of naturally occurring hydrogen.

Two other isotopes are known: , whose nucleus contains a proton and a neutron, and , whose nucleus contains a proton and two neutrons. The  isotope, deuterium, makes up 0.0156% of naturally occurring hydrogen. It is not radioactive and is often given the symbol D in chemical formulas, as in D2O (deuterium oxide), which is known as heavy water.

Because an atom of deuterium is about twice as massive as an atom of protium, the properties of deuterium-containing substances vary somewhat from those of the protium-containing analogs. For example, the normal melting and boiling points of D2O are 3.81°C and 101.42 °C, respectively, versus 0.00 °C and 100.00 °C for H2O. Not surprisingly, the density of D2O at 25 °C (1.104 g/mL) is greater than that of H2O (0.997 g/mL). Replacing protium with deuterium (a process called deuteration) can also have a profound effect on reaction rates, a phenomenon called a kinetic-isotope effect. For example, heavy water can be obtained from the electrolysis  of ordinary water because the small amount of naturally occurring D2O in the sample undergoes electrolysis more slowly than H2O and, therefore, becomes concentrated during the reaction.

The third isotope, tritium, is radioactive, with a half-life of 12.3 yr:

Because of its short half-life, only trace quantities of tritium exist naturally. The isotope can be synthesized in nuclear reactors by neutron bombardment of lithium-6:

Deuterium and tritium are useful in studying reactions of compounds containing hydrogen. A compound is “labeled” by replacing one or more ordinary hydrogen atoms with deuterium or tritium at specific locations in a molecule. By comparing the locations of the label atoms in reactants and products, the reaction mechanism can often be inferred. When methyl alcohol (CH3OH) is placed in D2O, for example, the H atom of the O — H bond exchanges rapidly with the D atoms, forming CH3OD. The H atoms of the CH3 group do not exchange. This experiment demonstrates the kinetic stability of C — H bonds and reveals the speed at which the O — H bond in the molecule breaks and re-forms.

Properties of Hydrogen

Hydrogen is the only element that is not a member of any family in the periodic table. Because of its ls 1 electron configuration, it is generally placed above lithium in the table. However, it is definitely not an alkali metal. It forms a positive ion much less readily than any alkali metal. The ionization energy of the hydrogen atom is 1312 kJ/mol, whereas that of lithium is 520 kJ/mol.

Hydrogen is sometimes placed above the halogens in the periodic table because the hydrogen atom can pick up one electron to form the hydride ion, H, which has the same electron configuration as helium. However, the electron affinity of hydrogen, E = – 73 kJ/mol, is not as large as that of any halogen. In general, hydrogen shows no closer resemblance to the halogens than it does to the alkali metals.

Elemental hydrogen exists at room temperature as a colorless, odorless, tasteless gas composed of diatomic molecules. We can call H2dihydrogen, but it is more commonly referred to as either molecular hydrogen or simply hydrogen. Because H2 is nonpolar and has only two electrons, attractive forces between molecules are extremely weak. As a result, its melting point (–259 °C) and boiling point (–253 °C) are very low.

The H — H bond enthalpy (436 kJ/mol) is high for a single bond.  (Table 8.4) By comparison, the Cl — Cl bond enthalpy is only 242 kJ/mol. Because H2 has a strong bond, most reactions involving H2 are slow at room temperature. However, the molecule is readily activated by heat, irradiation, or catalysis. The activation generally produces hydrogen atoms, which are very reactive. Once H2 is activated, it reacts rapidly and exothermically with a wide variety of substances.


If H2 is activated to produce H+, what must the other product be?

Hydrogen forms strong covalent bonds with many other elements, including oxygen; the O — H bond enthalpy is 463 kJ/mol. The formation of the strong O — H bond makes hydrogen an effective reducing agent for many metal oxides. When H2 is passed over heated CuO, for example, copper is produced:

When H2 is ignited in air, a vigorous reaction occurs, forming H2O. Air containing as little as 4% H2 by volume is potentially explosive. Combustion of hydrogen-oxygen mixtures is used in liquid-fuel rocket engines such as those of the Space Shuttle. The hydrogen and oxygen are stored at low temperatures in liquid form.

Production of Hydrogen

When a small quantity of H2 is needed in the laboratory, it is usually obtained by the reaction between an active metal such as zinc and a dilute strong acid such as HCl or H2SO4:

Large quantities of H2 are produced by reacting methane with steam at 1100 °C. We can view this process as involving two reactions:

Carbon heated with water to about 1000 °C is another source of H2:

This mixture, known as water gas, is used as an industrial fuel.


The reaction of hydrogen with oxygen is highly exothermic:

Because the only product of the reaction is water vapor, the prospect of using hydrogen as a fuel in fuel cells is attractive.  (Section 20.7) Alternatively, hydrogen could be combusted directly with oxygen from the atmosphere in an internal combustion engine. In either case, it would be necessary to generate elemental hydrogen on a large scale and arrange for its transport and storage.

FIGURE 22.4 The “hydrogen economy” would require hydrogen to be produced from various sources and would use hydrogen in energy-related applications.

FIGURE 22.4 illustrates various sources and uses of H2 fuel. The generation of H2 through electrolysis of water is in principle the cleanest route, because this process—the reverse of Equation 22.11— produces only hydrogen and oxygen.  (Figure 1.7 and Section 20.9) However, the energy required to electrolyze water must come from somewhere. If we burn fossil fuels to generate this energy, we have not advanced very far toward a true hydrogen economy. If the energy for electrolysis came instead from a hydroelectric or nuclear power plant, solar cells, or wind generators, consumption of nonrenewable energy sources and undesired production of CO2 could be avoided.

RELATED EXERCISES: 22.29, 22.30, 22.91

Electrolysis of water consumes too much energy and is consequently too costly to be used commercially to produce H2. However, H2 is produced as a by-product in the electrolysis of brine (NaCl) solutions in the course of commercial Cl2 and NaOH manufacture:


What are the oxidation states of the H atoms in Equations 22.7-22.12?

Uses of Hydrogen

Hydrogen is commercially important. About 5.0 × 1010 kg (50 million metric tons) is produced annually across the world. About half of the H2 produced is used to synthesize ammonia by the Haber process.  (Section 15.2) Much of the remaining hydrogen is used to convert high-molecular-weight hydrocarbons from petroleum into lower-molecular-weight hydrocarbons suitable for fuel (gasoline, diesel, and others) in a process known as cracking. Hydrogen is also used to manufacture methanol via the catalytic reaction of CO and H2 at high pressure and temperature:

Binary Hydrogen Compounds

Hydrogen reacts with other elements to form three types of compounds: (1) ionic hydrides, (2) metallic hydrides, and (3) molecular hydrides.

The ionic hydrides are formed by the alkali metals and by the heavier alkaline earths (Ca, Sr, and Ba). These active metals are much less electronegative than hydrogen. Consequently, hydrogen acquires electrons from them to form hydride ions (H):

The hydride ion is very basic and reacts readily with compounds having even weakly acidic protons to form H2:

Ionic hydrides can therefore be used as convenient (although expensive) sources of H2.

Calcium hydride (CaH2) is used to inflate life rafts, weather balloons, and the like where a simple, compact means of generating H2 is desired (FIGURE 22.5).


This reaction is exothermic. Is the beaker on the right warmer or colder than the beaker on the left?

FIGURE 22.5 The reaction of CaH2 with water.

The reaction between H and H2O (Equation 22.15) is an acid–base reaction and a redox reaction. The H ion, therefore, is a good base and a good reducing agent. In fact, hydrides are able to reduce O2 to OH:

For this reason, hydrides are normally stored in an environment that is free of both moisture and air.

Metallic hydrides are formed when hydrogen reacts with transition metals. These compounds are so named because they retain their metallic properties. They are not molecular substances, just as metals are not. In many metallic hydrides, the ratio of metal atoms to hydrogen atoms is not fixed or in small whole numbers. The composition can vary within a range, depending on reaction conditions. TiH2 can be produced, for example, but preparations usually yield TiH1.8. These nonstoichiometric metallic hydrides are sometimes called interstitial hydrides. Because hydrogen atoms are small enough to fit between the sites occupied by the metal atoms, many metal hydrides behave like interstitial alloys.  (Section 12.3)

The most interesting interstitial metallic hydride is that of palladium. Palladium can take up nearly 900 times its volume of hydrogen, making it very attractive for hydrogen storage in any possible future “hydrogen economy.” However, to be practical, any hydrogen-storage compound will have to contain 75% or more hydrogen by mass and be able to charge and discharge hydrogen quickly and safely near room temperature.


Palladium has a density of 12.023 g/cm3. Can a sample of Pd that has a volume of 1cm3 increase its mass to over 900 g by adsorbing hydrogen?

The molecular hydrides, formed by nonmetals and metalloids, are either gases or liquids under standard conditions. The simple molecular hydrides are listed in FIGURE 22.6, together with their standard free energies of formation,  (Section 19.5) In each family the thermal stability (measured as ) decreases as we move down the family. (Recall that the more stable a compound is with respect to its elements under standard conditions, the more negative  is.)


Which is the most thermodynamically stable hydride? Which is the least thermodynamically stable?

FIGURE 22.6 Standard free energies of formation of molecular hydrides. All values are kilojoules per mole of hydride.