Hydrogen
Hydrogen is a chemical element with atomic number Z = 1 and chemical symbol H. It is the most abundant element in the universe, (about 90% in number of atoms, about 70% in mass). It exists primarily in interstellar gas clouds and stars, but it is also thought that hydrogen is a major component of the planet Jupiter. On earth hydrogen is mainly bound to other elements, to carbon in carbohydrates, to oxygen in water, and to nitrogen in ammonia. Measured in number of atoms, hydrogen is the third abundant element (15%) on earth. Atomic hydrogen is very reactive and does not naturally appear in free form. The hydrogen molecule H2 is a stable gas at ambient temperature and pressure.
Isotopes
The hydrogen atom has two stable and one unstable isotope. Its averaged (weighted by abundance) mass is 1.00794 u. Only protium (1H) and deuterium (2H = D) occur in sizeable quantities in nature. The isotope tritium (3H = T) is largely man-made and has a half-life of about 12.5 years. Masses (in unified atom mass unit u), and abundances (percentage in brackets) are:
- 1H 1.0078250321 (99.9885%)
- 2H 2.0141017780 (0.0115%)
- 3H 3.0160492675 (10−15%)
Physical properties
Under normal temperature and pressure H2 is a colorless, tasteless, and odorless gas. The most commonly occurring molecule 1H2 has a boiling point of 20.28 K and a melting point of 14.01 K. For the deuterium molecule these temperatures are about 4 K higher: D2 has boiling point 23.65 K and melting point 18.65 K.
At ambient temperature and pressure the density of gaseous hydrogen is 0.08988 kg/m³. The density of the liquid (at boiling point) is 70.8 kg/m³ and the density of the solid is 70.6 kg/m³ (at 11 K).
The triple point is at temperature 13.81 K with a pressure 7.042 kPa. The critical point is at 33.15 K, 1.298 MPa. The critical density is 30.09 kg/m3.
Chemical properties
The neutral atom has nuclear charge e (the elementary charge) and one electron (charge e). The electronic wave functions of the ground and excited states of the H-atom are known exactly, because the Schrödinger equation can be solved analytically for one-electron atoms, see hydrogen-like atoms. In the ground state the electron is in a 1s orbital (a nodeless spherically symmetric function with energy -13.59 eV).
The molecule H2 has an equilibrium distance re = 1.4011 bohr, an equilibrium dissociation energy De: 38 297.1 cm−1 (458.135 kJ/mol) and a dissociation energy D0: 36118.069 cm−1 (is 432.068 kJ/mol, the energy for dissociation from from the vibrational ground state).[1] Because of this relatively high binding energy, molecular hydrogen is fairly inert. Many reactions of hydrogen require high temperatures and a catalyst.
Hydrogen resembles alkali metals, like Li and Na, in that it can lose its (outer) electron and become the cation H+. Strong acids, dissolved in water, dissociate and release H+, for instance HCl → Cl− + H+. Note, however, that the free hydrogen nucleus, being a bare charge, will not remain free in solution, but will immediately associate with solvent molecules, e.g., with water it will associate to H3O+.
Hydrogen also resembles halogens, like Cl, in that it can form an anion. It does so in saline hydrides (LiH, NaH, CaH2, etc.), in which to a large extent the binding is ionic, i.e., it is of the type M+–H−. These compounds are solids with a high melting point.
Hydrogen also forms bonds that are mainly covalent, as in carbohydrates, water, and ammonia.
Hydrogen gas is highly flammable and will burn at low concentrations in air. It combusts according to the following exothermic equation:
- 2 H2(g) + O2(g) → 2 H2O(l) + ΔH
where the enthalpy of combustion ΔH is 571.6 kJ[2] (per mole O2).
Preparation
A classical way of preparing hydrogen is by pouring sulphuric acid on zinc:
- H2SO4 + Zn → ZnSO4 + H2
As was already known by Henry Cavendish in the second half of the 18th century, zinc may be replaced in this recipe by iron or tin and sulphuric acid by hydrochloric acid (HCl in water).
Another way of preparing hydrogen gas is by immersing aluminum trimmings in a solution of sodium hydroxide in water:
- 2Al + 2NaOH + 2H2O → 2NaAlO2 + 3H2.
This reaction runs violently.
By electrolysis of water we observe at the cathode (the negative electrode with a voltage difference with the positive electrode of at least 0.83V) the following reaction:
- 4H2O + 4e− → 2H2 + 4OH−.
The OH− anions move to the anode (positive electrode) and decompose according to
- 4OH− → O2 + 2H2O + 4e−,
so that the net effect is electron transport from cathode to anode and the formation of gaseous hydrogen at the cathode and gaseous oxygen at the the anode. This process, which requires electrical power, is energetically not very efficient and at present not competitive with thermal cracking of hydrocarbons and steam reforming of natural gas (CH4).
In steam reforming of natural gas (SMR—steam methane reforming) water vapor reacts with methane to yield carbon monoxide and H2 at high temperatures (700 – 1000 °C),
This process takes place under 3-25 bar pressure (1 bar = 100 kPa = 0.987 atm) in the presence of a nickel catalyst. SMR is endothermic—that is, heat must be supplied to the process for the reaction to proceed. The product mixture is known as synthesis gas ("syngas") because it is often used directly for the production of methanol and related compounds. Hydrocarbons other than methane can be used to produce synthesis gas
- CxHy + x H2O → x CO + (x + 0.5y) H2,
but steam reforming of heavier hydrocarbons has the problem that it gives pure carbon (coke) as a contaminating byproduct. The steam reforming of methane is by far the largest source of hydrogen gas today.
With the use of an iron oxide catalyst it is possible to generate hydrogen by the water-gas shift reaction, which is slightly exothermic,
- CO + H2O → CO2 + H2.
The carbon dioxide can be frozen out from the product mixture.
Usage
It is expected that in the future hydrogen will replace gasoline as a fuel for cars and trucks (in the so-called hydrogen economy). At this moment the main use of hydrogen is in the manufacturing of ammonia by catalytic reaction of hydrogen with nitrogen. The produced ammonia is used for the manufacturing of nitrogen-containing fertilizers.
Another important application of H2 is hydrocracking. This is a catalytic cracking process of heavy hydrocarbons in the presence of hydrogen to lighter, more desirable, hydrocarbons. Hydrogen prevents the formation of polycyclic aromatic compounds. Another important role of hydrogen in the cracking is the reduction of tar formation and the prevention of buildup of coke on the catalyst. Hydrogenation also converts sulfur and nitrogen atoms present in the feedstock to hydrogen sulfide and ammonia, which can easily be removed from the product. The products of hydrocracking are saturated hydrocarbons ranging from ethane, LPG to high-octane gasoline.
Other uses of hydrogen are in high temperature welding and as rocket fuel.
Since the density of air is about 15 times that of hydrogen, (the molecular mass of hydrogen being about 1/15 of that of nitrogen and oxygen), hydrogen is a good gas for balloons. However, since it is very inflammable, it is not used to lift people, but it is regularly used in weather balloons.
Hydrogen's isotopes also have applications. Deuterium (2H) is used in nuclear reactors as a moderator to slow neutrons. Deuterated compounds have applications in chemistry and biology in studies of isotope effects in molecules. Tritium (3H), produced in nuclear reactors, is used in hydrogen bombs.
Liquid hydrogen is important in cryogenics and in the study of superconductivity, as its boiling point is only 20.28 degrees above absolute zero.