Lecture 3. Hydrogen.

In 1671, Robert Boyle discovered and described the reaction between iron filings and dilute acids, which resulted in the production of hydrogen gas. In 1766-81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, and that it produced water when burned. He named it "flammable air". In 1783, Antoine Lavoisier gave the element the name hydrogen (from the Greek υδρο- hydro meaning "water" and -γενης genes meaning "creator") when he and Pierre-Simon Laplace reproduced Cavendish's finding that water was produced when hydrogen was burned.
Hydrogen was liquefied for the first time by James Dewar in 1898 by using regenerative cooling and his invention, the vacuum flask. He produced solid hydrogen the next year. Deuterium was discovered in December 1931 by Harold Urey, and tritium was prepared in 1934 by Ernest Rutherford, Mark Oliphant, and Paul Harteck. Heavy water, which consists of deuterium in the place of regular hydrogen, was discovered by Urey's group in 1932.

The nickel hydrogen battery was used for the first time in 1977 aboard the U.S. Navy's Navigation technology satellite-2 (NTS-2). It had two caesium atomic clocks on board and helped to show that satellite navigation based on precise timing was possible. In the dark part of its orbit, the Hubble Space Telescope is powered by nickel-hydrogen batteries, which were finally replaced in May 2009, more than 19 years after launch, and 13 years passed their design life.

Hubble Space Telescope
from NASA (accessed 2 Feb 2015)

Isotopes of hydrogen
Hydrogen has three naturally occurring isotopes, denoted 1H, 2H and 3H. Other, highly unstable nuclei (4H to 7H) have been synthesized in the laboratory but are not observed in nature.
Hydrogen is the only element that has different names for its isotopes in common use today. During the early study of radioactivity, various heavy radioactive isotopes were given their own names, but these names are no longer used, except for deuterium and tritium.

nuclide symbol Z(p) N(n) isotopic mass (u) half-life decay mode Daughter Isotope representative isotopic composition
1H 1 0 1.00782503207(10) Stable     0.999885(70)
2H - D 1 1 2.0141017778(4) Stable     0.000115(70)
3H - T 1 2 3.0160492777(25) 12.32(2) y β 3He <1 in 1017 atoms

Properties of hydrogen

The difference of mass between isotopes of most elements is only a small fraction of the total mass and so this has very little effect on their properties, this is not the case for hydrogen. Consider chlorine with Z=17, there are 2 stable isotopes 35Cl (75.77%) and 37Cl (24.23%). The increase is therefore 2 in 35 or less than 6%. Deuterium and tritium are about double and triple the mass of protium and show significant physical and chemical differences particularly with regard to those properties related to mass, e.g. rate of diffusion, density, etc.

Some physical properties of the hydrogen isotopes.
isotope MP /K BP /K ΔHdiss /kJmol-1 Interatomic Distance /pm
H2 13.99 20.27 435.99 74.14
D2 18.73 23.67 443.4 74.14
T2 20.62 25.04 446.9 74.14

Differences between H2O and D2O

Property H2O D2O
Melting point /K 273.15 276.97
Boiling point /K 373.15 374.5
Temperature of maximum density /K 277 284.2
Maximum density /g cm3 0.99995 1.1053
Relative permittivity (at 298 K) 78.39 78.06
Kw (at 298 K) 1 *1014 2 * 1015
Symmetric stretch, ν1 /cm-1
(gaseous molecule)
3657 2671

Given that the boiling point of D2O is 101.4 °C (compared to 100.0 °C for H2O), evaporation or fractional distillation can be used to increase the concentration of deuterium in a sample of water by the selective removal of the more volatile light water, H2O. Thus bodies of water that have no outlet, such as the Great Salt Lake in Utah, USA and the Dead Sea in the Jordan Rift Valley, which maintain their level solely by evaporation, have significantly higher concentrations of deuterated water than do lakes or seawater with at least one outlet.

The isotopic ratio for H and D is not fixed and so a range is given
for the standard atomic weight in the IUPAC Periodic Table of isotopes.

Heavy water is 10.6% denser than ordinary water, a difference not immediately obvious since they are otherwise physically and chemically similar. The difference can be observed by freezing a sample and dropping it into normal water, where it sinks.

With respect to taste and smell, rats given a choice between distilled normal water and heavy water avoided the heavy water, based on smell, and it may be that they detected a different taste as well.

The difference in weight increases the strength of water's hydrogen-oxygen bonds, and this in turn is sufficient to cause differences that are important to some biochemical reactions. The human body naturally contains deuterium equivalent to about five grams of heavy water, which is harmless. When a large fraction of water (> 50%) in higher organisms is replaced by heavy water, the result is cell dysfunction and death.

In normal water, about 1 molecule in 3,200 is HDO (one hydrogen in 6,400 is in the form of D), and heavy water molecules (D2O) only occur in a proportion of about 1 molecule in 41 million (i.e. one in 6,4002). Thus semiheavy water molecules are far more common than "pure" (homoisotopic) heavy water molecules.

Deuterium oxide was initially obtained by the electrolysis of ordinary water over a considerable period of time. This method of production requires a large cascade of stills or electrolysis chambers and consumes large amounts of power, so that chemical methods are generally now preferred. The most important chemical method is the Girdler sulfide process.

Girdler sulfide process

In this process, demineralised and deaerated water is trickled through a series of perforated (seive) plates in a tower, while hydrogen sulfide gas (BP -60 °C) flows upward through the perforations. Deuterium migration preferentially takes place from the gas to the liquid water. This "enriched" water from the cold tower (maintained at 32 °C) is fed to the hot tower (at 130 °C) where deuterium transfer takes place from the water to the hydrogen sulfide gas. An appropriate "cascade" setup accomplishes enrichment via the reversible reaction:

H2O +HDS ⇄ HDO + H2S

The equilibrium constant, K for, this reaction in terms of the concentrations, can be written as:

K = ([HDO][H2S]) / ([H2O][HDS]) or alternatively:

K = ([HDO]/[H2O]) / ([HDS]/[H2S])

If H and D were the same chemically, the equilibrium constant for the reaction would be equal to unity. However, what is found is that K is not equal to unity, and furthermore it is temperature dependent:

at 25 °C,  K = 2.37
at 128 °C, K = 1.84

From the above information, at 32 °C, the equilibrium favours the concentration of deuterium in water. However, at around 130 °C, the equilibrium is now relatively more favorable to the concentration of deuterium in the hydrogen sulfide. In other words, the concentration of HDO in H2O is greater than the concentration of HDS in H2S but the relative concentration of HDS in H2S increases with increasing temperature, making it possible to separate D from H.

In the first stage, the gas is enriched from 0.015% deuterium to 0.07%. The second column enriches this to 0.35%, and the third column achieves an enrichment between 10% and 30% deuterium oxide, D2O. Further enrichment to "reactor-grade" heavy water (> 99% D2O) still requires distillation or electrolysis. The production of a single litre of heavy water requires ~340,000 litre of feed water.

In 1934, Norway built the first commercial heavy water plant with a capacity of 12 tonnes per year. From 1940 and throughout World War II, the plant was under German control and the Allies decided to destroy the plant and its heavy water to inhibit German development of nuclear weapons. In late 1942, a planned raid by British airborne troops failed, both gliders crashing. The raiders were killed in the crash or subsequently executed by the Germans. On the night of 27 February 1943 Operation Gunnerside succeeded. Norwegian commandos and local resistance managed to demolish small, but key parts of the electrolytic cells, dumping the accumulated heavy water down the factory drains. Had the German nuclear program followed similar lines of research as the United States Manhattan Project, the heavy water would not have been crucial to obtaining plutonium from a nuclear reactor, but the Germans did not discover the graphite reactor design used by the allies for this purpose.

ortho- and para-dihydrogen

In dihydrogen, the two electrons in the molecule will be spin paired but there is no similar requirement for the two nuclei; they may be parallel or opposed. There are therefore two nuclear spin isomers possible which are called ortho and para.

para and ortho dihydrogen

In the parahydrogen form the nuclear spins of the two protons are antiparallel and form a singlet (2I+1= 1) with a molecular spin quantum number, I, of 0 (½ - ½). In the orthohydrogen form, the spins are parallel and form a triplet state (2I+1= 3) with a molecular spin quantum number, I, of 1 (½ + ½). At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form, also known as the "normal form".

The para form has slightly lower energy:-
o-H2 ⇄ p-H2; ΔH = -1.5 kJmol-1
but due to the small difference this has little effect at room temperature.

The amount of ortho and para hydrogen varies with temperature:
It is possible then to get pure para hydrogen by cooling ordinary hydrogen gas to very low temperatures (close to 20 K) but it is not possible to get a sample of hydrogen containing more than 75% of ortho (triplet) hydrogen. The first synthesis of pure parahydrogen was achieved in 1929.

This conversion of ortho- to para-hydrogen liberates some heat which can cause evaporation of hydrogen within storage vessels. Since orthohydrogen molecules make up 75% of "normal" hydrogen at room temperature, this can considerably complicate the performance of storing liquid hydrogen. Without an ortho-para conversion catalyst, (such as hydrous ferric oxide) extra refrigeration equipment is required to remove the heat generated by the natural conversion to para hydrogen.

Production of hydrogen

Laboratory preparations
In the laboratory, H2 can be prepared by the action of a dilute non-oxidizing acid on a reactive metal such as zinc, with a Kipp's apparatus.
Zn + 2Haq+ ⇄ Znaq2+ + H2

Aluminium can produce H2 upon treatment with bases:
2Al + 6 H2O + 2 OH- ⇄ 2 Al(OH)4- + 3 H2

The electrolysis of water is another simple method of producing hydrogen. A low voltage current is passed through the water, and gaseous dioxygen forms at the anode while gaseous hydrogen forms at the cathode. Typically the cathode is made from platinum or other inert metal when producing hydrogen for storage. If, however, the gas is to be burnt on site, oxygen is desirable to assist the combustion, and so both electrodes would be made from inert metals. (Iron, for instance, would oxidize, and thus decrease the amount of oxygen given off.) The theoretical maximum efficiency (electricity used versus energetic value of hydrogen produced) is in the range 80-94%.
2 H2O(l) ⇄ 2 H2(g) + O2(g)

In 2007, it was discovered that an alloy of aluminium and gallium in pellet form added to water could be used to generate hydrogen. The process creates alumina, but the expensive gallium, which prevents the formation of an oxide skin on the pellets, can be re-used. This has important potential implications for a hydrogen economy, as hydrogen could be produced on-site without the need of being transported.

Industrial preparation of hydrogen
Steam reforming is a method for producing hydrogen, carbon monoxide or other useful products from hydrocarbon fuels such as natural gas. This is achieved in a processing device called a reformer which reacts steam at high temperature with the fossil fuel.

At high temperatures (700 - 1100 °C) and in the presence of a metal-based catalyst (nickel), steam reacts with methane to yield carbon monoxide and hydrogen.

CH4 + H2O → CO + 3 H2

In order to produce more hydrogen from this mixture, more steam is added and the water gas shift reaction is carried out:
CO + H2O → CO2 + H2

The mixture of CO and H2 is called "synthesis gas or syngas". Syngas is used as an intermediate in producing synthetic petroleum for use as a fuel or lubricant via the Fischer-Tropsch process and previously the Mobil methanol to gasoline process.

Enzymatic route from xylose
In 2013 a low-temperature, 50 °C, atmospheric-pressure, enzyme-driven process to convert xylose into hydrogen with nearly 100% of the theoretical yield was announced. The process employed 13 enzymes, including a novel polyphosphate xylulokinase (XK).

It was noted that: "Approximately 50 million metric tons of dihydrogen are produced annually from nonrenewable natural gas, petroleum, and coal. H2 production from water remains costly. Technologies for generating H2 from less costly biomass, such as microbial fermentation, enzymatic decomposition, gasification, steam reforming, and aqueous phase reforming, all suffer from low product yields."

Compounds of Hydrogen

The chemistry of hydrogen depends mainly on four processes:
  1. donation of the valency electron to form the hydrogen ion, H+
  2. accepting an electron to form the hydride ion H-
  3. sharing the electron with a partner atom to form a pair bond (covalent bond) H-H
  4. sharing the electron with an ensemble of atoms to form a metallic bond H.
While H2 is not very reactive under standard conditions, it does form compounds with most elements. Hydrogen can form compounds with elements that are more electronegative, such as halogens (e.g., F, Cl, Br, I), or oxygen; in these compounds hydrogen takes on a partial positive charge. When bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a form of medium-strength noncovalent bonding called hydrogen bonding, which is critical to the stability of many biological molecules. Hydrogen also forms compounds with less electronegative elements, such as the metals and metalloids, in which it takes on a partial negative charge. These compounds are often known as hydrides.

The term "hydride" suggests that the H atom has acquired a negative or anionic character, denoted H-, and is used when hydrogen forms a compound with a more electropositive element. The existence of the hydride anion, suggested by Gilbert N. Lewis in 1916 for group I and II salt-like hydrides, was demonstrated by Moers in 1920 by the electrolysis of molten lithium hydride (LiH), producing a stoichiometry quantity of hydrogen at the anode.

Although hydrides can be formed with almost all main-group elements, the number and combination of possible compounds varies widely; for example, there are over 100 binary borane hydrides known, but only one binary aluminium hydride. A simple binary indium hydride has not yet been identified, although larger complexes exist.

The position of H in the Periodic Table

In some respects, H does not seem to have a perfect position in the Periodic Table and so many designers have it in more than one position, e.g. in Group 1 or Group 17 and even in Group 14.

Ionization energy of hydrogen

Hydrogen has a single outer electron, like the alkali metals, but they all form positive ions quite readily whereas hydrogen has little tendency to do so. Hydrogen often tends to share its electron with nonmetals rather than losing it to them.

The first ionization energies for H, Li, Na and K are 1312, 520.2, 495.8 and 418.8 kJmol-1. The high IE for H (even bigger than for Xe) can be attributed to the very small size of the atom and the strong attractive force between the proton and electron.

H(g) → H+(g) + e        ΔH = 1312 kJmol-1

Xe(g) → Xe+(g) + e      ΔH = 1170 kJmol-1

The free proton can only be obtained under extreme conditions such as by an electric arc or in a discharge tube and even then only exists for about half a second. H+ can be found in solvated form where the solvation energy provides the energy needed to overcome the very high ionization energy. Examples are in ammonia, alcohol or water with species like NH4+, ROH2+ and H3O4+ being formed.

The hydrated proton (H3O4+) will be covered in Lecture 6 on acid-base chemistry.

Electron affinity of hydrogen

Hydrogen, like the halogens, exists as diatomic molecules and H atoms have electron configurations with one electron short of a filled outer shell hence the idea of placing H in Group 17. However unlike the halogens with large EA values, the EA for hydrogen is quite small. The formation of H- is much less favourable than the formation of a chloride ion, as seen from the thermodynamic profiles below and it is rare whereas halide ions are common and stable. In addition H has a lower electronegativity value than any of the halogens.

comparison of EA for H and Cl

Much more energy is required as well to break the H-H bond compared to the Cl-Cl bond where the steps for comparison are:
½H2 (g) → H. (g)          ΔH = 218 kJmol-1
H. (g) + e → H- (g)          ΔH = -72.8 kJmol-1
so overall for hydrogen
½H2 (g) + e → H- (g)          ΔH = +145.2 kJmol-1
½Cl2 (g) → Cl. (g)          ΔH = 121 kJmol-1
Cl. (g) + e → Cl- (g)          ΔH = -348.6 kJmol-1
overall for chlorine
½Cl2 (g) + e → Cl- (g)          ΔH = -227.6 kJmol-1

As a result, only the most active elements, whose Ionization Energies are low, can form ionic hydrides, e.g. NaH.

The covalent radius for H is 37 pm and the estimated radius for H- is ~140 pm indicating a substantial increase. This comes about as a result of the interelectronic repulsion when a second electron is added to the 1s atomic orbital. All the alkali metal hydrides crystallize with the NaCl-type structure and are all considered ionic. They are sometimes called "saline" hydrides.

Saline hydrides

The instability of the hydride ion compared to the halide ions can be seen by comparison of the ΔHf for alkali metal hydrides and chlorides.

Cation ΔHf MH
/ kJmol-1
/ kJmol-1
Li -90.5 -409
Na -56.3 -411
K -57.7 -436
Rb -52.3 -430
Cs -54.2 -433

Saline hydrides are formed by the group 1 and 2 metals when heated with dihydrogen (H2). They are white, high melting point solids that react immediately with protic solvents, for example:

NaH + H2O → NaOH + H2

(Their moisture sensitivity means that reaction conditions must be water-free.)

Evidence for the ionic nature of these hydrides is:
1) molten salts show ionic conductivity.
2) X-ray crystal data gives reasonable radius ratios expected for ionic compounds.
3) Observed and calculated Lattice Energies (from Born-Haber cycles etc.) are in good agreement (i.e. show little covalency).

NaH is capable of deprotonating a range of even weak Brønsted acids to give the corresponding sodium derivatives.

NaH + Ph2PH → Na[PPh2] + H2

Sodium hydride is sold by many chemical suppliers as a mixture of 60% sodium hydride (w/w) in mineral oil. Such a dispersion is safer to handle and weigh than pure NaH. The compound can be used in this form but the pure grey solid can be prepared by rinsing the oil with pentane or tetrahydrofuran, THF, care being taken because the washings will contain traces of NaH that can ignite in air. Reactions involving NaH require an inert atmosphere, such as nitrogen or argon gas. Typically NaH is used as a suspension in THF, a solvent that resists deprotonation but solvates many organosodium compounds.

Hydride reducing agents

LiH and Al2Cl6 gives lithium aluminium hydride (lithal LiAlH4), NaH reacts with B(OCH3)3 to give sodium borohydride (NaBH4). These find wide scope and utility in organic chemistry as reducing agents.

LiAlH4 is commonly used for the reduction of esters and carboxylic acids to primary alcohols; previously this was a difficult conversion that used sodium metal in boiling ethanol (the Bouveault-Blanc reduction). The solid is dangerously reactive toward water, releasing gaseous hydrogen (H2). Some related derivatives have been discussed for hydrogen storage.

NaBH4 is used in large amounts for the production of sodium dithionite from sulfur dioxide: Sodium dithionite is used as a bleaching agent for wood pulp and in the dyeing industry. NaBH4 consists of the tetrahedral BH4- anion in the crystalline form and is found to exist as three polymorphs: α, β and γ. The stable phase at room temperature and pressure is α-NaBH4, which is cubic and adopts an NaCl-type structure. Millions of kilograms are produced annually, far exceeding the production levels of any other hydride reducing agent.

NaBH4 will reduce many organic carbonyls, depending on the precise conditions. Most typically, it is used in the laboratory for converting ketones and aldehydes to alcohols. For example, reduction of acetone (propanone) to give propan-2-ol.
reaction of acetone to propan-2-ol

Molecular hydrides - covalent hydrides and organic compounds

Hydrogen forms a vast number of compounds with carbon, (the hydrocarbons), and an even larger array with heteroatoms that, because of their general association with living things, are called organic compounds. The study of their properties is covered in organic chemistry and their study in the context of living organisms is covered in biochemistry. By some definitions, "organic" compounds are only required to contain carbon. However, most of them also contain hydrogen, and because it is the carbon-hydrogen bond which gives this class of compounds most of its particular chemical characteristics, carbon-hydrogen bonds are required in some definitions of the word "organic" in chemistry. Millions of hydrocarbons are known, and they are usually formed by complicated synthetic pathways, which seldom involve direct reaction with elementary hydrogen.

Most molecular hydrides are volatile and many have simple structures that can be predicted by the VSEPR model. There are a large number of B hydrides known (boranes) and although the simplest BH3 has been found in the gas phase it readily dimerises to give B2H6.

In inorganic chemistry, hydrides can serve as bridging ligands that link two metal centers in a coordination complex. This function is particularly common in group 13 elements, especially in boranes (boron hydrides) and aluminium complexes, as well as in clustered carboranes, (composed of boron, carbon and hydrogen atoms). The bonding of the bridging hydrogens in many of the boranes is explained in terms of 3 centre - 2 electron bonds.

Diborane is a colourless and highly unstable gas at room temperature with a repulsively sweet odour. Diborane mixes well with air, easily forming explosive mixtures. Diborane will ignite spontaneously in moist air at room temperature.
B2H6 3 centre 2 electron model
MP = -164.85 °C, BP= -92.5 °C
B-H (terminal) 119 pm, (bridge) 131 pm

Metallic (interstitial) hydrides

Many transition metal elements form metallic (interstitial) hydrides, in which H2 molecules (and H atoms) can occupy the holes in the metal's crystal structure. They are traditionally termed 'compounds', even though they do not strictly conform to the definition of a compound; more closely resembling common alloys such as steel. These systems are usually non-stoichiometric, with variable amounts of hydrogen atoms in the lattice.

Palladium is unique in its ability to reversibly absorb large amounts of H2 or D2 (up to 900 times its own volume of hydrogen, but no other gases, at room temperature) to form palladium hydride. Structural studies show that the absorbed H fits into octahedral holes in the cubic close packed Pd lattice with a non-stoichiometric formula approximating to PdH0.6 for the β-form. This material has been considered as a means to carry hydrogen for vehicular fuel cells. Interstitial hydrides show some promise as a way for safe hydrogen storage. During the last 25 years many interstitial hydrides have been developed that readily absorb and discharge hydrogen at room temperature and atmospheric pressure. At this stage their application is still limited, as they are capable of storing only about 2 weight percent of hydrogen, insufficient for automotive applications.
H2 with Pd - interstitial PdH0.6

Hydrogen bonds

A hydrogen bond is the name given to the electrostatic attraction between polar molecules that occurs when a hydrogen (H) atom bound to a highly electronegative atom such as nitrogen (N), oxygen (O) or fluorine (F) experiences attraction to some other nearby highly electronegative atom. The name is something of a misnomer, as it represents a particularly strong dipole-dipole attraction, rather than a typical covalent bond.

The 2011 IUPAC definition specifies that "The hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X-H in which X is more electronegative than H, and an atom or a group of atoms in the same or a different molecule, in which there is evidence of bond formation."

These hydrogen-bond attractions can occur between molecules (intermolecular) or within different parts of a single molecule (intramolecular). The hydrogen bond (5 to 30 kJ/mole) is stronger than a van der Waals interaction, but weaker than covalent or ionic bonds. This type of bond can occur in inorganic molecules such as water and in organic molecules like DNA and proteins.

Intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared to the other group 16 hydrides that have no hydrogen bonds. Intramolecular hydrogen bonding is partly responsible for the secondary and tertiary structures of proteins and nucleic acids. It plays an important role in the structure of polymers, both synthetic and natural.

BP's of MG hydrides
BP's of MG hydrides with Noble gases for comparison /K

Hydrogen bonding in biological systems.

Base pairs, which form between specific nucleobases (also termed nitrogenous bases), are the building blocks of the DNA double helix and contribute to the folded structure of both DNA and RNA. Dictated by specific hydrogen bonding patterns, Watson-Crick base pairs (guanine-cytosine and adenine-thymine) allow the DNA helix to maintain a regular helical structure that is subtly dependent on its nucleotide sequence. The complementary nature of this based-paired structure provides a backup copy of all genetic information encoded within double-stranded DNA. The regular structure and data redundancy provided by the DNA double helix make DNA well suited to the storage of genetic information, while base-pairing between DNA and incoming nucleotides provides the mechanism through which DNA polymerase replicates DNA, and RNA polymerase transcribes DNA into RNA. Many DNA-binding proteins can recognize specific base pairing patterns that identify particular regulatory regions of genes.

Hydrogen bonding in base pairs
GC base pairs AT base pairs

Applications of hydrogen

Large quantities of H2 are used by the petroleum and chemical industries. The largest application of H2 is for the processing ("upgrading") of fossil fuels, and in the production of ammonia. The key consumers of H2 in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking. H2 has several other important uses. H2 is used as a hydrogenating agent, particularly in increasing the level of saturation of unsaturated fats and oils (found in items such as margarine), and in the production of methanol. It is similarly the source of hydrogen in the manufacture of hydrochloric acid. H2 is used as a reducing agent of metallic ores.

Nitrogen is a strong limiting nutrient in plant growth. Carbon and oxygen are also critical, but are more easily obtained by plants from soil and air. Even though air is 78% nitrogen, atmospheric nitrogen is nutritionally unavailable because nitrogen molecules are held together by strong triple bonds. Nitrogen must be 'fixed', i.e. converted into some bioavailable form, through natural or man-made processes. It was not until the early 20th century that Fritz Haber developed the first practical process to convert atmospheric nitrogen to ammonia, which is nutritionally available.

Fertilizer generated from ammonia produced by the Haber process is estimated to be responsible for sustaining one-third of the Earth's population. It is estimated that half of the protein within human beings is made of nitrogen that was originally fixed by this process; the remainder was produced by nitrogen fixing bacteria and archaea.

Dozens of chemical plants worldwide produce ammonia, consuming more than 1% of all man-made power. Ammonia production is thus a significant component of the world energy budget. Modern ammonia-producing plants depend on industrial hydrogen production to react with atmospheric nitrogen using a magnetite catalyst or over a promoted Fe catalyst under high pressure (100 standard atmospheres (10,000 kPa)) and temperature (450 °C) to form anhydrous liquid ammonia. This step is known as the ammonia synthesis loop (also referred to as the Haber-Bosch process):

3 H2 + N2 ⇄ 2 NH3 (ΔH = -92.4 kJmol-1)

Nitrogen (N2) is very unreactive because the molecules are held together by strong triple bonds. The Haber process relies on catalysts that accelerate the cleavage of this triple bond.

At room temperature, the equilibrium is strongly in favor of ammonia, but the reaction doesn't proceed at a detectable rate. Thus two opposing considerations are relevant to this synthesis. One possible solution is to raise the temperature, but because the reaction is exothermic, the equilibrium quickly becomes quite unfavourable at atmospheric pressure. Low temperatures are not an option since the catalyst requires a temperature of at least 400 °C to be efficient. By increasing the pressure to around 200 atm the equilibrium concentrations are altered to give a profitable yield.

The reaction scheme, involving the heterogeneous catalyst, is believed to involve the following steps:

  1. N2 (g) → N2 (adsorbed)
  2. N2 (adsorbed) → 2 N (adsorbed)
  3. H2(g) → H2 (adsorbed)
  4. H2 (adsorbed) → 2 H (adsorbed)
  5. N (adsorbed) + 3 H (adsorbed) → NH3 (adsorbed)
  6. NH3 (adsorbed) → NH3 (g)
Energy Profile for Haber process

Reaction 5 actually consists of three steps, forming NH, NH2, and then NH3. Experimental evidence suggests that reaction 2 is the slow, rate-determining step. This is not unexpected given that the bond broken, the nitrogen triple bond, is the strongest of the bonds that must be broken.

Return to the course outline or move on to Lecture 4: Oxygen and its compounds.


Much of the information in these course notes has been sourced from Wikipedia under the Creative Commons License.
DNA modelling
'Inorganic Chemistry' - C. Housecroft and A.G. Sharpe, Prentice Hall, 4th Ed., 2012, ISBN13: 978-0273742753, pps 24-27, 43-50, 172-176, 552-558, 299-301, 207-212
'Basic Inorganic Chemistry' - F.A. Cotton, G. Wilkinson and P.L. Gaus, John Wiley and Sons, Inc. 3rd Ed., 1994.
'Introduction to Modern Inorganic Chemistry' - K.M. Mackay, R.A. Mackay and W. Henderson, International Textbook Company, 5th Ed., 1996.

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