Lecture 4. Oxygen

Oxygen is a member of the chalcogen group (Group 16) of the periodic table and is a highly reactive nonmetallic element and oxidizing agent that readily forms compounds (notably oxides) with most elements. By mass, oxygen is the third-most abundant element in the universe, after hydrogen and helium. At standard temperature and pressure, STP, two atoms of the element bind to form dioxygen, O2, a diatomic gas that is colourless, odourless, and tasteless.

Elemental abundance


Oxygen was discovered independently by Carl Wilhelm Scheele, in Uppsala, Sweden in 1773 or earlier, and Joseph Priestley in Wiltshire, UK in 1774, but Priestley is generally given the credit since his work was published first. The name oxygen was coined in 1777 by Antoine Lavoisier, whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. Its name derives from the Greek roots οξυζ oxys, "acid", literally "sharp", referring to the sour taste of acids and -γενης genes meaning "creator", because at the time of naming, it was mistakenly thought that all acids required oxygen in their composition.

Allotropes of oxygen

There are 2 main allotropes of oxygen although a third has recently been shown to form under high pressures.

Some allotropes of oxygen
O2 - dioxygen O3 - ozone O8 - red oxygen

Dioxygen - O2

The common allotrope of elemental oxygen is often just called oxygen, O2, but to help distinguish it from the element may be called dioxygen or molecular oxygen. Elemental oxygen is most commonly encountered in this form, as about 21% (by volume) of the Earth's atmosphere is O2, the remainder largely being dinitrogen, N2. At STP, dioxygen is a colourless, odourless gas, in which the two oxygen atoms are chemically bonded to each other giving rise to two unpaired electrons occupying two degenerate molecular orbitals. The electron configuration of O2 molecules in this form, a diradical, indicates that they should be paramagnetic. This is a classic example of where a simple Lewis structure fails to account for the properties and where an MO approach correctly provides the explanation.

Lewis structure for dioxygen              MO diagram for dioxygen

The electron configuration (ignoring 1s orbitals) is:
2s)22s*)22p)22p)42p*)2 and from this the Bond Order is found to be ½(2 -2 + 2 + 4 -2 ) = 2 that is, a double bond as shown in the Lewis structure as well. The difference though is that the Lewis structure does not predict the molecule to be paramagnetic. The bond length is 121 pm and the bond energy is 498 kJmol-1.

A video clip showing liquid dioxygen being poured between the faces of a magnet and attracted into the magnet field has been prepared as a Harvard Natural Sciences Lecture Demonstration.

With 2 electrons to be placed in 2 degenerate orbitals, a number of variations are possible and the arrangement above where the 2 electrons are parallel is considered to be the most stable. Note that the spin multiplicity is given by the formula, 2S+1 and so for S=1 from s=½ + ½ then 2S+1 = 3 i.e. a spin triplet.

MO diagram for 2e 2o
a) e-s parallel (triplet 3Σg),     b) 2 e-s in 1 orbital (singlet 1Δg),    c) e-s opposed (singlet 1Σg)

Singlet oxygen is the name commonly used for the electronically excited state shown in b) and it is less stable than the normal triplet state a) by 94.7 kJmol-1. In isolation, singlet oxygen can persist for over an hour at room temperature. The other singlet state at 157.8 kJmol-1 shown in c) is very short lived and relaxes quickly to b). Because of differences in their electron shells, singlet and triplet oxygen differ in their chemical properties. Singlet oxygen is highly reactive.

Reactions of triplet dioxygen are restricted by conservation of spin state rules and at ambient temperatures this prevents direct reaction with all but the most reactive substrates, e.g. white phosphorus. At higher temperatures, or in the presence of suitable catalysts, reactions proceed more readily. For instance, most flammable substances are characterised by an autoignition temperature above which they will undergo combustion in air without an external flame or spark.

The energy difference between the ground state and singlet oxygen is 94.7 kJmol-1 which would correspond to a transition in the near-infrared at ~1263 nm. In the isolated molecule, this transition is strictly forbidden by spin, symmetry and parity selection rules, making it one of nature's most forbidden transitions. In other words, direct excitation of ground state oxygen by light to form singlet oxygen is very improbable. As a consequence, singlet oxygen in the gas phase is extremely long lived (72 minutes). Interaction with solvents however, can reduce the lifetime to microseconds or even nanoseconds.

Various methods for the production of singlet oxygen exist. A photochemical method involves the irradiation of normal oxygen gas in the presence of an organic dye as a sensitizer, such as methylene blue. Singlet oxygen can be produced chemically as well. One of the chemical methods is by the reaction of hydrogen peroxide with sodium hypochlorite. This is convenient in small laboratories and for demonstrative purposes:

H2O2 + NaOCl → O2(1Δg) + NaCl + H2O

In photosynthesis, singlet oxygen can be produced from the light-harvesting chlorophyll molecules. One of the roles of carotenoids in photosynthetic systems is to prevent damage caused by any singlet oxygen that is produced by either removing excess light energy from chlorophyll molecules, or quenching the singlet oxygen molecules directly.

Biological considerations

Molecular dioxygen is a potentially strong oxidizing agent, based on its position in the electrochemical series. The standard redox potential is:
O2 + 4H+ + 4e- ⇄ 2 H2O    E = +1.229 V

which is comparable to potassium dichromate (1.33 V). Nevertheless, reactions of dioxygen with most substrates tend to proceed very slowly at room temperature, in the gas phase or in solution. This relative inertness (unexpected considering it is a diradical) is critical for sustaining life in a dioxygen atmosphere. Perhaps one of the reasons for this stablilty is that the reaction above shows that a 4 electron change is required and this is highly improbable so that it is more likely that a sequence of steps involving 1 electron changes actually takes place.

O2 + e- → O2-* (superoxide)
O2-* + e- + 2H+ → H2O2 (peroxide)
H2O2 + e- + H+ → OH* (hydroxy radical)
OH* + e- + H+ → H2O (water)

The first step, that generates superoxide, has an adverse potential of -0.32 V (ΔG= +30.9 kJmol-1) which imparts some kinetic stability to the molecule. Note that both dioxygen and superoxide ion are free radicals that exhibit paramagnetism. Species like superoxide, peroxide and the hydroxy radical are far too reactive to be allowed to accumulate in living systems and the primary defence of the cell appears to be to destroy them as soon as they are formed. This is done by a variety of enzymes including superoxide dismutase (SOD) and catalases. Parts of the immune system of higher organisms, however, create peroxide, superoxide, and singlet oxygen to destroy invading microbes. Reactive oxygen species also play an important role in the hypersensitive response of plants against pathogen attack.

Another issue with dioxygen is its low solubility in water at room temperature and atmospheric pressure. Dioxygen is more soluble in water than is dinitrogen. Water in equilibrium with air contains approximately 1 molecule of dissolved O2 for every 2 molecules of N2, despite the atmospheric ratio of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mgL-1) dissolves at 0 °C than at 20 °C (7.6 mgL-1). At 25 °C and 1 standard atmosphere (101.3 kPa) of air, freshwater contains about 6.04 milliliters (mL) of dioxygen per liter, whereas seawater contains about 4.95 mgL-1. At 5 °C the solubility increases to 9.0 mgL-1 (50% more than at 25 °C) for water and 7.2 mgL-1 (45% more) for sea water.

The oxygen in water is unavailable to mammals so that divers (and diving mammals such as whales and seals) are entirely dependent on the oxygen carried in the air in their lungs or their gas supply. For divers this is complicated since at higher partial pressures oxygen can cause acute toxicity leading to convulsions. Scuba (self-contained underwater breathing apparatus) divers using compressed air are restricted to diving above 30 m. If the diver descends to depths near 100 m they can become unconsciousness from nitrogen narcosis and death usually results. A dive to 30 m for 20 minutes puts the scuba diver at risk of nitrogen narcosis and decompression illness. By contrast, the elephant seal can dive to 1 km for 1 hour without risk of either condition.

Respiratory pigments are capable of fixing dioxygen from the atmosphere, transporting it to the reacting site and then releasing it. In addition they are able to counteract small fluctuations in supply or demand by storing dioxygen.

O2 binding to metals eg in Myoglobin
O2 can bind to a single metal center either "end-on" (η1-) or "side-on" (η2-). In Haemoglobin and Myoglobin it is "end-on".

fully oxygenated Haemoglobin
(with 4 Fe porphins and O2's highlighted)

Iron based proteins like haemoglobin and myglobin as well as copper based proteins like haemocyanin perform this role in mammalian systems and lobsters, etc.

Molecular dioxygen, O2, is essential for cellular respiration in all aerobic organisms. Oxygen is used in mitochondria to help generate adenosine triphosphate (ATP) during oxidative phosphorylation. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as:

C6H12O6 + 6 O2 → 6 CO2 + 6 H2O      ΔH= -2880 kJmol-1

In vertebrates, O2 diffuses through membranes in the lungs and into red blood cells. Haemoglobin binds O2, changing its colour from bluish red to bright red (CO2 is released from another part of haemoglobin through the Bohr effect). Other animals use haemocyanin (molluscs and some arthropods) or haemerythrin (spiders and lobsters). A liter of blood can dissolve 200 cm3 of O2.

Ozone - O3

Ozone is colourless or slightly bluish gas (blue when liquified), slightly soluble in water and much more soluble in inert non-polar solvents such as carbon tetrachloride or fluorocarbons, where it forms a blue solution. At 161 K (-112 °C), it condenses to form a dark blue liquid. It is dangerous to allow this liquid to warm to its boiling point, because both concentrated gaseous ozone and liquid ozone can detonate.

Most people can detect about 0.01 μmol/mol of ozone in air where it has a very specific sharp odour somewhat resembling chlorine bleach. Exposure of 0.1 to 1 μmol/mol produces headaches, burning eyes and irritation to the respiratory passages. Even low concentrations of ozone in air are very destructive to organic materials such as latex, plastics and animal lung tissue.

Seaside air was once considered to be healthy because of its "ozone" content. It is now recognised that the smell in reality is released from bacteria following the breakdown of species from rotting seaweed! (dimethyl sulfoxide, DMS). Birds appear to be attracted to the smell since for them it indicates a plankton bloom, and therefore the presence of fish feeding on the marine plants.

Ozone may be formed from O2 by electrical discharges and by action of high energy electromagnetic radiation. Unsuppressed arcing breaks down the chemical bonds of the atmospheric oxygen surrounding the contacts [O2 → 2O]. Free radicals of oxygen in and around the arc recombine to create ozone [O3]. Certain electrical equipment generate significant levels of ozone. This is especially true of devices using high voltages, such as ionic air purifiers, laser printers, photocopiers, tasers and arc welders. Electric motors using brushes can generate ozone from repeated sparking inside the unit. Large motors that use brushes, such as those used by elevators or hydraulic pumps, will generate more ozone than smaller motors. Ozone is similarly formed in the Catatumbo lightning storms phenomenon on the Catatumbo River in Venezuela, which helps to replenish ozone in the upper troposphere. It is the world's largest single natural generator of ozone. The lightning originates from a mass of storm clouds at a height of more than 5 km, and occurs during 140 to 160 nights a year, 10 hours per day and up to 280 times per hour.

Lightning at Catatumbo River in Venezuela
From the imgkid gallery (accessed 31 Jan 2015).

Ozone-Oxygen cycle

Ninety percent of the ozone in the atmosphere sits in the stratosphere, the layer of atmosphere between about 10 and 50 kilometers altitude. The ozone layer is mainly found in the lower portion of the stratosphere, from approximately 20 to 30 kilometres (12 to 19 mi) above Earth, though the thickness varies seasonally and geographically. Here the concentration of O3 varies from two to eight ppm which is much larger than the average ozone concentration in the Earth's atmosphere which is only about 0.3 parts per million.

Ozone in the atmosphere

The natural level of ozone in the stratosphere can be considered to be a result of a balance between sunlight that creates ozone and chemical reactions that destroy it.

Creation: photolysis of an oxygen molecule by high energy UV light splits it into two oxygen atoms. The wavelength in the UV needed to achieve this can be estimated from the O=O Bond Energy of 498 kJmol-1.
By dividing this by 6.022 * 1023 bonds/mol the Energy per bond is found to be 8.27 * 10-19 J/bond.
The wavelength required for photolysis can be evaluated from E= hc / λ

λ= hc /E = (6.626 * 10-34 Js   x 3 * 108 m/s) / ( 8.27 * 10-19 J)

so λ= 240 nm, i.e. visible light cannot break the O=O bond, but UV-C light with energy > 240 nm can break the O=O bond.
{UV ranges are: UV-A (400-315 nm), UV-B (315-280 nm), and UV-C (280-200 nm).}

O2 + hν → 2O
Each oxygen atom then rapidly combines with an oxygen molecule to form an ozone molecule:
O + O2 → O3

The ozone-oxygen cycle: the ozone molecules formed by the reaction above absorb radiation again having an appropriate wavelength in the UV. The triatomic ozone molecule becomes diatomic molecular oxygen plus a free oxygen atom:
O3 + hν (240-310 nm) → O2 + O
The atomic oxygen produced quickly reacts with another oxygen molecule to reform ozone:
O + O2 → O3 + K.E.
where "K.E." denotes the excess energy of the reaction which is manifested as extra kinetic energy. These two reactions form the ozone-oxygen cycle, in which the chemical energy released when O and O2 combine is converted into kinetic energy of molecular motion. The overall effect is to convert penetrating UV-B light into heat, without any net loss of ozone. This cycle keeps the ozone layer in a stable balance while absorbing 97-99% of the Sun's medium-frequency ultraviolet light (from about 200 nm to 315 nm wavelength) and preventing it from reaching the Earth's surface, to the benefit of both plants and animals. This reaction is one of two major sources of heat in the stratosphere (the other being the kinetic energy released when O2 is photolyzed into O atoms).

Removal: reaction of an oxygen atom with an ozone molecule leads to production of two oxygen molecules:
O3 + O. → 2 O2
and if two oxygen atoms meet, they can react to form one oxygen molecule:
2 O. → O2
The overall amount of ozone in the stratosphere is determined by a balance between production by solar radiation and removal. The removal rate is slow, since the concentration of O atoms is very low. Certain free radicals, the most important being hydroxyl (OH), nitric oxide (NO) and atoms of chlorine (Cl) and bromine (Br), catalyze the recombination reaction, leading to an ozone layer that is thinner than it would be if the catalysts were not present.

Most of the OH and NO are naturally present in the stratosphere, but human activity, especially emissions of chlorofluorocarbons (CFCs) and halons, has greatly increased the Cl and Br concentrations, leading to ozone depletion. Each Cl or Br atom can catalyze tens of thousands of decomposition reactions before it is removed from the stratosphere.

Follow the Ozone Hole Watch at NASA and see the ozone depletion wikipedia page.

Ozone depletion

Ozone depletion describes two distinct but related phenomena observed since the late 1970s: a steady decline of about 4% per decade in the total volume of ozone in the ozone layer, and a much larger springtime decrease in stratospheric ozone over Earth's polar regions. Note that the Antarctic would be in darkness (no sunshine) for the winter months, sunshine returning in springtime (September). The latter phenomenon is referred to as the ozone hole.

The most important process in depletion is catalytic destruction of ozone by atomic halogens. The main source of these halogen atoms in the stratosphere is photodissociation of man-made halocarbon refrigerants, solvents, propellants, and foam-blowing agents (CFCs, HCFCs, freons, halons), substances that are referred to as ozone-depleting substances (ODS).

In the simplest example of such a cycle, a chlorine atom reacts with an ozone molecule, taking an oxygen atom with it (forming ClO) and leaving a normal oxygen molecule. The chlorine monoxide (ClO) produced can react with a second molecule of ozone to yield another chlorine atom and two molecules of oxygen. These gas-phase reactions can be written such that in the first step a chlorine atom changes an ozone molecule to ordinary oxygen:
Cl. + O3 → ClO + O2
The ClO generated can then destroy a second ozone molecule and recreate the original chlorine atom, which can repeat the first reaction and continue a cycle to destroy ozone.
ClO + O3 → Cl. + 2 O2

In theory, a single chlorine atom could keep on destroying ozone (acting as a catalyst) for up to two years (the time scale for transport back down to the troposphere) were it not for reactions that remove Cl by forming reservoir species such as hydrogen chloride (HCl). In practise, the average chlorine atom reacts with 100,000 ozone molecules before it is removed from the catalytic cycle.

On a per atom basis, bromine is even more efficient than chlorine at destroying ozone, but at present there is much less bromine in the atmosphere. Both chlorine and bromine significantly contribute to the overall depletion of ozone.

The Montreal Protocol on Substances that Deplete the Ozone Layer (a protocol to the Vienna Convention for the Protection of the Ozone Layer) is an international treaty designed to protect the ozone layer by phasing out the production of numerous substances that are responsible for ozone depletion. It was agreed on in 1987, and entered into force on January 1, 1989, with numerous revisions since then.

The bans on the production of CFCs, halons, and other ozone-depleting chemicals such as carbon tetrachloride and trichloroethane have led to the expectation of a recovery of the ozone layer to 1980 values somewhere between 2050 and 2070. This was the estimate given in a summary document of the Scientific Assessment of the Ozone Depletion 2014 published by the UN Environment Programme (UNEP) and the UN World Meterological Organization (WMO).

Among the key findings of the report were that the authors noted that what happens after 2050 will then largely depend on concentrations of CO2, methane and nitrous oxide - the three main long-lived greenhouse gases in the atmosphere.

Ozone hole above the Antarctic 24 Sep 2006
NASA view of the largest ozone hole observed above the Antarctic on 24 Sep 2006.

Ozone is much less stable than the diatomic allotrope O2, breaking down in the lower atmosphere to normal dioxygen. The O - O distances in O3 are 127.2 pm and the O - O - O angle is 116.78°. The bonding can be expressed as a resonance hybrid with a single bond on one side and double bond on the other producing an overall average bond order of 1.5 for each side.

resonance structure of ozone

Ozone is a powerful oxidant (far more so than dioxygen) and has many industrial and consumer applications related to oxidation. This same high oxidizing potential, however, causes ozone to damage mucous and respiratory tissues in animals, and also tissues in plants, above concentrations of about 100 ppb. This makes ozone a potent respiratory hazard and pollutant near ground level.

Ozone oxidizes nitric oxide to nitrogen dioxide:
NO + O3 → NO2 + O2
This reaction is accompanied by chemiluminescence. The NO2 can be further oxidized:
NO2 + O3 → NO3 + O2
The NO3 formed can react with NO2 to form N2O5.

Ozone does not react with ammonium salts, but it oxidizes ammonia to ammonium nitrate:
2 NH3 + 4 O3 → NH4NO3 + 4 O2 + H2O

Ozone reacts with carbon to form carbon dioxide, even at room temperature:
C + 2 O3 → CO2 + 2 O2

Red oxygen - O8

On increasing the pressure of dioxygen at room temperature, the "solid oxygen β-phase" undergoes phase transitions to the δ-phase at 9 GPa and the ε-phase at 10 GPa. Its volume decreases significantly, and it changes colour from blue to deep red. The structure of this red ε-phase, determined in 2006, is based on an O8 cluster (see structure above). It was confirmed that this structure is formed up to 96 GPa. The box-like cluster is an unusual conformation first recognised for oxygen, and as yet has not been experimentally or theoretically reported for any other diatomic molecule.

Liquid oxygen is already used as an oxidant in rockets, and it has been speculated that red oxygen could make an even better oxidant, because of its higher energy density.

Hydrogen peroxide

Hydrogen peroxide (H2O2) is the simplest peroxide (a compound with an oxygen-oxygen single bond) and in its pure form is a colourless liquid that is slightly more viscous than water. It is a strong oxidizer and is used as a bleaching agent and disinfectant. For safety reasons it is normally used as an aqueous solution, also colourless. For consumers, it is usually available from pharmacies at 3 and 6 wt% concentrations. The concentrations are sometimes described in terms of the volume of oxygen gas generated; one milliliter of a 20-volume solution generates twenty milliliters of oxygen gas when completely decomposed. For laboratory use, 100 volume, 30 wt% solutions are the most common. Concentrated H2O2, or 'high-test peroxide' is a reactive oxygen species and has been used as a propellant in rocketry. Diluted H2O2 (between 1.9% and 12%) mixed with ammonium hydroxide has been used to bleach human hair. The chemical's bleaching property lends its name to the phrase "Hollywood peroxide blonde". Hydrogen peroxide can be used for tooth whitening and when mixed with baking soda and salt forms a recipe for home-made toothpaste.


Hydrogen peroxide was first described in 1818 by Louis Jacques Thénard, who produced it by treating barium peroxide with nitric acid. An improved version of this process used hydrochloric acid, followed by addition of sulfuric acid to precipitate the barium sulfate byproduct. Thénard's process was used from the end of the 19th century until the middle of the 20th century.

Pure hydrogen peroxide was long believed to be unstable as early attempts to separate it from the water, which is present during synthesis, all failed. This instability was due to traces of impurities (transition metals salts) which catalyze the decomposition of the hydrogen peroxide. Pure hydrogen peroxide was first obtained in 1894 - almost 80 years after its discovery - by Richard Wolffenstein, who produced it via vacuum distillation.

Determination of the molecular structure of hydrogen peroxide proved to be very difficult. In 1892 the Italian physical chemist Giacomo Carrara (1864-1925) determined its molecular weight by freezing point depression, which confirmed that its molecular formula was H2O2. At least half a dozen hypothetical molecular structures seemed to be consistent with the available evidence. In 1934, the English mathematical physicist William Penney and the Scottish physicist Gordon Sutherland proposed a molecular structure for hydrogen peroxide which was very similar to the presently accepted one and which subsequent evidence cumulatively proved to be correct.

The structure of hydrogen peroxide

H2O2 - gas phase
(O-O 147.4 pm)

H2O2 - solid phase
(O-O 145.8 pm)

Although the O-O bond is a single bond, the molecule has a relatively high barrier to rotation of 29.45 kJmol-1 (2460 cm-1); for comparison, the rotational barrier for ethane is just 12.5 kJmol-1. The increased barrier is thought to be due to the repulsion between the lone pairs of the adjacent oxygen atoms.

The molecular structures of gaseous and crystalline H2O2 are significantly different. This difference is attributed to the effects of hydrogen bonding, which is absent in the gaseous state.

Preparation via the anthraquinone process

The anthraquinone process is a process for the production of hydrogen peroxide, which was developed by BASF. The industrial production of hydrogen peroxide is based on the reduction of oxygen with hydrogen, by the direct synthesis from the elements. Instead of hydrogen itself, however, a 2-alkyl-anthrahydroquinone (generated from the corresponding 2-alkyl-anthraquinone by catalytic hydrogenation with palladium) is used. Oxygen and the organic phase react to give hydrogen peroxide and the anthraquinone can be recycled.

H2O2 prep via anthraquinone

The hydrogen peroxide is then extracted with water in a second step and separated by fractional distillation. The anthraquinone thus acts as a catalyst with the overall reaction given by the equation:

H2 + O2 → H2O2

If ozone is used instead of oxygen, dihydrogen trioxide, H2O3 can be produced by this method.


Hydrogen peroxide is thermodynamically unstable and decomposes to form water and oxygen with a ΔH of -98.2 kJmol-1 and a ΔS of 70.5 Jmol-1K-1.

2 H2O2 → 2 H2O + O2

The rate of decomposition increases with rising temperature, concentration and pH, with cool, dilute, acidic solutions showing the best stability. Decomposition is catalysed by various compounds, including most transition metals and their compounds (e.g. manganese dioxide, silver, and platinum). Certain metal ions, such as Fe2+ or Ti3+, can cause the decomposition to take a different path, with free radicals such as (HO.) and (HOO.) being formed.

The decomposition of hydrogen peroxide liberates oxygen and heat; this can be dangerous as spilling high concentrations of hydrogen peroxide on an inflammable substance can cause an immediate fire.

Metal oxides, peroxides and superoxides

When the group 1 metals are heated in an excess of air or in O2, the principal products obtained depend on the metal: lithium oxide, Li2O, sodium peroxide, Na2O2, and the superoxides KO2, RbO2 and CsO2.

4Li + O2 → 2Li2O oxide formation
2Na + O2 → Na2O2 peroxide formation
K + O2 → KO2 superoxide formation
The superoxides and peroxides contain the paramagnetic [O2]- (μ ~ 1.73 BM) and diamagnetic [O2]2- ions respectively. See the MO diagram for O2 above for comparison.

All of the monoxides are known, M2O, and the heavier peroxides and superoxides will decompose to form the oxides. They are all ionic and strong bases with the basicity increasing down the group.
O2- + H2O(aq) → 2 OH- (aq)

Sodium peroxoborate is a solid peroxygen compound with exceptional storage stability and no shock sensitivity. It is cheap and readily available, being produced mainly as a solid ingredient of domestic washing formulations, in which it acts as sources of H2O2 in solution for stain bleaching. World annual production in 1995 was approximately 750,000 tonnes, and its use in washing powders dates back to 1907 with Henkel's original "Persil" product in Germany.

Sodium peroxoborate has the empirical formula "NaBO3.xH2O". Two forms that are commercially available correspond stoichiometrically to x = 1 or 4, and are known as the "monohydrate" and "tetrahydrate', respectively. Structurally, however, sodium peroxoborate was shown in 1961 to be the disodium salt of a 1,4.-diboratetroxane dianion. Hence, the "monohydrate" really corresponds to the anhydrous salt, and the "tetrahydrate" to a hexahydrated form of it. The monohydrate form dissolves better than the tetrahydrate and has higher heat stability; it is prepared by heating the tetrahydrate. The compound exists as a dimer as shown below.

The reagent offers low toxicity and a long shelf life. Sodium peroxoborate is a useful reagent in synthetic chemistry as a substitute for the unstable, highly concentrated hydrogen peroxide solutions that can pose a significant explosion hazard and are not commercially available.

Persil and the structure of Na2B2(O2)2(OH)4.6H2O
Persil Advert

Return to the course outline or move on to Lecture 5: Structure of the elements (Groups 1 and 2 metals, Boron, Carbon and Phosphorus, Sulfur).


Much of the information in these course notes has been sourced from Wikipedia under the Creative Commons License.
'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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