Oxygen preparation

Preparation in laboratory

Preparation in industry

Preparation in laboratory

Decomposition of rich in oxygen unstable compounds:

KСlО₃ = 2KСl + 3О₂ (catalysts МnO₂, Fе₂О₃, Сr₂О₃)

(Berthollet's Salt)

2KМnО₄ = K₂МnО₄ + МnО₂ + О₂

2KNО₃ = 2KNО₂ + О₂

2HgО = 2Hg + О₂

2H₂O₂ = 2H₂O + O₂ ( catalyst МnO₂)

2. Interaction of МnО₂ and concentrated H₂SO₄:

2МnО₂ + 2H₂SO₄ = 2МnSО₄ + O₂ +2H₂O

3. Interaction of peroxides and superoxides with water (carbon dioxide):

4KO₂ + 2H₂O = 4KOH+ 3O₂

2Na₂O₂ + 2CO₂ = 2Na₂CO₃ + O₂

4. The former industrial (till 1911) A. Breen method consists of two steps:

a) heating (300°С) of ВаО together with O₂

2ВаО + О₂ = 2ВаО₂

b) ВаО₂ further thermal decomposition (at 800 °С)

2ВаО₂ = 2ВаО + О₂

5. Electrolysis of water

PRODUCTION IN INDUSTRY

At present oxygen is produced in industry from liquid air. Rectification-fractional distillation of liquid air allows to separate oxygen, nitrogen, and noble gases in large amounts. To liquefy gaseous oxygen it is cooled below critical temperature (-118°С) and then strongly compressed.

T o liquefy the air is compressed up to 100 atm. Gas temperature grows simultaneously. Hot compressed air is passed through a refrigerator where it is cooled to the room temperature. Then, the gas pressure is decreased to 10 atm and its temperature is decreased by 30°С due to expansion (this physical phenomenon is referred to as Joule-Thompson effect). Repeted compression-expanding cycles allow to achieve a liquid state of air. Since O₂ has higher t_boiling (-183°С), compared with N₂ (-196°С) therefore N₂ is evaporated at distillation first and pure oxygen remains in a liquid phase.

Electrolysis of water (especially pure oxygen is produced)

О₂ is stored and transported at the pressure of 150 atm in blue balloons made of steel, with the inscription “Oxygen” in black.

Notwithstanding its low boiling point (about — 190 °C), liquid air can be stored at atmospheric pressure for quite a long time in the Dewar flasks.

OXYGEN. STANDARD ELECTRODE POTENTIALS, E^O, V

0		-1		-2
Acid medium
O2				H2O		O3			O2
		H2O2
Alkaline medium
O2				OH-		O3		O2

Chemical properties of oxygen

Oxygen as a ligand

Oxygen has a high chemical activity. It enters reactions with most elements, most frequently forming oxides:

4Аl + 3О₂ = 2Аl₂О₃ DG_f° = -1582.3 kJ/mol

4P + 5О₂ = 2P₂О₅ DG_f° = -1371.7kJ/mol

Significantly negative G_f values are an evidence of a considerable affinity of elements to oxygen.

Dioxygen does not react only with noble gases, some metals (Au, Pt) and elemental halogens. Gibbs energy change of the mentioned elements reaction with dioxygen is positive (G_f>0), for example, DG_f° (Аu₂О₃) = 79 kJ/mol.

Nevertheless the oxides of these elements, except for He, Ne and Ar, can be obtained indirectly:

НСlО₄ + Р₂О₅ = Сl₂О₇ + 2НРО₃

2Аu(ОН)₃ = Аu₂О₃ + ЗН₂О

Dioxygen also interacts with a lot of compounds. Oxides, sometimes oxides with elemental substances, are formed:

СН₄ + 2О₂ = СО₂ + 2Н₂О;

SiН₄ + 2О₂ = SiО₂ + 2Н₂О;

4NН₃ + 3О₂ = 6Н₂О + 2N₂.

In most cases oxidation reactions proceed at a sufficient rate only at

high temperatures. Initial activation of О₂ molecules is usually required.

If oxidation of compounds proceeds vigorously with the formation of large quantity of heat and light it is named burning. It takes place at the excess of air or pure О₂ and is completed with the formation of the final products of oxidation (СО₂, Н₂О, N₂ and others). For example, the flame temperature can reach 3200C at burning of acetylene С₂Н₂ in oxygen medium.

A smoldering splinter lights up at oxygen content in a mixture with N₂ not less than 28%. If a mixture contains less than 16%, a burning splint extincts.

Content of oxygen	Conditions of burning combustible compounds and life of man
17% О2	An oily miner's flashlight goes out; breathing remains normal
14%	A candle extincts
12%	A flame of acetylene extincts; breathing is deeper than normal one
10%	A man begins to feel strangled
6%	A Hydrogen flame extincts
5%	A man’s noticeable convulsions and sometimes fatal
2%	Sudden fainting fit less than in a 1 minute. During 10 minutes life can be rescued by means of artificial respiration.

Relatively slow reactions of oxidation depending on the character of substance appear as corrosion (oxidation of metals), and decay (oxidation of organic residues), more generally they are simply called oxidation.

In the reactions of oxidation there can be present the products of incomplete oxidation. So, burning Н₂S in air (in excess of О₂) gives the product of complete oxidation:

2Н₂S + 3О₂ = 2Н₂O + 2SO₂

and at gradual oxidation H₂S in aqueous solution with the dissolved O₂ (the lack of dioxygen) elementary S forms (appearance of H₂S solutions opalescence at their storage):

2Н₂S + О₂ = 2Н₂O + 2S

The high standard E value of the half-reaction

О₂ + 4Н⁺ + 4е = 2Н₂О E = 1,23 V (at pH=0)

is the evidence of considerable oxidising activity of oxygen.

Oxides. Formation of monoatomic multicharge anions of Е^n- type is energetically unprofitable. For these reasons there are no compounds, which have a free ion О^2- in the composition. Even in the crystalline oxides of the most active metals (Na₂О, СаО) an effective negative charge of oxygen is substantially less than -2.

In the oxides of non-metals differences in the electronegativities of oxygen and non-metal atoms are small, therefore the chemical bond is polar covalent. Most of such compounds at STP are gases, volatile liquids or fusible compounds: m.p. (SO₂) = -75°С, (СlО₂) = -121C, sublimation temperature (Р₄О₁₀) = 35C. It can be explained by the fact that in the solid state they form molecular crystalline structures with weak intermolecular forces. On the contrary, oxides with polymeric structures are considerably stronger and more refractory. So, SiO₂ is polymer, unlike monomeric SiO_2, at STP is chemically inert and has high m.p. (1700С).

The oxides of metals can be basic, amphoteric and acidic. Generally the oxides of non-metals according to the chemical nature are acidic, the most of them dissolve in water forming acids, and therefore they are named the anhydrides of acids:

СO₂ + Н₂О Û Н₂СO₃

SO₂ + Н₂О Û Н₂SO₄

Acidic oxides react with bases, basic and amphoteric oxides.

The oxides of the most active metals have mostly ionic bonds and according to their chemical nature are basic (oxides of alkaline and alkali earth metals). Their hydroxides are strong bases (dissociate at bond M—OH, which is ionic):

ВаО + Н₂О = Ва(ОН)₂ G_f = -97,9 kJ/mol

Basic oxides react with acids, acidic and amphoteric oxides.

MgО + SiO₂ = MgSiO₃ H₂₈₉ = -59 kJ/mol

Basic oxides are formed by active metals with low oxidation states (+1 or +2). They have an ionic bond, their crystalline structures have coordination number with oxygen atoms 6 or 8.

With the growth of metal oxidation state, the covalent part in its chemical bond with oxygen is increased. Such oxides with ionic-covalent bond are amphoteric. Similarly to amphoteric oxides amphoteric hydroxides can dissociate along the bond M—OH (like weak bases) or along the bond MO—H like weak acids. Such oxides include mainly the oxides of those metals that have a comparatively small EN, for example, Al₂O₃, Cr₂O₃, PbО, ZnO, Fe₂O₃. Depending on conditions they reveal properties of basic or acid oxides. The amphoteric oxides do not interact with water (are not dissolved), but can react both with acids, and with alkalis:

ZnO + 2HCl = ZnCl₂ + H₂O

ZnO + 2NaOH + H₂O = Na₂[Zn(OH)₄],

and at heating - with basic and acid oxides:

PbО + SO₃ = PbSО₄ G_f = -255 kJ/mol

PbО + Na₂O = Na₂PbО₂ G_f = -84 kJ/mol

PEROXIDE TYPE COMPOUNDS

A dioxygen molecule can add or lose electrons during chemical transformations, forming molecular ions O₂^-, O₂^2-, O₂⁺. Their stability is predicted by MO method. They are called peroxide type compounds since these species have two-bonded atoms of oxygen (O—O).

Table 1. Molecular orbitals of dioxygen, O₂^-, O₂^2-, O₂⁺ (MO method)

Molecular orbitals	O2+	O2	O2-	O22-
sр*	—	—	—	—
pр* pр*	­	­ ­	­¯ ­	­¯ ­¯
pр pр	­¯ ­¯	­¯ ­¯	­¯ ­¯	­¯ ­¯
ss	­¯	­¯	­¯	­¯
ss*	­¯	­¯	­¯	­¯
ss	­¯	­¯	­¯	­¯
Bond order	2.5	2	1.5	1
O oxidation state	+ ½	0	- ½	-1
Distance O—O, nm	0.112	0.1207	0.132	0.149
Dissociation energy, kJ/mol	642	494	394	210

Adding one electron onto *_2p molecular orbital, the molecule of oxygen forms a molecular ion O₂^-, which is named superoxide-ion, and its compounds are superoxides. The process O₂ + e^- = O₂^- is accompanied by the liberation of heat (H₂₈₉ = -48.1 kJ/mol) and decreases the bond order to 1.5 (the third electron is added to unpaired electrons of antibonding *_2p orbitals of the О₂ molecule).

Superoxides can be obtained directly only for the most active reductants (alkali metals below potassium in IA group):

K + O₂ = KO₂

Dioxygen forms an ion О₂⁺ (bond order 2.5) at elimination of one electron from the *_2p molecular orbital. This ion is named a dioxygenyl-ion. Dioxygen is only oxidized by very strong oxidants, for example by Pt(VI) fluoride:

О₂ + PtF₆ = О₂[Pt⁵⁺F₆]

Dioxygenil hexafluoroplatinate (V)

These salts are also synthesized by heating a mixture of О₂, F₂ and a powdered element to 150-500C:

О₂ + 3F₂ + M = О₂⁺[MF₆],

where M = As, Sb, Bi, Nb, Pt, Au, Ru, Rh.

A molecule О₂ can accept two electrons forming the peroxide-ion О₂^2-, derivatives of which are named peroxides. Additional electrons are added onto *_2p MO. Thus the bond order lowers down to 1. Absence of the unpaired electrons determines diamagnetism of peroxides. Peroxides are formed during combustive oxidation of some active metals:

2Na + O₂ = Na₂О₂

Ва + О₂ = ВаО₂,

Hydrogen Peroxide H₂O₂

It is one of the most important compounds of oxygen.

Production of peroxides in industry:

1. Electrolysis of 50%-Н₂SO₄ solution or some sulfates (in particular, (NН₄)₂SO₄) at high current density and cooling. Anode oxidation of hydrogen sulfate-ions to peroxodisulfuric acid, Н₂S₂O₈, occurs on a platinum anode:

cathode: 2H⁺ + 2е = Н₂

anode: 2НSO₄^- = Н₂S₂O₈ + 2е

The slight heating of H₂S₂O₈ solution leads to its hydrolysis:

Н₂S₂O₈ + Н₂O = Н₂O₂ + Н₂SO₄

In industry Н₂O₂ is produced and sold with concentration up to 60%, as more concentrated solutions are impossible to transport due to their high explosiveness. Concentrated to 30% and stabilized Н₂О₂ solution is named perhydrol. 3% solution of Н₂О₂ is used for medical purposes. World production of Н₂О₂ is a 0.5 million of tons per year.

2. Chemical methods are based on organic compounds oxidation with oxygen. Н₂О₂ together with other useful product - acetone are obtained at catalytic oxidation of isopropanol:

(СН₃)₂СНОН + О₂ ® СН₃СОСН₃ + Н₂О₂

Preparation in a laboratory:

It is easily obtained from BaO₂ at interaction with diluted Н₂SO₄:

ВаО₂ + Н₂SO₄ = Н₂О₂ + ВаSO₄¯

Structure. A molecule of Н₂О₂ is nonlinear, two bonds O—H are located in two planes:

Owing to the asymmetrical location of bonds the O—H molecule Н₂О₂ is strongly polar (electric dipole moment  = 0.7•10^-29 C•m), it exceeds water polarity  = 0.61•10^-29 C•m. The presence of unshared electron pairs of oxygen atoms in hydrogen peroxide creates a possibility for donor-acceptor interaction.

Properties. Instability of solutions. The Н₂О₂ of high purity and its solutions at STP are quite stable and can be stored for a long time, but at elevated temperatures, UV-irradiation, and also in the presence of ions of transition metals Н₂О₂ intensively decomposes:

2Н₂О₂ = 2Н₂О + О₂

Decomposition of pure Н₂О₂ and its concentrated solutions proceeds violently, explosion-like.

Additives (stabilizers, usually Na₃PO₄) are used to prevent Н₂О₂ decomposition in aqueous solutions. Stabilizers bind the ions of metals and as a result break undesirable catalytic process of decomposition. Solutions Н₂О₂ are stored in a cool place in dark vessels. It is worth knowing that even the minor content of alkalis leached from ordinary glass bottles substantially speed up the decomposition of Н₂О₂.

In the blood of a man and animals and juices of plants there is a specific enzyme of catalase, which decomposes Н₂О₂. The active component of catalase is ion of iron bound to complex organic molecules. One molecule of catalase decomposes at STP approximately 100 thousand molecules of Н₂О₂ per 1 sec.

Physical properties. Properties of liquid and solid Н₂О₂, as well as its solutions are determined by strong hydrogen bonds, which cause association of molecules. Therefore pure Н₂О₂ at STP is a syrup-like viscous, pale-blue, odourless liquid, which is almost 1.5 times heavier than water ( = 1.44 g/cm³ at 25С), t_boil = 150.2C (to determine t_boil Н₂О₂ at atmospheric pressure is impossible, as already at 90C it decomposes).

Н₂О₂ is miscible in water at any proportion due to H-bonds formation.

Hydrogen peroxide is a very weak acid in aqueous solutions:

Н₂О₂ Н⁺ + НО₂^- К = 2.24•10^-12

but it is stronger than water. Dissociation at the second stage,

НО₂^- Н⁺ + О₂^2-

virtually does not occur. It is suppressed by the presence of water — a substance that dissociates with the formation of hydrogen ions to a greater extent than hydrogen peroxide.

Peroxides of metals which belong to the class of salts (not oxides) can be obtained by the interaction of a hydrogen peroxide solution with alkalis:

Н₂О₂ + Ва(ОН)₂ = ВаО₂ + 2Н₂О

Unlike normal oxides which form salt and water with acids, peroxides form salt and hydrogen peroxide in similar reactions:

ВаО₂ + Н₂SO₄ = BaSО₄ + Н₂О₂

ВаО + Н₂SO₄ = BaSО₄ + Н₂О

SnО₂ + 2Н₂SO₄ = Sn(SО₄)₂ + 2Н₂О

In aqueous solution the peroxides of metals, being the salts of a weak acid, are unstable. It is obvious at its strong hydrolysis and decomposition of Н₂О₂in alkaline medium:

Na₂О₂ + 2Н₂О Û 2NaОН + Н₂О₂

2Н₂О₂ ® 2Н₂О + О₂

The hydrogen atoms in hydrogen peroxide can be substituted not only by a metal but, for example, by acid residues with the formation of peroxide compounds of various types:

Peroxodisulfuric peroxosulfuric acids

peroxonitric acid ( aquafortis)

НО—ОН

hydrogen peroxide

 

Bond O—O is weak due to the mutual electrostatic repulsion of two unshared pairs of electrons of each bound atom of oxygen. It is almost three times weaker than the bond O—H. Reactions of O—O bond destruction with the formation of compounds of oxygen (-2) or 0 are typical. Therefore the most characteristic Н₂О₂ transformations are namely redox reactions, where Н₂О₂ can be an oxidant or reductant with the following E’s:

Н₂О₂ + 2Н⁺ + 2е^- = 2Н₂О E = 1.78 V

О₂ + 2Н⁺ + 2е^- = Н₂О E = 0.68 V

Oxidizing activity (E = 1.78 V) of H₂O₂ is considerably stronger, than its reducing agent strength (E= 0.68 V).

Indeed, peroxides are strong oxidants, water is the product of reduction:

2KI + Н₂О₂ + Н₂SO₄ = I₂ + 2Н₂О + K₂SO₄

KNO₂ + H₂O₂ = KNO₃ + H₂O

2K₃[Cr(OH)₆] + 3Н₂О₂ = 2K₂CrO₄ + 8Н₂О + 2KOH

Hydrogen peroxide can demonstrate also reducing properties (only with strong oxidants), oxygen is the product of oxidation of Н₂О₂, for example:

Н₂О₂ + Cl₂ = О₂ + 2НCl

Н₂О₂ + О₃ = 2О₂ + Н₂О

3Н₂О₂ + 2KMnО₄ = 3О₂ + 2MnО₂ + 2KOH + 2Н₂О

5Н₂О₂ + 2KMnО₄ + 3H₂SO₄ = 5О₂ + 2MnSO₄ + K₂SO₄ + 3Н₂О

The last reaction is used in chemical analysis for quantitative determination of Н₂О₂.

Н₂О₂ decomposition belongs to the disproportionation reactions type (autooxidation-autoreduction):

2 Н₂О₂ = 2 Н₂О + О₂.

in the atmosphere into ozone and thereby provides a protective shield against short-wavelength UV light:

Practical activity of a man often predetermines destruction of ozone layer. Chlorine presence in atmosphere stands here in the first place. One molecule of chlorine is able to annihilate up to 100 thousands molecules of О₃ (comparison: 1 molecule NO destroys 10 molecules О₃ only). One launch of spaceship «Shuttle» destroys 0.3% О₃ from its common quantity in an earthly atmosphere.

Properties. О₃ is a toxic gas of a dark blue color with a strong characteristic smell, b.p. = -110C, m.p. = -192.7C. Liquid ozone has a navy blue color, solid ozone is black.

The molecule of О₃ diamagnetic, it has bent structure, it is polar ( = 0.17•10^-29 Cm). Length of bond between the atoms of oxygen (0.128 nm) has intermediate value compared to single (0.149 nm) and double (0.121 nm) bonds. Therefore the bond order in a molecule О₃ equals 1.5. The valence angle value (116.5) is the evidence of sp²-hybridization of valence orbitals of the central oxygen atom. It is also the substantiation of delocalized three-center -bond formation in О₃ molecule. Thus the central sp²-hybridized atom of oxygen forms 2 -bonds with neighboring atoms. Its nonhybridized 2р_z-orbital is located perpendicularly to the plane of three oxygen atoms like 2р_z-orbitals of neighboring oxygen atoms forming together delocalized three-centered -bond.

The asymmetrical structure of molecule О₃ predetermines its polarity. Therefore the energy of О₃ intermolecular interaction in the condensed state exceeds the latter of О₂ molecules. Consequently, t_boil and t_melt of ozone is higher than in О₂. These factors explain also ozone’s better solubility in polar solvents like water (45 volumes of О₃ is soluble in 100 volumes of water at STP).

О₃ easily decomposes and forms dioxygen and atomic oxygen, therefore it has a high oxidizing activity:

О₃ = О₂ + О

That is why О₃ easily oxidizes various substances, transforming itself into О². О₂ molecule remains as a stable fragment of unstable ozone molecule:

2Ag + О₃(OО₂) = Ag₂O + О₂

¯1e ­2e

PbS + 4О₃ = PbSO₄ + 4О₂

2KI + О₃ + H₂O = I₂ + О₃ + 2KOH

The latter reaction is used as the method of determination of ozone in quantitative chemical analysis.

In some case ozonides are the products of reactions with participation of ozone:

К + О₃ = КО₃

4КОН + 4О₃ = 4КО₃ + О₂ + 2Н₂О

Such compounds contain a molecular ion О₃^-. Unlike О₃ an О₃^-ion has an unpaired electron that predetermines its paramagnetizm. The ozonides of alkali metals have a red color and are very strong oxidants.