Oxidation states of oxygen and its characteristics

Oxy­gen is an el­e­ment of the 6ᵗʰ group (ac­cord­ing to the new clas­si­fi­ca­tion an el­e­ment of the 16th group) of the main sub­group of the pe­ri­od­ic ta­ble. It is a mem­ber of the group of chalco­gens (be­sides oxy­gen, they in­clude sul­fur, se­le­ni­um, tel­luri­um and polo­ni­um). In trans­la­tion, “chalco­gen” means “giv­ing birth to ore”

Phys­i­cal prop­er­ties of oxy­gen, its pres­ence in na­ture

Oxy­gen is a col­or­less di­atom­ic gas in nor­mal con­di­tions. It has no taste or smell, but its pres­ence can be iden­ti­fied with a smol­der­ing stick: when it is placed in an at­mos­phere of oxy­gen it flares up, as oxy­gen sup­ports com­bus­tion well. In an at­mos­phere of oxy­gen, breath­ing and de­com­po­si­tion are also pos­si­ble.

In na­ture it is in en­coun­tered in the form of three iso­topes (el­e­ments with an iden­ti­cal num­ber on the pe­ri­od­ic ta­ble, but with dif­fer­ent atom­ic num­bers). Oxy­gen can be en­coun­tered most fre­quent­ly with the atom­ic num­bers 16, 17 or 18.

In na­ture, the max­i­mum quan­ti­ty of oxy­gen is con­tained in the air (around 21% of its vol­ume), in wa­ter and in the earth’s crust (up to 47% oxy­gen by mass).

Al­lotrop­ic mod­i­fi­ca­tions

Molec­u­lar oxy­gen has the al­lotrop­ic mod­i­fi­ca­tion O₃ (ozone). Ozone is a poi­sonous gas of a bluish col­or with a spe­cif­ic smell. If it is com­pressed, it ac­quires a rich blue col­or. In a sol­id state, the sub­stance forms dark blue, al­most black crys­tals.

It forms from the break­down of molec­u­lar oxy­gen into atom­ic oxy­gen (this takes place with an elec­tri­cal charge, harsh ul­tra­vi­o­let ir­ra­di­a­tion of the air or the break­down of per­ox­ides). The O₃ mol­e­cule is rather un­sta­ble, so for a rel­a­tive­ly short time (sev­er­al tens of min­utes), it breaks down into molec­u­lar oxy­gen. To pre­vent this from hap­pen­ing, it must be cooled to -78 ᵒC or -108.4 ᵒF in a her­met­ic glass or met­al con­tain­er, or have a lit­tle ni­tric acid added to it. Ozone is a very re­ac­tive sub­stance, and sur­pass­es molec­u­lar oxy­gen by its re­ac­tiv­i­ty and ox­i­da­tion prop­er­ties.

Click here for ex­per­i­ments with pure oxy­den.

Ob­tain­ing oxy­gen

There are nu­mer­ous meth­ods for ob­tain­ing oxy­gen:

• dis­til­la­tion of liq­uid air (its rec­ti­fi­ca­tion);

• elec­trol­y­sis of wa­ter ac­cord­ing to the fol­low­ing equa­tion:

2H₂O = O₂ + 2H₂ (pres­ence of an al­ka­li is nec­es­sary);

• elec­trol­y­sis of salts of acids con­tain­ing oxy­gen:

2Cu­SO₄ + 2H₂O₂ = 2Cu + 2O₂ + 2H₂­SO₄;

• break­down of ox­i­diz­ers:

• potas­si­um chlo­rate:

2K­ClO₃ = 3O₂ + 2KCl (with heat­ing);

• potas­si­um per­man­ganate:

2KM­nO₄ = O₂ + MnO₂ + K₂M­nO₄ (with heat­ing);

• potas­si­um ni­trate:

2KNO₃ = O₂ + 2KNO₂ (with heat­ing);

• hy­dro­gen per­ox­ide:

2H₂O₂ = O₂ + 2H₂O (in the pres­ence of man­ganese (IV) ox­ide MnO₂ or with heat­ing);

• re­ac­tion of potas­si­um per­ox­ide with car­bon diox­ide: 2Na₂O₂ + 2CO₂ = O₂ + 2Na₂­CO₃.

In na­ture, oxy­gen is ob­tained by pho­to­syn­the­sis from car­bon diox­ide and wa­ter in light:

mCO₂ + nH₂O = mO₂ + Cₘ(H₂O)ₙ (be­sides oxy­gen, car­bo­hy­drates also form when il­lu­mi­nat­ed with light).

Ox­i­da­tion states of oxy­gen

Be­ing a strong ox­i­diz­er, oxy­gen most fre­quent­ly dis­plays an ox­i­da­tion state of -₂ in com­pounds (oth­er chalco­gens have a low­er ox­i­da­tion state). Typ­i­cal ex­am­ples of com­pounds where oxy­gen has an ox­i­da­tion state of -2 are H₂­SO₄, H₂O, KNO₃ and CaO (in the hy­drox­o­ni­um ion H₃O⁺ the ox­i­da­tion state of oxy­gen is also -2). In com­pounds of this type (apart from the hy­drox­o­ni­um ion), the va­lence of oxy­gen (its abil­i­ty to form a cer­tain num­ber of chem­i­cal bonds) is equal to two. For H₃O⁺, the va­lence of oxy­gen is equal to three, as oxy­gen forms three bonds – two co­va­lent and one donor-ac­cep­tor.

In per­ox­ides, the ox­i­da­tion state of oxy­gen is -1. The for­ma­tion of per­ox­ides is char­ac­ter­is­tic for hy­dro­gen (hy­dro­gen per­ox­ide H₂O₂) and some met­als (for ex­am­ple, sodi­um per­ox­ide Na₂O₂, bar­i­um per­ox­ide BaO₂, cal­ci­um per­ox­ide CaO₂ etc.).

In a free state, oxy­gen has the ox­i­da­tion state of 0, like oth­er sim­ple sub­stances.

Only in one case can oxy­gen have a pos­i­tive ox­i­da­tion state of +2 – in a com­pound with flu­o­rine, with the for­mu­la OF₂. As flu­o­rine is a more elec­tri­cal­ly neg­a­tive el­e­ment than oxy­gen, it ac­quires the neg­a­tive ox­i­da­tion state (-1) in the com­pound. The com­pound is called “oxy­gen flu­o­ride II”. It is ob­tained in the com­bus­tion of wa­ter in an at­mos­phere of flu­o­rine, and also by the re­ac­tion:

2F₂ + 2NaOH = OF₂ + 2NaF + H₂O (ozone and hy­dro­gen per­ox­ide can also form in the re­ac­tion).

There is a com­pound with a pos­i­tive ox­i­da­tion state of oxy­gen (+1), O₂F₂ (oxy­gen monoflu­o­ride, ob­tained in the re­ac­tion of molec­u­lar oxy­gen and flu­o­rine). It is an un­sta­ble com­pound.

Prop­er­ties of molec­u­lar oxy­gen

In the ma­jor­i­ty of cas­es, oxy­gen dis­plays ox­i­diz­ing prop­er­ties in re­ac­tions, both with sim­ple and com­plex sub­stances. If the ox­i­da­tion of oxy­gen takes place vi­o­lent­ly and rapid­ly, this re­ac­tion is called a com­bus­tion re­ac­tion, dur­ing which oxy­gen ox­i­dizes any sub­stance with the vi­o­lent re­lease of en­er­gy (usu­al­ly in the form of heat and light). If the re­ac­tion with oxy­gen takes place gen­tly and slow­ly (of­ten with a cat­a­lyst), then this process is called ox­i­da­tion. Which process takes place in a spe­cif­ic case of­ten de­pends on the con­di­tions in which the re­ac­tion is car­ried out and the prop­er­ties of the ox­i­diz­ing sub­stance.

With non-met­als, oxy­gen re­acts with the for­ma­tion of acidic ox­ides (with heat­ing):

• S + O₂ = SO₂;

• 4P + 5O₂ = 2P₂O₅ (phos­pho­rus pen­tox­ide ex­ists in the form of a dimer in the com­po­si­tion of P₄O₁₀); https://com­mons.wiki­me­dia.org/wiki/Cat­e­go­ry:Phos­pho­rus_pen­tox­ide#/me­dia/File:Phos­pho­rus-pen­tox­ide-3D-balls.png

• C+O₂ = CO₂.

With met­als, oxy­gen re­acts with the for­ma­tion of ba­sic ox­ides or per­ox­ides (the re­ac­tion takes place with heat­ing):

• 4Li + O₂ = 2Li₂O;

• 2Zn + O₂ = 2ZnO;

• 2Na + O₂ = Na₂O₂.

With many com­plex sub­stances, oxy­gen also re­acts (in all cas­es in the sec­ond reagent there is an el­e­ment which is not in the high­est ox­i­da­tion state – it is ox­i­dized by oxy­gen to a high­er ox­i­da­tion state, or to the high­est pos­si­ble state for it). Oxy­gen also ox­i­dizes or­gan­ic and non-or­gan­ic com­pounds (the most sig­nif­i­cant re­ac­tions are giv­en, which re­flect the ox­i­diz­ing prop­er­ties of oxy­gen):

• СН₄ + О₂ = НСОН + Н₂О (in­com­plete ox­i­da­tion of meth­ane is pos­si­ble with the for­ma­tion of formalde­hyde with the use of cat­a­lysts);

• СН₄ + 2О₂ = СО₂ + 2Н₂О (com­plete ox­i­da­tion of meth­ane – com­bus­tion);

• C₂H₂OH + 3O₂ = 2CO₂ + 3H₂O (com­plete ox­i­da­tion of al­co­hol);

• C₂H₂OH + O₂ = CH₃­COOH + H₂O (gen­tle ox­i­da­tion of al­co­hol by oxy­gen to acetic acid);

• 2BaO + O₂ = 2BaO₂ (bar­i­um ox­ide ab­sorbs oxy­gen, form­ing per­ox­ide);

• Na₂O₂ + O₂ = 2NaO₂ (when re­act­ing with per­ox­ides, oxy­gen ox­i­dizes them to su­per­ox­ides with an ox­i­da­tion state of oxy­gen of -1/2);

• 4NH₃ + 3O₂ = 2N₂ + 6H₂O (com­bus­tion);

• 4NH₃ + 5O₂ = 4NO + 6H₂O (cat­alyt­ic ox­i­da­tion with heat­ing in the pres­ence of plat­inum);

• 2SO₂ + O₂ = 2SO₃ (re­ac­tion takes place with heat­ing and adding a cat­a­lyst);

• H₂S + 3O₂ = 2SO₂ + 2H₂O (sul­fur (IV) ox­ide forms in an abun­dance of oxy­gen);

• 2H₂S + O₂ = 2S + 2H₂O (molec­u­lar sul­fur forms in the re­ac­tion with a lack of oxy­gen);

Burn­ing sul­fur

• 4Fe(OH)₂ + O₂ + 2H₂O = 4Fe(OH)₃;

• 4FeS₂ + 11O₂ = 8SO₂ + 2Fe₂O₃ (re­ac­tion takes place with heat­ing with the for­ma­tion of two ox­ides);

• 4CH₃N­H₂ + 9O₂ = ₄CO₂ + 2N₂ + 10H₂O.

Qual­i­ta­tive re­ac­tion to ozone:

O₃ + 2KI + H₂O = I₂ + 2KOH + O₂ (molec­u­lar oxy­gen is re­leased and a dark vi­o­let sed­i­ment of molec­u­lar io­dine forms).

Oxy­gen has found wide ap­pli­ca­tion in in­dus­try (for ex­am­ple in smelt­ing cast-iron and steel). It is also used for cut­ting met­als (with acety­lene) and in medicine. Ozone is used for bleach­ing fab­rics and de­con­tam­i­nat­ing wa­ter.