Does the oxidation number of oxygen change?

The oxidation number of oxygen can vary.

According to oxidation state rules, any element that is not combined with other elements in a compound is zero.

Neutral compounds have net zero charge, so the charges of elements in a compound must equal zero.

Due to its high electronegativity, oxygen usually has a negative two charge. For example in the compound, calcium oxide, CaO, calcium has a oxidation number of +2 and the oxygen has -2 charge.

In peroxides, such as hydrogen peroxide, #H_2O_2#, each hydrogen has +1 charge, to give a combined oxidation number of +2.
That means that oxygen component, #O_2# ,has a combined charge of -2. Consequently, each oxygen must carry a negatives one charge.

Fluorine is the most electronegativity element and when it combines with oxygen to form oxygen difluoride, #OF_2#, the oxidation for each fluorine is negative and the combined oxidation number for two fluorine is negative two. Therefore to balance the charges, oxygen in this case is positive two.

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Oxidation Numbers

It is often useful to follow chemical reactions by looking at changes in the oxidation numbers of the atoms in each compound during the reaction. Oxidation numbers also play an important role in the systematic nomenclature of chemical compounds. By definition, the oxidation number of an atom is the charge that atom would have if the compound was composed of ions.

1. The oxidation number of an atom is zero in a neutral substance that contains atoms of only one element. Thus, the atoms in O2, O3, P4, S8, and aluminum metal all have an oxidation number of 0.

2. The oxidation number of simple ions is equal to the charge on the ion. The oxidation number of sodium in the Na+ ion is +1, for example, and the oxidation number of chlorine in the Cl- ion is -1.

3. The oxidation number of hydrogen is +1 when it is combined with a nonmetal as in CH4, NH3, H2O, and HCl.

4. The oxidation number of hydrogen is -1 when it is combined with a metal as in. LiH, NaH, CaH2, and LiAlH4.

5. The metals in Group IA form compounds (such as Li3N and Na2S) in which the metal atom has an oxidation number of +1.

6. The elements in Group IIA form compounds (such as Mg3N2 and CaCO3) in which the metal atom has a +2 oxidation number.

7. Oxygen usually has an oxidation number of -2. Exceptions include molecules and polyatomic ions that contain O-O bonds, such as O2, O3, H2O2, and the O22- ion.

8. The elements in Group VIIA often form compounds (such as AlF3, HCl, and ZnBr2) in which the nonmetal has a -1 oxidation number.

9. The sum of the oxidation numbers in a neutral compound is zero.

H2O: 2(+1) + (-2) = 0

10. The sum of the oxidation numbers in a polyatomic ion is equal to the charge on the ion. The oxidation number of the sulfur atom in the SO42- ion must be +6, for example, because the sum of the oxidation numbers of the atoms in this ion must equal -2.

SO42-: (+6) + 4(-2) = -2

11. Elements toward the bottom left corner of the periodic table are more likely to have positive oxidation numbers than those toward the upper right corner of the table. Sulfur has a positive oxidation number in SO2, for example, because it is below oxygen in the periodic table.

SO2: (+4) + 2(-2) = 0

Oxy­gen is an el­e­ment of the 6ᵗʰ group (un­der the new clas­si­fi­ca­tion the 16ᵗʰ group) of the main sub­group of the pe­ri­od­ic ta­ble. It is a rep­re­sen­ta­tive of the chalco­gens group (they also in­clude sul­fur, se­le­ni­um, tel­luri­um and polo­ni­um). Oxy­gen is a di­atom­ic col­or­less gas with­out smell or taste. It sup­ports breath­ing, com­bus­tion and de­com­po­si­tion. It is en­coun­tered in the form of 3 iso­topes – in na­ture, oxy­gen with the atom­ic num­bers of 16, 17 and 18 is en­coun­tered.

Oxy­gen is a strong ox­i­diz­er (only flu­o­rine dis­plays stronger ox­i­da­tion prop­er­ties be­cause of its greater elec­tri­cal neg­a­tive­ly and its more pro­nounced non-metal­lic prop­er­ties (by its po­si­tion in the pe­ri­od­ic ta­ble)). Oxy­gen is ca­pa­ble of dis­play­ing sev­er­al ox­i­da­tion states in chem­i­cal re­ac­tions: -2, -1, 0, +2.

Oxy­gen in the ox­i­da­tion state of -2

The low­est ox­i­da­tion state of oxy­gen is -2. As this non-met­al is a strong ox­i­diz­er, it fre­quent­ly dis­plays this ox­i­da­tion state in com­pounds. We may pro­vide many ex­am­ples of such com­pounds among salts, acids, ox­ides and bases: KClO₄, H₂­SO₄, N₂O₃, NaOH etc. In wa­ter and in a hy­dro­ni­um ion, the ox­i­da­tion state of oxy­gen is also two.

The va­lence of oxy­gen in these two com­pounds is dif­fer­ent, how­ev­er. In wa­ter oxy­gen shows typ­i­cal va­lence of 2, and in the hy­dro­ni­um ion, from the for­ma­tion of the third, donor-ac­cep­tor bond, va­lence (abil­i­ty to form a cer­tain num­ber of bonds) grows to three. The donor-ac­cep­tor bond forms be­cause the un­shared pair of elec­trons in the oxy­gen atom are lo­cat­ed on the free or­bital of the hy­dro­gen cation Н⁺.

Many re­ac­tions take place with­out a change in ox­i­da­tion states:

H₂­SO₄ + 2NaOH = Na₂­SO₄ + 2H₂O;

CaO + H₂O = Ca(OH)₂;

But usu­al­ly, even in ox­i­da­tion-re­duc­tion re­ac­tions, oxy­gen does not ox­i­dize to high­er ox­i­da­tion states, and pre­serves the val­ue of -2:

10KI + 2KM­nO₄ + 8H₂­SO₄ = 5I₂ + 2Mn­SO₄ + 6K₂­SO₄ + 8H₂O.

Ox­i­da­tion of oxy­gen takes place in the break­down of sub­stances (for ex­am­ple, wa­ter or ox­i­diz­ers), or when wa­ter re­acts with flu­o­rine:

• 2H₂O = O₂ + 2H₂ (car­ried out in the pres­ence of an al­ka­li);

• 2KClO₃ = 3O₂ + 2KCl (with heat­ing);

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

• 2KNO₃ = O₂ + 2KNO₂ (with heat­ing);

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

• 2F₂ + H₂O = 4HF + O₂.

Oxy­gen in the ox­i­da­tion state of -1

In per­ox­ides, the ox­i­da­tion state of oxy­gen is -1. The for­ma­tion of per­ox­ides is char­ac­ter­is­tic for hy­dro­gen (H₂O₂) and cer­tain met­als (Na₂O₂, BaO₂, CaO₂ etc.).

In the case with per­ox­ides and su­per­ox­ides (such as KO₂, where the ox­i­da­tion state of oxy­gen is -0.5), both an in­crease and a de­crease of the ox­i­da­tion state of oxy­gen in re­ac­tions are pos­si­ble:

• 2H₂O₂ = O₂ + 2H₂O (in the pres­ence of man­ganese (IV) ox­ide MnO₂ or when heat­ed);

• 2Na₂O₂ + 2CO₂ = O₂ + 2Na₂­CO₃;

• Ag₂O + Н₂О₂ = 2Ag + H₂O + O₂.

• 4KO₂ + 2H₂­SO₄ = 2H₂O + 3O₂ + 2K₂­SO₄;

• 4KO₂ + 2H₂O = 4KOH + 3O₂;

• 2КM­nO₄ + 5Н₂О₂ + 3H₂­SO₄ = K₂­SO₄ + ₂Mn­SO₄ + 5O₂ + 8H₂O;

Here there are oth­er ex­per­i­ments with potas­si­um per­man­ganate.

• Na₂O₂ + CO = Na₂­CO₃;

• Na₂O₂ + SO₂ = Na₂­SO₄ (in the pres­ence of sul­fu­ric acid or hy­dro­gen per­ox­ide the re­ac­tion takes place more quick­ly);

• 4Na₂O₂ + NH₃ = NaNO₃ + 3NaOH + 2Na₂O.

The ox­i­da­tion state of oxy­gen does not change in the im­pact on per­ox­ides of di­lut­ed acids:

Na₂O₂ + H₂­SO₄ = H₂O₂ + Na₂­SO₄.

As hy­dro­gen per­ox­ide has weak­ly pro­nounced acidic prop­er­ties, it can re­act with al­ka­lis with­out a change in the ox­i­da­tion state of oxy­gen:

Ва(ОН)₂ + Н₂О₂ = ВаО₂ + 2Н₂О.

Oxy­gen in the ox­i­da­tion state of 0

In a free state, oxy­gen has an ox­i­da­tion state of 0, like oth­er sim­ple sub­stances. As oxy­gen is a strong ox­i­diz­er, it re­acts with many met­als and non-met­als, and also com­pounds, dis­play­ing ox­i­diz­ing prop­er­ties (the ox­i­da­tion state of oxy­gen drops to -2, but if per­ox­ide forms, to -1). De­pend­ing on the con­di­tions, the same sub­stances may re­act with oxy­gen dif­fer­ent­ly:

• S + O₂ = SO₂ (with heat­ing);

• 4Li + O₂ = Li₂O (with heat­ing);

• 2Na + O₂ = Na₂O₂ (prod­uct of com­bus­tion of sodi­um in air – sodi­um per­ox­ide);

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

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

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

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

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

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

• 2H₂S + O₂ = 2S + 2H₂O (with lack of oxy­gen);

• 4Fe(OH)₂ + O₂ + 2H₂O = 4Fe(OH)₃;

• 4FeS₂ + 11O₂ = 8SO₂ + 2Fe₂O₃ (takes place with heat­ing);

• СН₄ + 2О₂ = СО₂ + 2Н₂О (com­plete ox­i­da­tion of met­al - com­bus­tion);

• C₂H₅OH + 3O₂ = 2CO₂ + 3H₂O (com­plete ox­i­da­tion of al­co­hol);

• C₂H₅OH + O₂ = CH₃­COOH + H₂O (mild ox­i­da­tion of al­co­hol by oxy­gen to acetic acid).

Oxy­gen in the ox­i­da­tion state of +2

As a sin­gle atom, oxy­gen has a pos­i­tive ox­i­da­tion state of +2 – in a com­pound with flu­o­rine, OF₂. As flu­o­rine is a more elec­tri­cal­ly neg­a­tive el­e­ment, it and not oxy­gen ac­quires the neg­a­tive ox­i­da­tion state of -1 in the com­pound.

Oxy­gen flu­o­ride forms by the re­ac­tion:

2F₂ + 2NaOH = OF₂ + 2NaF + H₂O (in the re­ac­tion, ozone and hy­dro­gen per­ox­ide H₂O₂ can also form).

The com­pound with the pos­i­tive ox­i­da­tion state of oxy­gen +1 ex­ists, O₂F₂ (oxy­gen monoflu­o­ride). Oxy­gen monoflu­o­ride is an un­sta­ble com­pound, and it can be ob­tained in the re­ac­tion of molec­u­lar gas­es – oxy­gen and flu­o­rine.

Oxy­gen and re­ac­tions with it have found wide ap­pli­ca­tion in lab­o­ra­to­ry prac­tice (for ob­tain­ing ox­ides and oth­er sub­stances) and in in­dus­try (for ex­am­ple in smelt­ing cast-iron and steel). It is also used for cut­ting met­als (with acety­lene) and in medicine.

Do oxidation numbers change?

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