Redox Reactions in Chemistry

Introduction

The redox or redox reactions are reactions in which there is a variation in the oxidation number ( n. O. ) Of ions or atoms.

The chemical species that oxidizes yields electrons and increases the oxidation number.
The chemical species that is reduced acquires those electrons, decreasing the oxidation number.

Obviously the chemical species that oxidizes acts as a reducing agent while, on the contrary, the one that is reduced acts as an oxidant .

In a redox reaction the charge balance must be equal to zero.

Redox reactions can be proposed in two ways: in molecular form and in ionic form .

Redox reactions in molecular form: they describe all the atoms, mostly in the form of indissociated molecules, which participate in the overall reaction, even those that do not enter the redox, as they do not undergo variations in the no .

Redox reactions in ionic form: they only show the ions and the undissociated molecules in which a change of the no takes place; among the reactants there are also + ions , or OH ions  or molecules of 2 O depending on whether the reaction takes place in an acid, basic or neutral environment; also in the reaction products there are + ions or OH  ions , ie molecules of 2 O for the overall balance of the charges. The molecules of 2 Othey can come from a combination of + ions with oxygen released by the oxidizing species or from the 4OH  —> 2H 2 O + O 2 reaction .

Whenever possible it is therefore preferable to make the molecular reactions in ionic form with the following method:

  • The no is attributed to each atom and it is verified in which it undergoes a variation.
  • Molecules in which atoms have undergone no modifications are dissociated into ions. This dissociation occurs mostly for salts, acids and bases while the diatomic molecules of gases, oxides of any type and some binary molecules such as NH 3 , PH 3 do not dissociate .
  • It is observed in which environment the reaction takes place (acid, basic or neutral).
  • The reaction is written in net ionic form, thus including:
  • a) – Ions and molecules in which the no
  • b) – H + ions or OH  ions , or molecules of H 2 O depending on the reaction environment.
  • c) – H + ions or OH  ions , or molecules of H 2 O for balancing the charges.
  1. 2Cr 2 O 7 + KI + HNO 3 —-> KNO 3 + Cr (NO 3 ) 3 + I 2 + H 2 O :

the oxidation numbers are assigned:

2 +1 Cr 2 +6 7 -2 + K +1 -1 + H +1 +5 3 -2 —-> K +1 +5 3 -2 + Cr +3 ( N +5 3 -2 3 + I 2 0 + H 2 +1 -2

The chemical species that modify the no and that, therefore, enter the redox are those indicated.
In the case of 2 Cr 2 O 7 only Cr modifies the no but it is necessary to indicate all the polyatomic ion Cr 2 O +2 rather than the single ion. This procedure must also be followed in the case of other acid residues (eg NO – 3 ,, SO 2- CO 2- , etc.).

The environment is acidic due to the presence of HNO 3 , so the H + ions are reported .

The reaction is thus rendered in a net and unbalanced ionic form:

Cr 2 O -2 + I  + H + —-> Cr 3+ + I 2 + H 2 O .

Redox balancing with the half-reaction method

alancing a reaction means attributing stoichiometric coefficients to each substance present, so that conservation of mass and conservation of electrical charges is possible.

In other words, the number of atoms, for each chemical species, present in the reactants must be equal to the number of atoms of the same chemical species present in the reaction products; the total electric charge of the reactant substances must be equal to the total charge of the products.

The balancing procedures are various; the one based on the half-reaction method or ionic-electronic method can be described as follows, using the reaction in an acid environment already proposed:

1) – The oxidation and reduction half-reactions are written separately:

 —-> I 2 (oxidation)
Cr 2 O -2 —-> Cr 3+ (reduction)

2) – Atoms and ions are balanced; we indicate the moving electrons:

2I  —-> I 2 + 2e

Cr 2 O -2 + 6e —-> 2Cr 3+

3) – If the number of electrons involved in the two half-reactions is not equal, the mcm (least common multiple) of the two values ​​is calculated and divided by the number of electrons in each half-reaction.

The coefficient obtained must be multiplied by the number of electrons, atoms and ions of each half-reaction:

mcm between 6 and 2 = 6; divide this value by the number of electrons in the two half-reactions:

6I  —-> 3I 2 + 6e 6: 2 = 3 (multiplier coefficient)
Cr 2 O -2 + 6e —-> 2Cr 3+ 6: 6 = 1 (no need to multiply)

4) – Since oxygen is present, it is necessary to balance it with molecules of 2 O :

6I  —-> 3I 2 + 6e

Cr 2 O -2 + 6e —-> 2Cr 3+ + 7H 2 O

5) – Hydrogen in water is balanced with + ions (acid environment); the balance of the charges is checked:

6I  —-> 3I 2 + 6e charges: -2 —> -2
Cr 2 O -2 + 6e —-> 2Cr 3+ charges : 0 —> 0

6) – The algebraic sum of the two half reactions is performed, making the necessary simplifications:

6I  —-> 3I 2 + 6e
Cr 2 O -2 + 14H + + 6e —-> 2Cr 3+
————————————————– ———-
6I  + Cr 2 O -2 + 14H + —-> 3I 2 + 2Cr 3+ + 7H 2 O

The ionic reaction is thus balanced.

At this point it is possible to write the balanced reaction also in molecular form:

2 Cr 2 O 7 + + 6KI 14HNO 3 —-> 8KNO 3 + 2Cr (NO 3 ) 3 + 3I 2 + 7H 2 O .

The same procedure must also be used in the case of reactions in a basic environment :

 

Consider the reaction in a basic environment:

Cl 2 + I  + OH  —-> Cl  + IO  + H 2 O

1) – The oxidation and reduction half-reactions are written separately:

 —-> IO 

Cl 2 —-> 2Cl 

2) – Atoms are balanced and electrons in motion are indicated:

 —-> IO  + 6e

Cl 2 + 2e —-> 2Cl 

 

3) – The mcm is calculated and with this value the electrons are balanced, modifying the number of atoms;

mcm between 6 and 2 = 6:

 —-> IO  + 6e 6: 6 = 1 (no need to multiply)
3Cl 2 + 6e —-> 6Cl  6: 2 = 3 (multiplier coefficient)

4) – Oxygen is balanced with 2 O molecules and water hydrogen with OH  ions (basic environment), checking the balance of the charges:

 + 6OH  —-> IO  + 6e + 3H 2 O charges: -7 ¾¾® -7
3Cl 2 + 6e —-> 6Cl  charges: -6 ¾¾® -6

5) – It is added member to member with the appropriate simplifications:

 + 6OH  —-> IO  + 6e + 3H 2 O
3Cl 2 + 6e —-> 6Cl 
————————————————– ———-
 + 3Cl 2 + 6OH  —-> IO 3 + 6Cl  + 3H 2 O

6) – If required, the net ionic reaction is transformed into a balanced molecular reaction:

3Cl 2 + KI + 6KOH v 6KCl + KIO 3 + 3H2O .

 The same procedure is used in reactions that take place in a neutral environment :

AsO 3- + I 2 + H 2 O —-> AsO 3- + H + + I 

1) – The oxidation and reduction half-reactions are written separately:

AsO 3- —-> AsO 3-

2 —-> 2I 

2) – Atoms are balanced and electrons in motion are indicated:

AsO 3- —-> AsO 3- + 2e

2 + 2e —-> 2I 

3) – Since the number of moving electrons is equal, no balancing must be carried out.

4) – Oxygen is balanced in the reactants with 2 O and in the hydrogen products, thus added, with + ions , controlling the balance of the charges:

AsO 3- + H 2 O —-> AsO 3- + 2e + 2H +

2 + 2e —-> 2I 

5) – It is added member to member with the appropriate simplifications:

AsO 3- + H 2 O —-> AsO 3- + 2e + 2H +
2 + 2e —-> 2I 
————————————————– ———-
AsO 3- + I 2 + H 2 O —-> AsO 3- + 2I  + 2H +

6) – If required, the net ionic reaction is transformed into a balanced molecular reaction:

KAsO 3 + I 2 + H 2 O —-> KAsO 4 + 2HI .

 

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