SSS 2: REDOX REACTION

REDOX REACTION is a reaction involving both oxidation and reduction occuring at the same time. Most chemical reactions are redox reactions except neutralization reactions and precipitation reactions.

Redox reactions can be explained in four different ways:

1. Redox in terms of electron transfer

2. Redox in terms of change in oxidation number (O.N.)

3. Redox in terms of addition or removal of oxygen and hydrogen 

4. Redox in terms of addition or removal of electronegative and electropositive elements 

      Redox in terms of electron transfer 

(I) Oxidation is the process of electron loss

(II) Reduction is the process of electron gain

(a)   H₂(g) + Cu²⁺ (aq) → Cu(s) + 2H⁺(aq)

O.N.:  0             +2                0             +1 

(I) Cu²⁺ (aq) is undergoing reduction (is reduced). Hence Cu²⁺ (aq) is an oxidizing agent.

(II) H₂(g) is undergoing oxidation (is oxidized), hence H₂(g) is a reducing agent 

(b)   H₂S(g) + Cl₂(g) → S(s) + 2HCl(g)

O.N.: +1 -2        0              0         +1 -1

 

(I) Cl₂(g) is undergoing reduction (is reduced); hence Cl₂(g) is an oxidizing agent.

(II) H₂S(g) is undergoing oxidation (is oxidized); hence, H₂S(g) is a reducing agent.

 Redox in terms of change in oxidation number

(I) Oxidation is Increase in oxidation number

(II) Reduction is decrease in oxidation number

(a) Zn(s) + Cu²⁺ (aq) →Zn²⁺(aq) + Cu(s)

O.N.: 0           +2                  +2              0 

(I) Zn(s) is undergoing oxidation (is increased from 0 to +2), hence Zn(s) is a reducing agent 

(II) Cu²⁺ (aq) is undergoing reduction (ON is reduced from +2 to 0), Hence Cu²⁺ (aq) is an oxidizing agent.

(b) Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

O.N.:  0           +1 -1            +2 -1               0

 

(I) Mg(s) is undergoing oxidation (is oxidized from 0 to +2); hence, Mg(s) is a reducing agent.

(II) HCl(aq) is undergoing reduction (is reduced from +1 to 0); hence HCl(aq) is an oxidizing agent.

Redox in terms of addition or removal of oxygen and hydrogen

(I) Oxidation is the addition of oxygen to a substance and the removal of hydrogen 

(II) Reduction is the addition of hydrogen to a substance and the removal of oxygen 

(a) CuO(s) + H₂(g) → Cu(s) + H₂O(g)

(I) CuO is undergoing reduction (is reduced); hence CuO is an oxidizing agent

(II) H₂(g) is undergoing oxidation (is oxidized); Hence, H₂(g) is a reducing agent 

(b) 2Mg(s) + O₂(g) → 2MgO(s)

(I) Mg(s) is undergoing oxidation (is oxidized); hence, Mg(s) is a reducing agent 

(II) O₂(g) is undergoing reduction (is reduced); hence, O₂(g) is an oxidizing agent 

Redox in terms of addition or removal of electronegative and electropositive elements 

(I) Oxidation is the addition of electronegative element to a substance and the removal of electropositive element

(II) Reduction is the addition of electropositive element to a substance and the removal of electronegative element.

(a) 2FeCl₂(aq) + Cl₂(g)  → 2FeCl3 (aq)

(I) FeCl₂ is undergoing oxidation ( is oxidized); hence, FeCl₂ is a reducing agent

(II) Cl₂ is undergoing reduction (is reduced); hence, Cl₂ is an oxidizing agent

                      Oxidizing Agent

An oxidizing agent is a substance which accepts electrons or is an electron acceptor.e.g. Chlorine, Fluorine, Conc. H2SO4, Conc. HNO3, KMnO4, K2Cr2O7 etc

                       Reducing agent 

Reducing agent is a substance which donates electrons or is an electron donor.

Hydrogen, Hydrogen sulphide, active metals etc.

Example: Consider the following redox reactions. State the substance that is

(I) Oxidized (II) Reduced (III) Oxidizing agent (IV) Reducing agent 

(a) Ag+ (aq) + Cl-(aq) → AgCl(s)

(b) CuO(s) + CO(g) → Cu(s) + CO2(g)

(c) Cr2O72- (aq) + 14H+(aq) + 6e- → 2Cr3+(aq) + 7H2O(l)

                             Solution

(a)    Ag+ (aq) + Cl-(aq) → AgCl(s)

O.N.: +1                -1                +1 -1

The O.N. of Ag+ (aq) = +1

The O.N. of Ag in AgCl (s) = +1

The O.N. of Cl- (aq) = -1

The O.N. of Cl in AgCl (s) = -1

The oxidation numbers of each substance remain unchanged, hence, the reaction is not a redox reaction.

(b)   CuO(s) + CO(g) → Cu(s) + CO2(g)

O.N.: +2 -2       +2 -2         0        +4 -2

(I) CO is oxidized; oxidation number of C in CO increases from +2 to +4

(II) CuO is reduced; oxidation number of Cu in CuO is reduced from +2 to 0

(III) CuO is an oxidizing agent; it undergoes reduction 

(IV) CO is a reducing agent; it undergoes oxidation

(c) Cr2O72- (aq) + 14H+(aq) + 6e- → 2Cr3+(aq) + 7H2O(l)

O.N.: +6  -2       +1  →   +3           +1  -2

(I) H+ is oxidized; it undergoes oxidation 

(II) Cr2O72- (aq) is reduced

(III)  Cr2O72- (aq) is an oxidizing agent; it undergoes reduction 

(IV) H+  is a reducing agent; it undergoes oxidation

                 Oxidation number 

Oxidation number, also known as oxidation state, is the electrical charge it appears to have as determined by a set of arbitrary rules.

Rules for determining oxidation number 

The set of rules used for the determination of oxidation number are as follows:

1. The oxidation number of all elements in the free state is zero. Eg. Na, Cl2, Mg, P4, H2 etc

2. The oxidation number of an ion is the charge on it. The oxidation number of K+, Ca2+, Al3+, Cl- and S2-are respectively +1, +2, +3, -1 and -2.

In the case of radicals, oxidation number is the algebraic sum of the oxidation numbers of all the elements in the ions.

The oxidation numbers of NH4+, OH-, NO3-, SO42-, PO43- and MnO4- are +1, -1, -1, -2, -3 and -1 respectively.

Example: For OH- 

(O.N. of O) + (O.N. of H) = (O.N. of OH-)

      (-2)        +        (+1)       =   -1

3. The algebraic sum of the oxidation number of all the elements in a compound is zero. e.g. AlCl3.

(O.N. of Al) + 3(O.N. of Cl) = (O.N. of AlCl3)

     +3.              3(-1)              =  0

4. Oxidation number of oxygen in a compound is -2 but in peroxides, it is -1

5. Oxidation number of hydrogen is +1 but in hydrides such has KH, NaH, CaH2 etc is -1

Example: Determine the oxidation number of the underlined elements 

(I) NH4+ (II) NO3- (III) HNO3 (IV) SO42-

(V) H2SO4 (VI) KMnO4 (VII) MnO4- (VIII) K2Cr2O7 (IX) KClO3  (X) H2SO3 (XI) NaMnO4 (XII) Na2ZnO2 (XIII) HCO3- (XIV) [K3Fe(CN)6, where C= -4, N=+3]

Solution:

(I) NH4+                          (II) NO3- 

NH4+ = +1                            NO3- = -1

N + 4(+1) = +1                      N + 3(-2) = -1

N+4=1                                   N - 6 = -1

N = 1- 4                                 N = -1 + 6

N= -3                                     N= +5

(III) HNO3

HNO3 = 0

(1x1) + N + 3(-2) = 0

1 + N - 6 = 0

N= 6-1

N= +5

(IV) SO42-

SO42- = -2

S + 4(-2) = -2

S - 8 = -2

S = -2 + 8 

S = +6

(V) H2SO4

H2SO4 = 0

(1x2) + S + 4(-2) = 0

2 + S - 8 = 0

S = 8 - 2

S = +6

(VI) KMnO4

KMnO4 = 0

(1x1) + Mn + 4(-2) = 0

1 +  Mn  - 8 = 0

Mn = 8 - 1

Mn = +7

(VII) MnO4- 

(VIII) K2Cr2O7 

(IX) KClO3 

 (X) H2SO3

(XI) NaMnO4

                     Class activities 

Consider the following redox reactions. State the substance that is

(I) Oxidized (II) Reduced (III) Oxidizing agent (IV) Reducing agent 

(a) Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu (s)

(b) MnO2(s) + 4HCl(aq) → MnCl2(aq) + 2H2O(l) + Cl2(g)