Balance the following redox reactions by ion-electron method:
(a)MnO4- (aq) + I - (aq) → MnO2(s) + I2(s) (in basic medium)
(b) MnO4- (aq) + SO2(g) → Mn2+(aq) +HSO4- (aq) (in acidic solution)
(c) H2O2(aq)+Fe2+(aq) → Fe3+ (aq) + H2O (l) (in acidic solution)
(d) Cr2O72-+ SO2(g) → Cr3+ (aq) +SO42- (aq) (in acidic solution)
(a)MnO4- (aq) + I - (aq) → MnO2(s) + I2(s) (in basic medium)
(b) MnO4- (aq) + SO2(g) → Mn2+(aq) +HSO4- (aq) (in acidic solution)
(c) H2O2(aq)+Fe2+(aq) → Fe3+ (aq) + H2O (l) (in acidic solution)
(d) Cr2O72-+ SO2(g) → Cr3+ (aq) +SO42- (aq) (in acidic solution)
Solution and Explanation
(a) Step 1: The two half-reactions involved in the given reaction are:
Oxidation half-reaction: I(aq) \(\xrightarrow[]{}\) I2( s)
Reduction half-reaction: MnO4- (aq) \(\xrightarrow[]{}\) MnO2(aq)
Step 2: Balancing I in the oxidation half-reaction, we have:
2I(aq)\(\xrightarrow[]{}\)I2(s)
Now, to balance the charge, we add 2 e - to the RHS of the reaction.
2I−(aq)\(\xrightarrow[]{}\) I2( s)+2e−
Step 3: In the reduction half-reaction, the oxidation state of Mn has reduced from +7 to +4. Thus, 3 electrons are added to the LHS of the reaction.
MnO4−+3e−\(\xrightarrow[]{}\)MnO2(aq)
Now, to balance the charge, we add 4 OH- ions to the RHS of the reaction as the reaction is taking place in a basic medium.
MnO4−(aq)+3e−\(\xrightarrow[]{}\)MnO2(aq)+4OH−
Step 4: In this equation, there are 6O atoms on the RHS and 4O atoms on the LHS. Therefore, two water molecules are added to the LHS.
MnO4−(aq)+2H2O\(\xrightarrow[]{}\)MnO2(aq)+4OH−
Step 5: Equalizing the number of electrons by multiplying the oxidation half-reaction by 3 and the reduction half-reaction by 2, we have:
6I−(aq)\(\xrightarrow[]{}\)3I2( s)+6e−
2MnO4−(aq)+4H2O+6e−→2MnO2(aq)+8OH−(aq)
Step 6: Adding the two half-reactions, we have the net balanced redox reaction as:
6I−(aq)+2MnO4(aq)+4H2O\(\rightarrow\)2MnO2(aq)+8OH−(aq)
(b) Following the steps as in part (a), we have the oxidation half-reaction as: And the reduction half reaction as:
SO2( g)+2H2O(l)\(\xrightarrow[]{}\)HSO4−(aq)+3H+(aq)
Multiplying the oxidation half-reaction by 5 and the reduction half-reaction by 2, and then by adding them, we have the net balanced redox reaction as:
(c) Following the steps as in part (a), we have the oxidation half-reaction as:
The reduction half-reaction as:
Fe2+(aq)\(\xrightarrow[]{}\)Fe3+(aq)+e−
Multiplying the oxidation half-reaction by 2 and then adding it to the reduction half-reaction, we have the net balanced redox reaction as:
H2O2(aq)+2Fe2+(aq)+2H+(ac)\(\xrightarrow[]{}\)2Fe(2q)3++2H2O(i)
(d) Following the steps in part (a), we have the oxidation half-reaction as:
SO2( g)+2H2O(1)\(\xrightarrow[]{}\)SO24(2q)+4H+(aq)+2e−
And the reduction half reaction as:
Cr2O72−(aq)+14H+(aq)+6e−\(\xrightarrow[]{}\)Cr3+(aq)+7H2O(i)
Multiplying the oxidation half-reaction by 3 and then adding it to the reduction half-reaction, we have the net balanced redox reaction as:
Cr2O72-(aq)+3SO2( g)+2H+(aq)\(\xrightarrow[]{}\)2Cr3+(a9)+3SO2−4(aq)+H2O(l)
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Concepts Used:
Redox Reactions
Redox Reaction:
Redox reactions are chemical reactions where oxidation and reduction take place simultaneously. In this type of reaction, there is a gain of electrons for one chemical species while the other loses electrons or simply involves transfer of electrons. The species that loses electrons is oxidized while the one that gains electrons is reduced.
Types of Redox Reactions:
Redox reactions can be differentiated into 4 categories namely combination reactions, decomposition reactions, displacement reactions, and disproportionation reactions. Each is explained separately below:
Combination Reaction:
In this, the molecules combine to form new compounds. For example, when magnesium reacts to nitrogen.
Decomposition Reaction:
Opposite to the combination reaction, here there is a breakdown of compounds to simpler substances. For example, electrolysis of water.
Displacement Reaction:
In this, the more reactive metal will displace the less reactive one in a chemical reaction. The reactivity of an element is represented in a series called the reactivity series (arranged in decreasing order of reactivity) which makes it easier to determine the chemical reaction and its products.
Disproportionation Reaction:
This is a peculiar type of reaction where an element showing a particular oxidation state will be oxidized and reduced simultaneously. Another thing to note is that these reactions will always have an element that can exhibit three oxidation states.