Chapter 21.2, Problem 21.2BFP

Interpretation Introduction

Interpretation:

The cell diagram for given voltaic cell with given half cells has to be drawn. The balanced half-cell reaction and the cell notation for the given cell has to be determined.

Concept Introduction:

Oxidation: The gain of oxygen or the loss of hydrogen or the loss of an electron in a species during a redox reaction is called as oxidation.

Reduction: The loss of oxygen or the gain of hydrogen or the gain of an electron in a species during a redox reaction is called as reduction.

Oxidizing agent: The substance that is oxidized is called as a reducing agent.

Reducing agent: The substance that is reduced is called as an oxidizing agent.

Steps in balancing redox reactions:

1. 1) Divide the overall reaction into an oxidation half-reaction and a reduction half-reaction.
2. 2) Balance atoms other than $\text{O}$ and $\text{H}$.
3. 3) Balance $\text{O}$ by adding ${\text{H}}_{\text{2}}\text{O}$ as needed
4. 4) Balance $\text{H}$ by adding ${\text{H}}^{\text{+}}$ ion at the required side.
5. 5) Balance charges by adding, as needed number of electrons,
6. 6) Multiply the oxidation half-reaction with the coefficient of electrons in the reduction part.
7. 7) Multiply the reduction half-reaction with the coefficient of electrons in the oxidation part.
8. 8) Combine the two half-reactions, cancel out the species that appears on both side, so that number of elements that appear on both sides become equal.
9. 9) For the reaction in acidic medium, the presence of $H^{+}$ ion allowed.
10. 10) For the reaction in basic medium, the $H^{+}$ ion should be neutralized by the addition of required amount of $O{H}^{-}$ ions on both the sides.

11) Cancel out the species that appears on both sides and ensure that the number of atoms on the reactant side is equal to the number of atoms on the product side.

Basic conditions to balance the half-reactions are as follows:

• Balance the oxidation number with electrons.
• Balance the elements except for O and H.
• Balance O with ${\text{H}}_{\text{2}}\text{O}$.
• Balance H with ${\text{H}}^{\text{+}}$ ion.
• Balance the net charge with the electrons.
• Add $\text{O}{\text{H}}^{\text{-}}$ on both sides to cancel all  ${\text{H}}^{\text{+}}$ ions.

Two electrodes connected by a metallic conducting wire, dipped in electrolytes, form an electrochemical cell. Reduction occurs at one electrode and oxidation occurs at the other.

Galvanic cell consists of two half-cells. The redox reaction occurs in these half-cells. The half-cell in which the reduction reaction occurs is known as the reduction half-cell, whereas the half-cell in which the oxidation reaction occurs is known as the oxidation half-cell.

The cell is described using a ‘cell notation’ as follows:

${\text{M}}_{\text{(s)}}{\text{|M}}^{\text{+}}{}_{\text{(aq)}}{\text{||H}}^{\text{+}}{}_{\text{(aq)}}\text{|}\frac{\text{1}}{\text{2}}{\text{H}}_{\text{2}}{}_{\text{(g)}}{\text{;Pt}}_{\text{(s)}}$

The direction of spontaneity is from left to right across the cell notation. Here, M oxidizes to ${\text{M}}^{\text{+}}$ spontaneously, while ${\text{H}}^{\text{+}}$ is reduced to hydrogen gas.

The ‘||’ divides the cell into two half cells. On the left of ‘||’, the oxidation half-cell; on the right the reduction half-cell.

The electrode potential of the oxidation half-cell is written as ${\text{E}}_{{\text{M/M}}^{\text{+}}}$.

The electrode potential of the reduction half-cell is written as ${\text{E}}_{{\text{H}}^{\text{+}}\text{/}\frac{\text{1}}{\text{2}}{\text{H}}_{\text{2}}}$.

The cell potential is calculated as $\text{E}{°}_{\text{cell}}\text{=}\text{E}{°}_{{\text{M/M}}^{\text{+}}}\text{+}\text{E}{°}_{{\text{H}}^{\text{+}}\text{/}\frac{\text{1}}{\text{2}}{\text{H}}_{\text{2}}}$ .