Chapter 21, Problem 21.42P

(a)

Interpretation Introduction

Interpretation:

The given cell reaction has to be balanced first and then, identified that whether it is spontaneous by using ${\text{E}}^{\text{o}}{}_{\text{cell}}$ calculation.

Concept Introduction:

An electrochemical cell is a device in which a redox reaction is used to convert chemical energy into electrical energy. Such device is also known as the galvanic or voltaic cell.

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.

Anode: The electrode where the oxidation occurs is called as an anode. It is a negatively charged electrode.

Cathode: The electrode where reduction occurs is called as a cathode. It is a positively charged electrode.

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.

The Standard Gibb’s free energy change and the standard cell potential are related as followed:

$\text{Δ}°\text{G}\text{=}\text{-nFE}{°}_{\text{cell}}$

n - Number of electrons involved per equivalent of the net redox reaction in the cell

F - Faraday’s Constant (96500 C)

$\text{E}{°}_{\text{cell}}$ - Standard cell potential.

Steps in balancing organic redox reactions:

1) Divide the overall reaction into an oxidation half-reaction and a reduction half-reaction

2) Balance atoms other than $\text{O}$ and $\text{H}$.

3)  Balance $\text{O}$ by adding ${\text{H}}_{\text{2}}\text{O}$ as needed

4) Balance $\text{H}$ by adding ${\text{H}}^{\text{+}}$ ion at the required side.

5) Balance charges by adding, as needed number of electrons,

6) Multiply the oxidation half-reaction with the coefficient of electrons in the reduction part.

7) Multiply the reduction half-reaction with the coefficient of electrons in the oxidation part.

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) For the reaction in acidic medium, the presence of $H^{+}$ ion allowed.

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.

(b)

Interpretation Introduction

Interpretation:

The given cell reaction has to be balanced first and then, identified that whether it is spontaneous by using ${\text{E}}^{\text{o}}{}_{\text{cell}}$ calculation.

Concept Introduction:

An electrochemical cell is a device in which a redox reaction is used to convert chemical energy into electrical energy. Such device is also known as the galvanic or voltaic cell.

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.

Anode: The electrode where the oxidation occurs is called as an anode. It is a negatively charged electrode.

Cathode: The electrode where reduction occurs is called as a cathode. It is a positively charged electrode.

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.

The Standard Gibb’s free energy change and the standard cell potential are related as followed:

$\text{Δ}°\text{G}\text{=}\text{-nFE}{°}_{\text{cell}}$

n - Number of electrons involved per equivalent of the net redox reaction in the cell

F - Faraday’s Constant (96500 C)

$\text{E}{°}_{\text{cell}}$ - Standard cell potential.

Steps in balancing organic redox reactions:

1) Divide the overall reaction into an oxidation half-reaction and a reduction half-reaction

2) Balance atoms other than $\text{O}$ and $\text{H}$.

3)  Balance $\text{O}$ by adding ${\text{H}}_{\text{2}}\text{O}$ as needed

4) Balance $\text{H}$ by adding ${\text{H}}^{\text{+}}$ ion at the required side.

5) Balance charges by adding, as needed number of electrons,

6) Multiply the oxidation half-reaction with the coefficient of electrons in the reduction part.

7) Multiply the reduction half-reaction with the coefficient of electrons in the oxidation part.

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) For the reaction in acidic medium, the presence of $H^{+}$ ion allowed.

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.