Equilibrium Constant from Nernst Equation

Q. The emf of the cell $Zn|Z{n}^{2+}\left(0.1M\right)||F{e}^{2+}|Fe\left(0.01M\right)$ is $0.2905V$. Equilibrium constant for the cell reaction is

1. 
2. 
3. 
4. 

Q.

Electra needs to answer this question to master this topic.
Help her calculate the equilibrium constant of the following reaction:
$Cu\left(s\right)+2A{g}^{+}\left(aq\right)\to C{u}^{2+}\left(aq\right)+2Ag\left(s\right)$
$E_{cell}^0=0.46V$

The options below are $logK$ values where $K$ is the equilibrium constant.

1. 15.93
2. 19.5
3. 15.59

4. 15.45

Q. Electra has mastered the basics of Nernst equation. She is now studying equilibrium constants relation with Nernst equation.

Electra was performing experiments with Nernst with a Daniel Cell. She noticed that when the reaction reaches equilibrium, the cell potential goes to 0.

She asked Nernst why this happened. His reply was

1. Because all the ions have been depleted.
2. Because there is change in temperature
3. Because there is no displacement reaction
4. Because there is no change in concentration of Cu²⁺ and Zn²⁺
   

Q. $E^{\circ }$ for the cell is $Zn|Z{n}^{2+}\left(aq\right)||C{u}^{2+}\left(aq\right)|Cu$ at ${25}^{\circ }C$, the equilibrium constant for the reaction $Zn+C{u}^{2+}\left(aq\right)\rightleftharpoons Cu+Z{n}^{2+}\left(aq\right)$ is of the order of

1. 10^-37
2. 10^-28
3. 10⁺¹⁸
4. 10⁺¹⁷

Q. In a cell that utilizes the reaction $Z{n}_{\left(s\right)}+2H^{+}\left(aq\right)\to Z{n}^{2+}\left(aq\right)+H_{2\left(g\right)}$ addition of $H_{2}S{O}_4$ to cathode compartment, will

1. Increase the E and shift equilibrium to the right
2. Lower the E and shift equilibrium to the right
3. Lower the E and shift equilibrium to the left
4. Increase the E and shift equilibrium to the left

Q. If the $E_{cell}^0$ for a given reaction has a negative value, which of the following gives the correct relationships for the values of $\Delta G^0$ and $K_{eq}$ ?

1. $\Delta G^{\circ }>0;K_{eq}<1$
2. $\Delta G^{\circ }>0;K_{eq}>1$
3. $\Delta G^{\circ }<0;K_{eq}>1$
4. $\Delta G^{\circ }<0;K_{eq}<1$

Q. In the acid-base titration
$[H_{3}P{O}_4\left(0.1M\right)+NaOH\left(0.1M\right)]$ , the e.m.f of the solution is measured by coupling this electrode with a suitable reference electrode. When an alkali is added, the pH of the solution is in accordance with the equation:
$E_{\text{cell}}=E_{\text{cell}}^{\circ }+0.059\text{pH}$
For $H_{3}P{O}_4:K{a}_1={10}^{-3},K{a}_2={10}^{-8},K{a}_3={10}^{-13}$
What is the e.m.f of the cell at the ${\text{II}}^{\text{nd}}$ end point of the titration if $E_{\text{cell}}^{\circ }$ at this stage is $1.3805\text{V}$

Q. Maxium work that can be done by a chemical cell is equal to the change in gibbs free energy.

1. True
2. False

Q. EMF of the cell, $Ag|AgN{O}_3\left(0.1M\right)|KBr\left(1N\right),AgBr\left(s\right)|Ag$ is $-0.6\text{V}$ at $298\text{K}$
Calculate the $K_{sp}$ of $AgBr$ at $298\text{K}$.

1. $K_{sp}=4.0\times {10}^{-12}$
2. $K_{sp}=3.5\times {10}^{-12}$
3. $K_{sp}=4.8\times {10}^{-12}$
4. $K_{sp}=3.8\times {10}^{-12}$

Q. The equilibrium constant for the following reaction is ${10}^{30}$. Calculate $E^{\circ }$ (in volts) for the cell at 298 K.
$2X_2\left(s\right)+3Y^{2+}\left(aq\right)$→$2X_2^{3+}\left(aq\right)+3Y\left(s\right)$

1. +0.295 V
2. -0.295 V
3. +0.105 V
4. +0.0985 V