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Nernst Equation

Question

If the $Δ$ G of a cell reaction, $A g C l + e^{-} \to A g + C l^{-}$ is $- 21.20$ kJ, the standard emf of cell is
$0.220$ V
$0.239$ V
$- 0.110$ V
$- 0.320$ V

A

- 0.320

V

B

0.239

V

C

0.220

V

D

- 0.110

V

Open in App

Solution

Verified by Toppr

$Δ G^{o} = - n F E_{c e l l}$
$- 21200 = - 1 \times 96500 \times E_{c e l l}$
$E_{c e l l} = 0.220 V$

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Similar Questions

Q1

If the

Δ

G of a cell reaction,

A g C l + e^{-} \to A g + C l^{-}

is

- 21.20

kJ, the standard emf of cell is

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Q2

Calculate the useful work done during the reaction in kJ/mol.
$A g (S) + \frac{1}{2} C l_{2} (g) \to A g C l (s)$
Given that $E_{C l_{2} / {C l}^{-}}^{0} = + 1.36 V$
and $E_{A g^{+} | A g C l | C l^{-}}^{0} = + 0.220 V$
$P_{C l_{2}} = 1 atm$ and T = 298 K

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Q3

The standard reduction potential of two reactions are given.

A g C l_{(s)} + e^{-} \to A g_{(s)} + C l_{(a q)}^{-}

;

E^{⊖} = 0.22

V (i)

A g_{(a q)}^{+} + e^{-} \to A g_{(s)}; E^{⊖} = 0.80

V (ii)
The solubility product of

A g C l

under standard conditions of temperature is:

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Q4

Calculate the EMF of the cell at 298 K.

P t | H_{2} (1 a t m) | N a O H (x M) | | N a C l (x M) | A g C l (s) | A g

If

E_{C l^{-} / A g C l / A g}^{0} = 0.222 V

.

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Q5

C d (s) | C d C l_{2} (0.10 M) | A g C l (s) | A g (s)

The EMF of the above cell is 0.6315 V at

0^{\circ} C

and 0.6753 V at

25^{\circ} C

. The

Δ H

of reaction in kJ at

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is:

View Solution