MEDIUM

Earn 100

A hydrogen electrode placed in a buffer solution of ${CH}_{3} COONa$ and acetic acid in the ratios $x : y$ and $y : x$ has electrode potential values $E_{1}$ volt and $E_{2}$ volt respectively at $25^{0} C$ . Calculate the $pKa$ value of acetic acid in terms of $E_{1}$ and $E_{2}$ ( $E_{1}$ and $E_{2}$ are oxidation potential).

Important Questions on Electrochemistry

HARD

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

The pressure of

H_{2}

required to make the potential of hydrogen electrode zero in pure water at 298 K is:

HARD

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

In the cell,

Pt (s) |H_{2} (g, 1 bar)| HCl (aq) |AgCl (s)| Ag (s) |Pt (s),

the cell potential is

0 .92 V

when a

10^{- 6}

molar

HCl

solution is used. The standard electrode potential of

Ag | AgCl | {Cl}^{-}

electrode is:

(Given,

\frac{2 .303 RT}{F} = 0 .06 V

298 K

)

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

The correct representation of Nernst’s equation for half- cell reaction

{Cu}^{2 +} (aq) + e^{-} \to {Cu}^{+} (aq)

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

For the given cell; $Cu (s) |{Cu}^{2 +} (C_{1} M)| |{Cu}^{2 +} (C_{2} M)| Cu (s)$

change in Gibbs energy $(∆ G)$ is negative, it :

HARD

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

For the electrochemical cell,

Mg(s) |{Mg}^{2+} (aq,1 M)| |{Cu}^{2+} (aq,1 M)| Cu (s)

the standard emf of the cell is

2 .70 V

300 K .

When the concentration of

{Mg}^{2+}

is changed to

x M,

the cell potential changes to

2 .67 V

300 K .

The value of

x

is ________.

(given,

\frac{F}{R} {=11500 K V}^{- 1},

where

F

is the Faraday constant and

R

is the gas constants,

ln (10) = 2 .30

)

HARD

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

The photoelectric current from Na (work function, $w_{0} = 2.3 eV$ ) is stopped by the output voltage of the cell

$Pt (s) ‖ H_{2} (g, 1 bar) | HCl (aq \cdot, pH = 1) | AgCl (s) ∣ Ag (s)$
the $p H$ of aq. $HCl$ required to stop the photoelectric current from $K (w_{0} = 2.25 eV)$ , all other conditions remaining the same, is $\dots \dots \dots . \times 10^{- 2}$ (to the nearest integer).

Given $2.303 \frac{RT}{F} = 0.06 V; E_{AgCl / Ag / C l}^{0} = 0.22 V$

HARD

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

For an electrochemical cell

Sn (s) |{Sn}^{2 +} (aq, 1 M)| |{Pb}^{2 +} (aq, 1 M)| Pb (s)

the ratio

\frac{[{Sn}^{2 +}]}{[{Pb}^{2 +}]}

when this cell attains equilibrium is _______.
(Given:

E_{{Sn}^{2 +}| Sn}^{0} = - 0 .14 V, E_{{Pb}^{2 +}| Pb}^{0} = - 0 .13 V, \frac{2 .303 RT}{F} = 0 .06 \log 2.15 = \frac{1}{3}

)

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

Find the electrode potential for the given half-cell reaction given below at $pH = 4$ .

$2 H_{2} O ⟶ O_{2} + 4 H^{+} + 4 e^{-}; E_{}^{°} = - 1.23 V$

(Given, $R = 8.314 J K^{- 1} {mol}^{- 1}$ , temperature $= 298 K$ , oxygen is under standard atmospheric pressure of $1 bar$ )

HARD

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

In the electrochemical cell:

Zn |{ZnSO}_{4} (0 .01 M)| |{CuSO}_{4} (1 .0 M)| Cu,

the emf of this Daniel cell is

E_{1}

. When the concentration

{ZnSO}_{4}

is changed to

1 .0 M

and that of

{CuSO}_{4}

changed to

0 .01 M,

the emf changes to

E_{2}

. From the following, which one is the relationship between

E_{1}

and

E_{2}

? (Given,

\frac{RT}{F} = 0 .059

)

HARD

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

The standard cell potential for

Zn |{Zn}^{2 +}| |{Cu}^{2 +}| Cu

is 1.10 V. When the cell is completely discharged,

\log [{Zn}^{2 +}] / [{Cu}^{2 +}]

is closest to

HARD

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

The electrode potential of a hydrogen electrode at

pH = 10

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

The voltage of the cell consisting of

Li (s)

and

F_{2} (g)

electrodes is

5.92 V

at standard condition at

298 K

. What is the voltage if the electrolyte consists of

2 M LiF

.
Given that,

\ln 2 = 0 .693, R = 8 .314 {JK}^{- 1} {mol}^{- 1} F = 96500 C {mol}^{- 1}

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

The electrode potentials for ${Cu}^{2 +} (aqueous) + e^{-} ⟶ {Cu}^{+} (aqueous)$ and ${Cu}^{+} (aqueous) + e^{-} ⟶ Cu (s)$ are $+ 0.15 V$ and $+ 0.50 V,$ respectively. The value of $E °_{{Cu}^{2 +} / Cu}$ will be:

HARD

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

298 K

, the standard reduction potentials are

1.51 V

for

{M n O}_{4}^{-} | M n^{2 +}

1.36 V

for

C l_{2} | C l^{-}

1.07 V

for

B r_{2} | B r^{-}

0.54 V

for

I_{2} | I^{-}

. At

pH = 3

, permanganate is expected to oxidize:

(\frac{R T}{F} = 0.059)

HARD

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

$E_{1}, E_{2}$ and $E_{3}$ are the emf of the following three galvanic cells respectively

(i)

Zn (s) |{Zn}^{2 +} (0.1 M)| |{Cu}^{2 +} (1 M)| Cu (s)

(ii) $Zn (s) |{Zn}^{2 +} (1 M)| |{Cu}^{2 +} (1 M)| Cu (s)$

(iii) $Zn (s) |{Zn}^{2 +} (1 M)| |{Cu}^{2 +} (0.1 M)| Cu (s)$

Which of the following is true?

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

The

EMF

of a galvanic cell consisting of two hydrogen electrodes is

0.17 V

. If the solution of one of the electrodes has

[H^{+}] = 10^{- 3} M

, the

pH

at the other electrode is:

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

Given

E_{{Fe}^{+ 3 / {Fe}^{2}}} = + 0.76 V

and

E °_{I_{2} / I^{-}} = + 0.55 V

.
The equilibrium constant for the reaction taking place in galvanic cell consisting of above two electrodes is

[\frac{2.303 RT}{F} = 0.06]

HARD

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

What would be the electrode potential for the given half-cell reaction at

p H = 5 ?

2 H_{2} O \to O_{2} + 4 \overset{\oplus}{H} + 4 e^{-}; E_{r e d}^{0} = 1.23 V

(R = 8 .314 J {mol}^{- 1} K^{- 1}; Temp = 298 K; oxygen under standard atm . pressure of 1 bar)

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

To find the standard potential of

M^{3 +} / M

electrode, the following cell is constituted:

Pt / M / M^{3 +} (0 .001 mol L^{- 1}) / {Ag}^{+} (0 .01 mol L^{- 1}) / Ag

The emf of the cell is found to be

0.421 volt

298 K

. The standard potential of half-reaction

M^{3 +} + 3 e^{-} \to M

298 K

will be:

(Given:

E_{\frac{{Ag}^{+}}{Ag}}^{⊝}

298 K = 0.80 volt

)

MEDIUM

Chemistry>Physical Chemistry>Electrochemistry>Nernst Equation

For the following electrochemical cell at

289 K,

Pt (s) | H_{2} (g, 1 bar) | H^{+} (aq, 1 M) | | M^{4 +} (aq .) {, M}^{2+} (aq .) | Pt (s)

E_{c e l l} = 0.092 V

when

\frac{[M^{2 +} (a q .)]}{[M^{4 +} (a q .)]} = 10^{x}

Given:

E_{M^{4 +} / M^{2 +}}^{0} = 0.151 V; 2.303 \frac{R T}{F} = 0.059

The value of

x

is -

A hydrogen electrode placed in a buffer solution of CH3COONa and acetic acid in the ratios x:y and y:x has electrode potential values E1 volt and E2 volt respectively at 250C. Calculate the pKa value of acetic acid in terms of E1 and E2 (E1 and E2 are oxidation potential).

Important Questions on Electrochemistry

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