The standard emf for the cell using the overall cell reaction below is +0.48 V:
$\quad$ $\quad$ Zn(s) + $\mathrm { Ni } ^ { 2 + }$ (aq) ? $\mathrm { Zn } ^ { 2 + }$ (aq) + Ni(s)
The emf generated by the cell when [ $\mathrm { Ni } ^ { 2 + }$ ] = 6.50 mol L^-1 and [ $\mathrm { Zn } ^ { 2 + }$ ] = 0.500 mol L^-1 is -------- $\text { V. }$

Question

The standard emf for the cell using the overall cell reaction below is +0.48 V:

\quad

\quad

Zn(s) +

\mathrm { Ni } ^ { 2 + }

(aq) ?

\mathrm { Zn } ^ { 2 + }

(aq) + Ni(s)
The emf generated by the cell when [

\mathrm { Ni } ^ { 2 + }

] = 6.50 mol L^-1 and [

\mathrm { Zn } ^ { 2 + }

] = 0.500 mol L^-1 is --------

\text { V. }

Quizplus · Accepted Answer

The Nernst equation is:$$E = E^0 - \frac{RT}{nF} \ln Q$$where:- E is the cell potential in volts (V)- E^0 is the standard cell potential in volts (V)- R is the ideal gas constant (8.314 J/mol·K)- T is the temperature in kelvins (K)- n is the number of moles of electrons transferred in the balanced chemical equation- F is the Faraday constant (96,485 C/mol)- Q is the reaction quotientIn this case, the balanced chemical equation is:$$Zn(s)+ $\mathrm { Ni } ^ { 2 + }$ (aq)? $\mathrm { Zn } ^ { 2 + }$ (aq)+ Ni(s)$$and n = 2.The reaction quotient is:$$Q = \frac{[\mathrm { Zn } ^ { 2 + }]}{[\mathrm { Ni } ^ { 2 + }]}$$Substituting the given values into the Nernst equation, we get:$$E = 0.48 V - \frac{0.0257 V}{2} \ln \frac{0.500 mol/L}{6.50 mol/L}$$$$E = 0.48 V - 0.0128 V \ln 0.0769$$$$E = 0.48 V - 0.0128 V (-2.63)$$$$E = 0.51 V$$Therefore, the emf generated by the cell when [ $\mathrm { Ni } ^ { 2 + }$ ] = 6.50 mol L^-1 and [ $\mathrm { Zn } ^ { 2 + }$ ] = 0.500 mol L^-1 is 0.51 V.

The Standard Emf for the Cell Using the Overall Cell \quad

The Standard Emf for the Cell Using the Overall Cell $\quad$