Chemistry

Problem

A galvanic cell employs the half-reactions Ag

^{+}

+ e

^{-}

→ Ag

E^{\circ}=+0.80\,\text{V}

^{2+}

+ 2 e

^{-}

→ Cu

E^{\circ}=+0.34\,\text{V}

Initially both ion concentrations are 1.0 M. After the cell has operated for some time,

[\text{Ag}^{+}]=0.050\,\text{M}

and

[\text{Cu}^{2+}]=2.0\,\text{M}

. Which statement best describes how the instantaneous cell potential

E

compares with the standard potential

E^{\circ}_{\text{cell}}

? a)

E > E^{\circ}_{\text{cell}}

because the

\log Q

term is added to

E^{\circ}

in the Nernst equation. b)

E > E^{\circ}_{\text{cell}}

because

Q = \dfrac{[\text{Ag}^{+}]^{2}}{[\text{Cu}^{2+}]}

, which is less than 1 under the given conditions. c)

E < E^{\circ}_{\text{cell}}

because the reaction quotient is greater than 1. d)

E = E^{\circ}_{\text{cell}}

because the increases and decreases in ion concentrations effectively keep

Q

equal to 1.