For a general electrochemical reaction of the type
`alphaA+bB overset(n e^(-))(hArr)cC+dD`
Nernst equation can be written as

Options

A. `E_("cell") = E_("cell")^(@)+(RT)/(nF) In ([C]^(c)[D]^d)/([A]^(a)[B]^(b)`
B. `E_("cell") = E_("cell")^(@)-(RT)/(nF) In ([C]^(c)[D]^d)/([A]^(a)[B]^(b)`
C. `E_("cell") = E_("cell")^(@)-(RT)/(nF) In ([A]^(a)[B]^b)/([C]^(c)[D]^d)`
D. `E_("cell") = -E_("cell")^(@)-(RT)/(nF) In ([C]^(c)[D]^d)/([A]^(a)[B])`

Answer

B

Explanation

The `emf` of a cell dependsd on the concentrations of ions and on gas pressures. For that reason, cell `emf` provide a way to measure ion concentrations. For example, the `pH` meter depends on the variation of cell `emf` with hydrogen-ion concentraion. We can relate cell `emfs` for various concentrations of ions and vrious gas pressures to standard electrode potential by means of an equation first derived by the Ferman chemist Walther Nernst:
`E_("cell") = E_("cell")^(@) - (Rt)/(nF) "ln"([C]^(c)[D]^(d))/([A]^(a)[B]^(b))`
or `E_("cell") = E_("cell")^(@) - (RT)/(nF) ln Q`
We can show from Nernst equation that the cell emf, `E_("cell")`, decreases as the cell reaction proceeds if `E_("cell")^(@)` is positive. As the reaction occurs in the voltaic cell, the concentrations of products increase and the concentrations of reactants decrease. Therefore, `Q` and log `Q` and log `Q` increases. The second term the difference `E_("cell")^(@) - (0.0592//n) log Q`, increases, so that the difference `E_("cell")^(@) - (0.0592//n) log Q`, decreases. Thus, the cell emf, `E_("cell")`, becomes smaller. Eventually the cell emf goes on the other hand , we can increase the `E_("cell")` by increasing the concentrations of reactants and decreasing the concentration of products.