Consider: Zn(s)+Cu^2+(aq)---->Zn^2+(aq)+Cu(s) Ecell=+1.10v Which would increase the cell potential? Increase Cu^2+ Decrease Cu^2+ Increase Zn^2+...

The standard cell potential for the given reaction is +1.10V. The standard cell potential is the cell potential at standard conditions of 1M of ions. Changing ion concentrations will affect the cell potential either increasing or decreasing it. The new cell potential can be calculated using the Nernst equation:

`E = E(std) - ((RT)/(nF)) ln(Q)` where:

E(std) is the standard cell potential,

R is the gas constant

T is temperature in Kelvin

n is the number of electrons transferred

F is Faraday's constant and

Q is the reaction constant.

The reaction constant for redox reactions is the ratio between concentrations of the ions (product over reactant). Thus, for this reaction, 

`Q = [(Zn^(2+))/(Cu^(2+))]` .

Incorporating this to the Nernst equation, we get the following equation specific to the given reaction:

 `E = 1.10 - (RT)/(nF)ln(([Zn^(2+)])/([Cu^(2+)]))` .

Using laws of logarithms, the ln portion can be expanded to yield the following:

`E = 1.10 - ((RT)/(nF))ln(Zn^(2+)) + ((RT)/(nF))ln(Cu^(2+))` .

In general, as x increases, the value of ln(x) increases too. Therefore, as the concentration of zinc ions increases, the potential decreases. Meanwhile, as the concentration of copper ions increases, the potential increases.