Plotting ΔG vs Temp for Electrochemical Reaction
My question is, theoretically, what should happen to the voltage as I heat the cell, assuming that the concentrations of the solutions were off?


This question is tough to answer because I don't know what you mean by "off".  So let me ask (and answer) a few related questions that I think will help you.

What should happen to voltage as a function of temperature in an electrochemical cell

The electrical potential of electrochemical cells varies with temperature, even when the cell is at standard conditions.  A 1985 NIST document explains the situation well.  Only the standard hydrogen electrode has a standard potential that does not vary with temperature.  (The SHE potential is defined as 0 volts at every temperature.)  For other cells, usually the assumption of a linear variation with temperature is usually very good.  That is, 

$${\cal E}_T^\circ = {\cal E}_{298}^\circ  + (T-298.15)\cdot \frac{d{\cal E}_{298}^\circ}{dT}$$

Note that this variation is completely unrelated to the Nernst equation.  The same reference gives values for $\frac{d{\cal E}_{298}^\circ}{dT}$ for zinc(II)/zinc metal of 0.119 milliVolts per Kelvin and for tin(II)/tin metal -0.32 milliVolts per Kelvin.

Don't be confused by the $T$ in the Nernst equation

The Nernst equation is 

$$ {\cal E}= {\cal E}^\circ + \frac{RT}{nF}\ln Q$$  

It does contain a $T$ in there, but don't be mislead into thinking that the $T$ you see is the sole means by which a cell potential ${\cal E}$ can vary with temperature.  ${\cal E}^\circ$ also varies with temperature (see above).


  According to the Nernst equation, it seems that the voltage or potential of the cell should not be affected by a change in temperature if the concentrations are equal because ln(0.1/0.1)=0.


Thus, hopefully my answer makes clear that this statement is wrong, because ${\cal E}^\circ$ varies with temperature.