Answer:
The change in cell potential under non-standard conditions can be predicted using the Nernst equation. The Nernst equation relates the measured cell potential to the reaction quotient and allows the accurate determination of equilibrium constants (including solubility constants).
In this case, the concentrations of Al³+ and Ni²+ are not standard. According to the Nernst equation, an increase in the concentration of the oxidizing agent ion in the cathodic side of the system (reduction reaction) should result in an increase in cell potential. Conversely, an increase in the concentration of the reducing agent in the anodic side of the system (oxidation reaction) should result in a decrease in cell potential.
Step-by-step explanation:
The Nernst equation enables the determination of cell potential under non-standard conditions. It relates the measured cell potential to the reaction quotient and allows the accurate determination of equilibrium constants (including solubility constants).
The Nernst equation is given by:
Ecell = E°cell - (RT/nF)lnQ
where Ecell is the cell potential under non-standard conditions, E°cell is the standard cell potential, R is the universal gas constant, T is the temperature in kelvins, n is the number of moles of electrons transferred in the balanced equation for the cell reaction, F is Faraday’s constant, and Q is the reaction quotient.
At room temperature (25°C or 298K), this equation can be simplified to:
Ecell = E°cell - (0.0592/n)logQ
where logQ is the base-10 logarithm of the reaction quotient.
To calculate the change in cell potential using this equation, you need to know the standard cell potential (E°cell), the number of moles of electrons transferred in the balanced equation for the cell reaction (n), and the reaction quotient (Q) under non-standard conditions.