Calculate the cell potential for the following reaction as written at 25.00 °C, given that [Mg2 ] = 0.763 M and [Sn2 ] = 0.0180 M
Mg(s)+Sn^2+(aq)-->Mg^(2+)+Sn(s)
I keep getting 2.57 but it says I'm wrong. Please help me!
http://www.jiskha.com/display.cgi?id=1400465598
To calculate the cell potential for the given reaction, you need to use the Nernst equation. The Nernst equation relates the cell potential to the concentrations of the species involved in the reaction.
The Nernst equation is given by:
Ecell = E°cell - (RT/nF) * ln(Q)
Where:
Ecell is the cell potential
E°cell is the standard cell potential
R is the gas constant (8.314 J/(mol*K))
T is the temperature in Kelvin
n is the number of electrons transferred in the balanced equation
F is the Faraday constant (96,485 C/mol)
Q is the reaction quotient, which is the ratio of the concentrations of the products to the concentrations of the reactants, each raised to the power of their stoichiometric coefficients.
In this case, the balanced equation is:
Mg(s) + Sn2+(aq) --> Mg2+(aq) + Sn(s)
The number of electrons transferred in this reaction is 2.
The standard cell potential, E°cell, can be obtained from reference tables. The given concentrations are [Mg2+] = 0.763 M and [Sn2+] = 0.0180 M.
Now, let's calculate the reaction quotient, Q. It is given by:
Q = ([Mg2+]/[Sn2+])^2
Substituting the given concentrations, we get:
Q = (0.763/0.0180)^2
Simplifying,
Q ≈ 182.36^2
Q ≈ 33280.3696
Now, we need to convert the temperature from Celsius to Kelvin. We have T = 25.00 °C, adding 273.15 to convert to Kelvin:
T = 298.15 K
Finally, we can substitute the given values into the Nernst equation and solve for Ecell:
Ecell = E°cell - (RT/nF) * ln(Q)
Assuming the E°cell value is given, if you provide that, I can help you calculate the cell potential.