At 25 C, RT/F = (8.314 J/mol/K *298 K) /(96500 C/mol) =0.0257 volts. My textbook lists 0.0592. What am I doing wrong?
Nevermind. They were also converting natural log to log10.
right.
To understand why you are getting a different answer than what is listed in your textbook, let's review the equation you are using and the constants involved.
The equation you are using is known as the Nernst equation, which relates the standard electrode potential (E°) to the actual electrode potential (E) in an electrochemical cell. The Nernst equation is given as follows:
E = E° - RT/nF * ln(Q)
Where:
- E is the actual electrode potential
- E° is the standard electrode potential
- R is the ideal gas constant (8.314 J/mol/K)
- T is the temperature in Kelvin (298 K in this case)
- n is the number of moles of electrons transferred in the balanced redox equation
- F is the Faraday constant (96500 C/mol)
- Q is the reaction quotient
To calculate the E° value, you need to use the equation:
E° = RT/nF * ln(K)
Where:
- K is the equilibrium constant for the redox reaction
In your case, it seems like you are calculating the term RT/F (8.314 J/mol/K * 298 K) / (96500 C/mol) correctly, which equals 0.0257 volts.
The issue seems to be with the value of E° listed in your textbook. It is important to note that E° values are specific to each redox reaction and may vary depending on the specific half-cell reactions used. Therefore, it is possible that the value given in your textbook is different from the one you are calculating based on a different set of redox reactions.
To compare your calculated value with the one in your textbook, you should double-check the specific redox reactions and corresponding equilibrium constants (K) used to obtain the E° value listed. By ensuring that you are using the same set of reactions and constants as your textbook, you should be able to match the given value.
In summary, it is likely that the discrepancy between your calculated value and the one listed in your textbook is due to different redox reactions or equilibrium constants being used. Double-check the specific reactions and constants in your textbook to ensure your comparisons are accurate.