- Why does the cell potential drop to zero at equilibrium?
- At equilibrium, the reaction quotient Q equals the equilibrium constant K. Substituting Q = K into the Nernst equation gives E = E° − (RT/nF) ln K. Since the standard relationship is E° = (RT/nF) ln K, these two terms cancel, resulting in E = 0. A zero cell potential means there is no net driving force for the reaction in either direction; the system is at balance.
- What's the practical use of the Nernst equation?
- The Nernst equation is fundamental to the operation of pH meters, ion-selective electrodes, and all potentiometric sensors. These devices measure voltage (E) to determine the concentration of an ion in solution. By holding E°, n, T, and other concentrations constant, the measured potential varies logarithmically with the concentration of the target ion, as described by the Nernst equation.
- Does the Nernst equation apply to both galvanic and electrolytic cells?
- Yes, it applies to both. For a galvanic (voltaic) cell, E is positive, indicating a spontaneous reaction that produces electrical work. For an electrolytic cell, an applied external voltage forces a non-spontaneous reaction; the Nernst equation calculates the minimum voltage that must be applied to drive the reaction, which corresponds to a negative calculated E for the forward cell reaction.
- Why do we use natural log (ln) in the equation, and can we use log base 10?
- The thermodynamic derivation naturally yields the natural logarithm. However, a common practical form uses base-10 log: E = E° − (2.303RT/nF) log Q. The constant 2.303 is the conversion factor (ln 10). Both forms are correct; the simulator uses the natural log form as it is the most fundamental, but the behavior is identical in principle.