Calculate the standard Gibbs free energy change for the reaction that occurs when a cell consisting of a nickel electrode and a silver electrode is connected, given that the nickel electrode is immersed in a solution containing 0.010 M Ni2+ ions and the silver electrode is immersed in a solution containing 0.20 M Ag+ ions. The standard reduction potential for the Ni2+/Ni half-cell is -0.25 V and for the Ag+/Ag half-cell is +0.80 V.

Question

Calculate the standard Gibbs free energy change for the reaction that occurs when a cell consisting of a nickel electrode and a silver electrode is connected, given that the nickel electrode is immersed in a solution containing 0.010 M Ni2+ ions and the silver electrode is immersed in a solution containing 0.20 M Ag+ ions. The standard reduction potential for the Ni2+/Ni half-c

Answer 1

To calculate the standard Gibbs free energy change, we first need to determine the cell potential  E_cell  under non-standard conditions using the Nernst equation:E_cell = E_cell -  RT/nF  * ln Q where E_cell is the standard cell potential, R is the gas constant  8.314 J/molK , T is the temperature in Kelvin  assuming 298 K , n is the number of moles of electrons transferred in the redox reaction, F is the Faraday constant  96,485 C/mol , and Q is the reaction quotient.First, let's find the standard cell potential  E_cell :E_cell = E_cathode - E_anodeE_cell =  +0.80 V  -  -0.25 V  = 1.05 VNow, let's find the reaction quotient  Q :Q = [Ni2+]/[Ag+]Q =  0.010 M / 0.20 M  = 0.05The balanced redox reaction is:Ni2+ + 2Ag  Ni + 2Ag+So, n = 2 moles of electrons are transferred in the reaction.Now, we can use the Nernst equation to find E_cell:E_cell = 1.05 V -  8.314 J/molK * 298 K  /  2 * 96,485 C/mol  * ln 0.05 E_cell = 1.05 V -  0.0257 V  * ln 0.05 E_cell  1.05 V -  0.0257 V  *  -2.9957 E_cell  1.05 V + 0.077 VE_cell  1.127 VFinally, we can calculate the standard Gibbs free energy change  G  using the following equation:G = -nFE_cellG = -2 * 96,485 C/mol * 1.127 VG  -217,000 J/molSo, the standard Gibbs free energy change for the reaction is approximately -217 kJ/mol.