An Ag/AgCl electrode dipping into 1.00 M HCI has a standard reduction potential of +0.2223 V. The half-reaction is AgCl(s) + e = Ag(s) +Cl(aq) A second Ag/AGCI electrode is dipped into a solution containing Cl at an unknown concentration. The cell generates a potential of 0.0402 V, with the electrode in the solution of unknown concentration having a negative charge. What is the molar concentration of Cl in the unknown solution? [CI] = M

Fe (s) +3 Au (aq) ^ + Fe log ^ 3+ +3 Au (s) For the first part, identify the oxidation half-reaction , identify the reduction half- reaction, identify the oxidizing agent, and identify the reducing agent under standard conditions. Next, calculate the standard cell potential, E for the reaction . cell E^ red +1.69 V Half-reaction Au^ + (aq) + e^ - Au (s) Fe^ 3+ (aq) + 3 e^ - Fe (s) -0.040 V Now , calculate the non- standard cell potential , when [Au^ + ]=0.0015 M and the [Fe^ 3+ ]= 0.033 M. After finding the cell potential , calculate the Gibb's free energy , AG Finally, calculate the equilibrium constant , K. Express your answers in two significant digits . Show your work for partial credit . Under non -standard conditions , is this reaction spontaneous or nonspontaneous ? Are reactants or products favored at equilibrium under non -standard conditions ?

Calculate Eº, E, and AG for the following cell reaction: 2+ 2+ Mg(s) + Sn²+ (aq) = Mg²+ (aq) + Sn (s) 2+ 2 where [Mg²+] = [Mg²+] = 0.040 M and [Sn²+] = 0.030 M. Round each of your answers to 4 significant digits. Note: Reference the Standard reduction potentials at 25 °C table for additional information. Note: The Faraday constant is 9.649 × 10ª Part: 0 / 3 Part 1 of 3 Eº = v V x10 J V. mol S I

A galvanic cell consists of a Ni/Ni²+ half-cell and a Co/Co²+ half-cell with initial concentrations: [Ni²+] = 0.80M and [Co²+] = 0.20 M at 298.15 K. half-reaction Eº (V) Ni²+ (aq) +2e→Ni(s) -0.25 Co²+ (aq) +2e→Co(s) -0.28