14) Which of the following statement(s)Ware correct for the reaction below. Pt,H_(2)(580torr)vert HCl(0.0016M)Vert Pd^-2(0.28)^-) (E^circ Pd^-1/dd=0.987V)(760torr:1atm) I- The reaction type is spontaneous Vert - Ecell=0.809V III- Keq=2.19times 10^33 A) Only I B) I and II C I and III D) II and III E) I, II and III

To determine which statements are correct for the given reaction, we need to analyze each statement individually.<br /><br />Statement I: The reaction type is spontaneous.<br />To determine if the reaction is spontaneous, we need to calculate the cell potential (E_cell) and compare it to zero. If E_cell is positive, the reaction is spontaneous.<br /><br />The cell potential can be calculated using the Nernst equation:<br /><br />E_cell = E^0_cell - (0.0592/n) * log(Q)<br /><br />where E^0_cell is the standard cell potential, n is the number of moles of electrons transferred, and Q is the reaction quotient.<br /><br />In this case, the standard cell potential (E^0_cell) can be calculated using the standard reduction potentials provided:<br /><br />E^0_cell = E^0_Pd^2+/Pd - E^0_HCl/Cl^- = 0.987V - 1.358V = -0.371V<br /><br />The reaction quotient (Q) can be calculated using the concentrations and partial pressures of the reactants and products:<br /><br />Q = (P_H2/P^0) * (C_HCl/C^0) * (C_Pd^2+/C^0)^2<br /><br />where P^0 is the standard pressure (1 atm) and C^0 is the standard concentration (1 M).<br /><br />Substituting the given values, we have:<br /><br />Q = (580 torr/760 torr) * (0.0016 M/1 M) * (0.28 M/1 M)^2 = 0.000686<br /><br />Using the Nernst equation, we can calculate E_cell:<br /><br />E_cell = -0.371V - (0.0592/2) * log(0.000686) = -0.371V + 0.0296 * 3.18 = -0.371V + 0.0947 = -0.2763V<br /><br />Since E_cell is negative, the reaction is not spontaneous. Therefore, statement I is incorrect.<br /><br />Statement II: E_cell = 0.809V.<br />We have already calculated E_cell to be -0.2763V, which is not equal to 0.809V. Therefore, statement II is incorrect.<br /><br />Statement III: Keq = 2.19 x 10^33.<br />The equilibrium constant (Keq) can be calculated using the standard cell potential and the number of moles of electrons transferred:<br /><br />Keq = exp(n * E^0_cell / 0.0592)<br /><br />Substituting the given values, we have:<br /><br />Keq = exp(2 * -0.371V / 0.0592) = exp(-12.52) = 4.79 x 10^-6<br /><br />This value is not equal to 2.19 x 10^33. Therefore, statement III is incorrect.<br /><br />Based on our analysis, none of the statements are correct. Therefore, the answer is none of the above.