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In an acidic solution, copper ion is oxidized to copper(II) ion by the nitrate ion.

3 Cu (aq) + NO (aq)+4H (aq)NO(g) 3Cu2+ (aq) + 2H2O (0)

Part A Using the values provided, calculate the standard call potential, Ecell in (V) for this reaction And round your answer in 2 decimal.if your answer in negative, include the sign

Part B Use the rounded value of Ecell from Part A to calculate the standard free energy, delG, (in kJ) of the reaction at 298 K. Round the answer to nearest whole number

Part C: Use the rounded value of Ecell frm Part A to calculate the equilibrium constant, K at 298K.

In an acidic solution, copper ion is oxidized to copper(II) ion by the nitrate ion-example-1
User Joesan
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Final answer:

The question involves a redox reaction in acidic solution. Calculate Ecell by adding the half-reaction potentials, then use Ecell to find ΔG with the equation ΔG = -nFEcell, and K using the Nernst equation. Ecell is positive, indicating spontaneity. The standard cell potential, Ecell, for the given redox reaction is 0.34 V.

Step-by-step explanation:

In order to calculate the standard cell potential, Ecell, for the given redox reaction, we need to use the Nernst equation:



Ecell = E°cell - (0.0592/n) log(Q)



where E°cell is the standard cell potential, n is the number of electrons transferred in the balanced equation, and Q is the reaction quotient.



In this case, the balanced equation is:



3Cu(s) + 8H+(aq) + 2NO3-(aq) → 3Cu2+(aq) + 2NO(g) + 4H2O(l)



n = 6 (since 6 electrons are transferred)



The standard cell potential, E°cell, can be determined by looking up the reduction potentials for each half-reaction:



Cu2+(aq) + 2e- → Cu(s) (E° = 0.34 V)



NO3-(aq) + 4H+(aq) + 3e- → NO(g) + 2H2O(l) (E° = 0.96 V)



Substituting the values into the Nernst equation:



Ecell = 0.34 V - (0.0592/6) log(Q)



where Q is the reaction quotient:



Q = [Cu2+]^3 / ([H+]^8 [NO3-]^2)



Given the standard concentrations of Cu2+, H+, and NO3- are all 1 M, we can substitute these values into the equation:



Q = (1^3) / (1^8 * 1^2) = 1



Plugging this value into the Nernst equation:



Ecell = 0.34 V - (0.0592/6) log(1) = 0.34 V



Therefore, the standard cell potential, Ecell, for this reaction is 0.34 V.

User Mehroz Munir
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