How do you calculate electrochemical cell potential?
1 Answer
Answer
Answer:
Explanation
Explanation:
Answer:
Warning! VERY long answer! You can calculate the cell potential for an electrochemical cell from the halfreactions and the operating conditions.
Explanation:
The first step is to determine the cell potential at its standard state — concentrations of 1 mol/L and pressures of 1 atm at 25°C.
The procedure is:

Write the oxidation and reduction halfreactions for the cell.

Look up the reduction potential,
#E⁰_"red"# , for the reduction halfreaction in a table of reduction potentials 
Look up the reduction potential for the reverse of the oxidation halfreaction and reverse the sign to obtain the oxidation potential. For the oxidation halfreaction,
#E⁰_text(ox) = "" E⁰_text(red)# . 
Add the two halfcell potentials to get the overall standard cell potential.
#E⁰_text(cell) = E⁰_text(red) + E⁰_text(ox)#
At the standard state
Let’s use these steps to find the standard cell potential for an electrochemical cell with the following cell reaction.
#"Zn(s)" + "Cu"^"2+""(aq)" → "Zn"^"2+""(aq)" + "Cu(s)"#
1. Write the halfreactions for each process.
2. Look up the standard potential for the reduction halfreaction.
3. Look up the standard reduction potential for the reverse of the oxidation reaction and change the sign.
4. Add the cell potentials to get the overall standard cell potential.
Nonstandard state conditions
If the conditions are not standard state (concentrations not 1 mol/L, pressures not 1 atm, temperature not 25°C), we must take a few extra steps.
1. Determine the standard cell potential.
2. Determine the new cell potential resulting from the changed conditions.
a. Determine the reaction quotient,
#Q# .
b. Determine#n# , the number of moles electrons transferred in the reaction.
c. Use the Nernst equation to determine#E_"cell"# , the cell potential at the nonstandard state conditions.
The Nernst equation is
#color(blue)(bar(ul(color(white)(a/a)E_"cell" = E⁰_"cell"  (RT)/(nF)lnQcolor(white)(a/a))))" "#
where
Note: the units of
The moles refer to the “moles of reaction”.
Since we always have 1 mol of reaction, we can write the units of
Example
Calculate the cell potential for the following reaction when the pressure of the oxygen gas is 2.50 atm, the hydrogen ion concentration is 0.10 mol/L, and the bromide ion concentration is 0.25 mol/L.
1. Write the halfreactions for each process.
2. Look up the standard potential for the reduction halfreaction
3. Look up the standard reduction potential for the reverse of the oxidation reaction and change the sign.
4. Add the cell potentials together to get the overall standard cell potential.
5. Determine the new cell potential at the nonstandard conditions.
a. Calculate the value for the reaction quotient,
#Q# .
#Q = 1/(P_"O₂"["H"^"+"]^4["Br"^""]^4) = 1/(2.50 × 0.10^4 × 0.25^4) = 1.0 × 10^6# b. Calculate the number of moles of electrons transferred in the balanced equation.
#n = "4 mol electrons"# c. Substitute values into the Nernst equation and solve for
#E_"cell"# .
#E_"cell" = E_"cell" = E⁰_"cell"  (RT)/(nF)lnQ = "+0.152 V" – (8.314 "V"·color(red)(cancel(color(black)("C·K"^"1"))) × 298 color(red)(cancel(color(black)("K"))))/(4 color(red)(cancel(color(black)("mol"))) × "96 485" color(red)(cancel(color(black)("C·mol"^"1")))) × ln(1.0 × 10^6) = "0.152 V – 0.089 V" = "0.063 V"#
Describe your changes (optional) 200