To what pH should you adjust a standard hydrogen electrode to get an electrode potential of -0.124V V ? (Assume that the partial pressure of hydrogen gas remains at 1 atm.)
2H^+ + 2e ==> H2
E = Eo -0.05916/n(log)pH2/(H+^2)
Plug in the numbers and solve for (H^+), then convert to pH.
pH-balanced enough to become a potential magnet! But seriously though, to get an electrode potential of -0.124V, you would need to adjust the pH of the standard hydrogen electrode to about 1.
To determine the pH at which you should adjust a standard hydrogen electrode to get an electrode potential of -0.124 V, you can use the Nernst equation. The Nernst equation relates the electrode potential of a half-cell to the concentrations (or activities) of the reactants and products involved.
The Nernst equation for the standard hydrogen electrode is:
E = E° - (0.0592/n) * log[H+]
Where:
E = electrode potential
E° = standard electrode potential (0 V for the standard hydrogen electrode)
[H+] = concentration of hydrogen ions (in this case, it represents pH)
n = number of electrons transferred in the half-reaction (1 for the standard hydrogen electrode)
Rearranging the equation to solve for [H+], we have:
log[H+] = (E° - E) * (n/0.0592)
Substituting the given values:
log[H+] = (0 - (-0.124)) * (1/0.0592)
log[H+] = (-0.124) * (1/0.0592)
log[H+] = -2.09
Now, to find the value of [H+], we need to take the antilog of -2.09:
[H+] = 10^-2.09
[H+] = 7.92 x 10^-3 M
Since we are working with pH, which is a logarithmic scale, we can convert the concentration [H+] into pH using the equation:
pH = -log[H+]
pH = -log(7.92 x 10^-3)
pH = 2.10
Therefore, you should adjust the standard hydrogen electrode to a pH of 2.10 to achieve an electrode potential of -0.124 V.
To determine the pH at which you should adjust a standard hydrogen electrode (SHE) to achieve a specific electrode potential, you need to use the Nernst equation. The Nernst equation relates the electrode potential to the concentration of the species involved in the redox reaction and the temperature. In the case of a standard hydrogen electrode, the redox reaction is:
2H+(aq) + 2e- → H2(g)
The Nernst equation for this reaction is:
E = E° - (RT/nF) * ln(Q)
Where:
- E is the electrode potential
- E° is the standard electrode potential (0 V for SHE)
- R is the gas constant (8.314 J/(mol·K))
- T is the temperature in Kelvin
- n is the number of electrons transferred in the redox reaction (2 in this case)
- F is Faraday's constant (96,485 C/mol)
- Q is the reaction quotient, which is the ratio of the concentrations of the products and reactants, raised to the power of their stoichiometric coefficients
In the case of the standard hydrogen electrode, the concentration of hydrogen gas (H2) is taken to be constant at 1 atm, which means the reaction quotient (Q) is equal to 1.
Rearranging the Nernst equation, we get:
E = E° - (RT/nF) * ln(1)
Since ln(1) = 0, the equation simplifies to:
E = E°
Therefore, to achieve an electrode potential of -0.124 V for a standard hydrogen electrode, the pH does not need to be adjusted. The standard hydrogen electrode is already at the desired potential for any pH value.