The correct statement(s) regarding the standard electrode potential, E° are

(A) The standard electrode potential changes with the change in the concentration of electrolytes
(B) The standard electrode potential is always zero at 273 K
(C) The concentration of each ion in the solution is unity.
(D) All the gases involved are at 1 bar

Choose the correct answer from the options given below:

(C) and (D) only

The correct answer is Option (2) → (C) and (D) only

The Standard Electrode Potential ($E^\circ$) is the potential of a half-cell reaction measured under standard conditions. These conditions are defined as:

(A) The standard electrode potential changes with the change in the concentration of electrolytes (Incorrect)

• $E^\circ$ is constant for a given reaction at a fixed temperature because it is a standard value.
• The electrode potential that changes with concentration is the Non-Standard Electrode Potential ($E$), which is calculated using the Nernst equation: $E = E^\circ - \frac{RT}{nF} \ln Q$.

(B) The standard electrode potential is always zero at $273 \text{ K}$ (Incorrect)

• $E^\circ$ is defined at a specified temperature, usually $298 \text{ K}$ ($\approx 25^\circ\text{C}$), but it is not always zero.
• The only potential that is defined as $0.00 \text{ V}$ by convention is the Standard Hydrogen Electrode ($\text{SHE}$) potential at all temperatures. All other standard electrode potentials are measured relative to the $\text{SHE}$.

(C) The concentration of each ion in the solution is unity. (Correct)

• Standard conditions require the concentration of all ions in the electrolyte solution to be $1 \text{ mol L}^{-1}$ (or unity).

(D) All the gases involved are at $1 \text{ bar}$ (Correct)

• Standard conditions require the pressure of all gases involved in the reaction (such as $\text{H}_2$ gas in $\text{SHE}$) to be $1 \text{ bar}$ (or $1 \text{ atm}$ in older conventions).