Class 12 Chemistry Electrochemistry | Nernst Equation |

__Nernst Equation__

- This equation was named after a German physicist Walther Nernst.

- The Nernst Equation empowers the assurance of cell potential under non-standard conditions and relates the measured cell potential to the reaction quotient and permits the exact measurement of equilibrium constants.
- Let us consider an electrochemical reaction of the following type:

aA +bB --> cC + dD

- Nernst equation for this can be written as follows:

- In case of daniel cell Nerst equation is as follows:

- The above equation implis that the value of increases with the increase in the concentartion of Cu
^{2+}ion increases and decrease in the concentration Zn^{2+} - Putting the values of R, F at T= 298 K. the equation becomes

- If the circuit in Daniel cell is closed:

Zn(s) + Cu^{2+} (aq) →Zn^{2+} (aq) + Cu(s)

- With time the concentration of Zn
^{2+} - The concentration of Cu
^{2+} - Voltage reading of the cell on the voltmeter decreases.
- After some time there is no alteration in the concentration of Cu
^{2+}and Zn^{2+}ions and the voltmeter gives zero reading. At this point of time equilibrium has been reach - The Nernst equation for the reaction is:

- But at equilibrium,

At T = 298 K.

The equation can be rewritten as

**Problem:**

**Calculate the emf of the cell in which the following reaction takes place: Ni _{(s)} + 2Ag^{+} (0.002 M) → Ni^{2+} (0.160 M) + 2Ag_{(s)}. Given that E^{ø}_{cell} = 1.05 V.**

**Solution:**

By using Nernst equation

= 1.05 - 0.02955 log 4 × 10^{4}

= 1.05 - 0.02955 (log 10000 + log 4)

= 1.05 - 0.02955 (4 + 0.6021)

= 0.914 V

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