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Electrode Potential

Zinc is coate...

Question

Zinc is coated over iron to prevent rusting of iron because:

A

it is cheaper than iron

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B

E_{(Z n^{2 +} / Z n)}^{\circ} = E_{(F e^{2 +} / F e)}^{\circ}

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C

E_{(Z n^{2 +} / Z n)}^{\circ} < E_{(F e^{2 +} / F e)}^{\circ}

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D

E_{(Z n^{2 +} / Z n)}^{\circ} > E_{(F e^{2 +} / F e)}^{\circ}

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Solution

The correct option is D $E_{(Z n^{2 +} / Z n)}^{\circ} < E_{(F e^{2 +} / F e)}^{\circ}$
Zinc is used for coating iron to prevent rusting because whenever the metal comes in contact with moisture and water then there is formation of electrochemical cell and oxidation of metal takes place, zinc is having lower reduction potential and acts as anode and oxidation of zinc take place rather than iron so, in this way we can prevent iron from rusting.

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Similar questions

Q. The standard oxidation potentials,

E^{\circ}

, for the half reactions are as,

Z n \to Z n^{2 +} + 2 e^{-}; E^{\circ} = + 0.76 v o l t

F e \to F e^{2 +} + 2 e^{-}; E^{\circ} = + 0.41 v o l t

The emf of the cell,

F e^{2 +} + Z n \to Z n^{2 +} + F e

Q. The standard reduction potentials

E^{⊝}

for the half reactions are follows :

Z n \to Z n^{2 +} + 2 e^{-}; E^{⊝} = + 0.76 V

F e \to F e^{2 +} + 2 e^{-}; E^{⊝} = 0.41 V

The EMF for the cell reaction

F e^{2 +} Z n \to Z n^{2 +} + F e

Q. The standard electrode potential of the half cells is given below:

Z n^{2 +} + 2 e^{-} \to Z n; E = - 0.76 V

F e^{2 +} + 2 e^{-} \to F e; E = - 0.44 V

The emf of the cell

F e^{2 +} = Z n^{-} \to Z n^{2 +} + F e

The standard reduction potentials, $E^{\circ}$ , for the half-reactions are as $Z n = Z n^{2 +} + 2 e$ , $E^{\circ}$ = + 0.76 V and $F e = F e^{2 +} + 2 e$ , $E^{\circ}$ = + 0.41 V. The e.m.f. for the cell reaction, $F e^{2 +} + Z n = Z n^{2 +} + F e$ is

$E^{\circ}$ value of $M g^{2 +} | M g, Z n^{2 +} | Z n$ and $F e^{2 +} | F e$ are – 2.37 V, - 0.76 V and 0.44 V respectively. Which of the following is correct?

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