At 298 K, the standard electrode potentials of Cu2+/Cu, Zn2+/Zn, Fe2+/Fe and Ag+/Ag are 0.34 V, - 0.76 V, - 0.44 V and 0.80 V, respectively.
On the basis of standard electrode potential, predict which of the following reaction cannot occur?
On the basis of standard electrode potential, predict which of the following reaction cannot occur?
- CuSO4(aq) + Zn(s) → ZnSO4(aq) + Cu(s)
- CuSO4(aq) + Fe(s) → FeSO4(aq) + Cu(s)
- FeSO4(aq) + Zn(s) → ZnSO4(aq) + Fe(s)
2CuSO4(aq) + 2Ag(s) → 2Cu(s) + Ag2SO4(aq)
The Correct Option is D
Solution and Explanation
To predict which reaction cannot occur based on standard electrode potentials, we need to consider the concept of spontaneity in redox reactions. A reaction can spontaneously occur if the standard cell potential, \( E_{\text{cell}}^\circ \), is positive. For a redox reaction: \[ \text{Cell Potential}: E_{\text{cell}}^\circ = E_{\text{cathode}}^\circ - E_{\text{anode}}^\circ \]
The given standard electrode potentials are:
- \( \text{Cu}^{2+}/\text{Cu}: 0.34 \, \text{V} \)
- \( \text{Zn}^{2+}/\text{Zn}: -0.76 \, \text{V} \)
- \( \text{Fe}^{2+}/\text{Fe}: -0.44 \, \text{V} \)
- \( \text{Ag}^{+}/\text{Ag}: 0.80 \, \text{V} \)
We'll analyze the reactions:
\(\text{CuSO}_4(\text{aq}) + \text{Zn(s)} \to \text{ZnSO}_4(\text{aq}) + \text{Cu(s)}\)
- \( E_{\text{cell}}^\circ = 0.34 - (-0.76) = 1.10 \, \text{V} \, (\text{positive}) \) - Could occur.
\(\text{CuSO}_4(\text{aq}) + \text{Fe(s)} \to \text{FeSO}_4(\text{aq}) + \text{Cu(s)}\)
- \( E_{\text{cell}}^\circ = 0.34 - (-0.44) = 0.78 \, \text{V} \, (\text{positive}) \) - Could occur.
\(\text{FeSO}_4(\text{aq}) + \text{Zn(s)} \to \text{ZnSO}_4(\text{aq}) + \text{Fe(s)}\)
- \( E_{\text{cell}}^\circ = -0.44 - (-0.76) = 0.32 \, \text{V} \, (\text{positive}) \) - Could occur.
\(2\text{CuSO}_4(\text{aq}) + 2\text{Ag(s)} \to 2\text{Cu(s)} + \text{Ag}_2\text{SO}_4(\text{aq})\)
- \( E_{\text{cell}}^\circ = 0.34 - 0.80 = -0.46 \, \text{V} \, (\text{negative}) \) - Cannot occur.
The reaction \( 2\text{CuSO}_4(\text{aq}) + 2\text{Ag(s)} \to 2\text{Cu(s)} + \text{Ag}_2\text{SO}_4(\text{aq}) \) cannot occur spontaneously due to its negative cell potential.
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Concepts Used:
Electrochemical Cells
An electrochemical cell is a device that is used to create electrical energy through the chemical reactions which are involved in it. The electrical energy supplied to electrochemical cells is used to smooth the chemical reactions. In the electrochemical cell, the involved devices have the ability to convert the chemical energy to electrical energy or vice-versa.
Classification of Electrochemical Cell:
Cathode
- Denoted by a positive sign since electrons are consumed here
- A reduction reaction occurs in the cathode of an electrochemical cell
- Electrons move into the cathode
Anode
- Denoted by a negative sign since electrons are liberated here
- An oxidation reaction occurs here
- Electrons move out of the anode
Types of Electrochemical Cells:
Galvanic cells (also known as Voltaic cells)
- Chemical energy is transformed into electrical energy.
- The redox reactions are spontaneous in nature.
- The anode is negatively charged and the cathode is positively charged.
- The electrons originate from the species that undergo oxidation.
Electrolytic cells
- Electrical energy is transformed into chemical energy.
- The redox reactions are non-spontaneous.
- These cells are positively charged anode and negatively charged cathode.
- Electrons originate from an external source.