A galvanic cell is set up from a zinc bar weighing 100 g and 1.0 litre of 1.0 M Cu $SO_4$ solution. How long would the cell run if it is assumed to deliver a steady current of 1.0 ampere ? (Atomic mass of Zn = 65)

$100 \,g$ of $Zn = \frac{100}{65}$ mol. Thus, according to reaction, $Zn + Cu^{2+} {->} Zn^{2+} + Cu, Cu^{2+}$ ions are limiting reagents. Now, charge flowing for reduction of $1$ mole of $Cu^{2+}$ ions $ = 2 \times 96500 \,C$ If $1$ ampere of current is to be delivered for $t$ hours, the quantity of electricity is $3600\,t\,C$ . Now $3600\,t = 2 \times 96500$ $ t = \frac{2\times 96500}{3600} = 53.61\,hr$

A galvanic cell is set up from a zinc bar weighing

Notes on Electrochemistry

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.