Galvanic cells and electrolytic cells are two types of electrochemical cells used to convert chemical energy into electrical energy or vice versa. Both types of cells rely on the redox reactions occurring at the electrodes to generate an electric current. There are some fundamental differences between these two types of cells that set them apart. In this answer, we will explore the differences between galvanic and electrolytic cells in detail.
Galvanic Cells:
Galvanic cells are also known as voltaic cells, named after the Italian physicist Alessandro Volta, who developed the first practical battery in 1800. Galvanic cells convert the chemical energy of a spontaneous redox reaction into electrical energy. In a galvanic cell, two different metals are immersed in solutions of their respective ions, which are called half-cells. These half-cells are separated by a salt bridge or a porous membrane that allows the exchange of ions between the two half-cells without mixing them.
The half-cell where oxidation occurs is called the anode, and the half-cell where reduction occurs is called the cathode. The anode is the negative electrode, and the cathode is the positive electrode. The anode and cathode are connected by a wire, which allows the flow of electrons from the anode to the cathode. The electrons flow through the wire from the anode to the cathode, creating a current. The salt bridge or porous membrane allows the flow of ions from the anode half-cell to the cathode half-cell to maintain neutrality.
The standard representation of a galvanic cell is:
Anode: M(s) → Mn+ + ne−
Cathode: M'(aq) + ne− → M'(s)
Overall cell reaction: M(s) + M'(aq) → M'(s) + Mn+ (aq)
In this reaction, M is the metal in the anode half-cell, M’ is the metal in the cathode half-cell, and n is the number of electrons transferred in the redox reaction. The standard cell potential, E°cell, is the difference between the standard electrode potentials of the two half-cells. A positive E°cell indicates that the reaction is spontaneous, and a negative E°cell indicates that the reaction is non-spontaneous.
Galvanic cells are used in batteries, fuel cells, and electrochemical sensors. The most common example of a galvanic cell is the zinc-copper cell, which is used in household batteries.
Electrolytic Cells:
Electrolytic cells are used to convert electrical energy into chemical energy by forcing a non-spontaneous redox reaction to occur. In an electrolytic cell, an external source of electrical energy, such as a battery, is used to force a redox reaction to occur. Electrolytic cells are used for electroplating, electrorefining, and the production of chemicals such as chlorine gas and sodium hydroxide.
In an electrolytic cell, there are two electrodes, an anode and a cathode, just like in a galvanic cell. However, the anode and cathode are connected to an external source of electrical energy, such as a battery or a power supply. The anode is the positive electrode, and the cathode is the negative electrode. The anode and cathode are immersed in an electrolyte, which is a solution of ions that can conduct electricity.
When a voltage is applied to the electrodes, the anode becomes the site of oxidation, and the cathode becomes the site of reduction. The electrons flow from the external source of electrical energy to the cathode, where they reduce the cations in the electrolyte. At the same time, the anode loses electrons and oxidizes the anions in the electrolyte. The oxidation and reduction reactions are non-spontaneous, and they require an external source of electrical energy to occur.
The standard representation of an electrolytic cell is:
Anode: M'(s) → M'(aq) + ne−
Cathode: Mn+(aq) + ne− → M(s)
Overall cell reaction: M'(s) + Mn+(aq) → M(s) + M'(aq)
In this reaction, M’ is the metal in the anode, and M is the metal in the cathode. n is the number of electrons transferred in the redox reaction. The voltage required to run the electrolytic cell is the opposite of the cell potential, Ecell, which is the difference between the standard electrode potentials of the two half-cells. A positive Ecell indicates that the reaction is non-spontaneous, and a negative Ecell indicates that the reaction is spontaneous.
Electrolytic cells are used for a variety of industrial applications, including electroplating of metals, production of chemicals such as chlorine gas and sodium hydroxide, and electrorefining of metals such as copper.
Differences between Galvanic and Electrolytic Cells:
Energy Conversion:
The most significant difference between galvanic and electrolytic cells is the direction of energy conversion. Galvanic cells convert chemical energy into electrical energy, while electrolytic cells convert electrical energy into chemical energy.
Electrode Potential:
The electrode potential of the two half-cells determines the direction of electron flow in a galvanic cell. The anode has a higher electrode potential than the cathode, and electrons flow from the anode to the cathode. In contrast, in an electrolytic cell, an external source of electrical energy is used to create a voltage that is applied to the electrodes. This voltage drives the oxidation of the anode and the reduction of the cathode.
Spontaneity:
Galvanic cells are spontaneous redox reactions, meaning they occur naturally and release energy. Electrolytic cells, on the other hand, are non-spontaneous redox reactions that require an external source of electrical energy to drive the reaction.
Electrolyte:
In a galvanic cell, the two half-cells are separated by a salt bridge or a porous membrane. This allows for the exchange of ions between the two half-cells without mixing them. In an electrolytic cell, the anode and cathode are immersed in an electrolyte solution that can conduct electricity.
Applications:
Galvanic cells are used in batteries, fuel cells, and electrochemical sensors. Electrolytic cells are used for electroplating, electrorefining, and the production of chemicals such as chlorine gas and sodium hydroxide.
Galvanic and electrolytic cells are two types of electrochemical cells that have fundamental differences in their energy conversion, electrode potential, spontaneity, electrolyte, and applications. Understanding these differences is essential for the design and implementation of various electrochemical technologies in different fields, including chemistry, biology, medicine, and materials science.
