Question

What is the difference between a galvanic cell and an electrolysis cell? Give an example of each.

Short answer

The main difference between a galvanic cell and an electrolysis cell is that a galvanic cell converts chemical energy into electrical energy through spontaneous redox reactions, while an electrolysis cell uses electrical energy to drive non-spontaneous redox reactions, converting it into chemical energy. A common example of a galvanic cell is the Daniell cell, involving zinc and copper electrodes in their respective sulfate solutions. An example of an electrolysis cell is the electrolysis of water, using inert electrodes immersed in an aqueous electrolyte solution, resulting in the production of hydrogen and oxygen gases.

Step-by-step solution

Galvanic Cell Definition

A galvanic cell, also known as a voltaic cell, is an electrochemical cell that converts chemical energy into electrical energy through spontaneous redox reactions. In a galvanic cell, the anode (oxidation) and cathode (reduction) are kept separate; they are connected through an external circuit, allowing electrons to flow. The two half-cells are balanced by the flow of cations through a salt bridge, which maintains the electrical neutrality.

Electrolysis Cell Definition

An electrolysis cell is an electrochemical cell in which an electric current is used to drive a non-spontaneous redox reaction. In contrast to a galvanic cell, the electrical energy is converted into chemical energy. Electrolysis cells typically consist of two electrodes (anode and cathode) immersed in an electrolyte solution.

Example of a Galvanic Cell

A common example of a galvanic cell is the Daniell cell. In a Daniell cell, the anode is made of zinc and is placed in a zinc sulfate solution, while the cathode is made of copper and is placed in a copper(II) sulfate solution. The two half-cells are connected by a salt bridge and an external wire, enabling the flow of electrons from the zinc anode to the copper cathode. The overall redox reaction in a Daniell cell is: \[Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)\]

Example of an Electrolysis Cell

A common example of an electrolysis cell is the electrolysis of water. In this process, an electric current is applied to two inert electrodes (commonly made of platinum or graphite) immersed in an aqueous solution of an electrolyte (e.g., sulfuric acid, sodium chloride), which facilitates the conduction of ions. The overall redox reaction in the electrolysis of water is: \[2H_2O(l) \xrightarrow{electric \: current} 2H_2(g) + O_2(g)\] At the cathode, the reduction of water occurs: \(2H_2O(l) + 2e^- \rightarrow H_2(g) + 2OH^-(aq)\) At the anode, the oxidation of water takes place: \(2H_2O(l) \rightarrow O_2(g) + 4H^+(aq) + 4e^-\)

Key concepts

Galvanic Cell

A galvanic cell, also known as a voltaic cell, plays a crucial role in converting chemical energy into electrical energy. This conversion happens through spontaneous redox reactions. A galvanic cell is cleverly designed with two separate compartments, each housing an electrode. These are known as the anode and cathode.
The anode is where oxidation (loss of electrons) occurs, while reduction (gain of electrons) happens at the cathode. Though these compartments are kept separate, they are connected through an external circuit, allowing the flow of electrons from the anode to the cathode.

A key feature of galvanic cells is the salt bridge, which maintains electrical neutrality by allowing the flow of ions between the two compartments. Without the salt bridge, the reaction would come to a halt due to charge buildup. Galvanic cells are found in many common batteries, providing the power that fuels our electronic devices.

Electrolysis Cell

An electrolysis cell is like the reverse of a galvanic cell. Instead of producing electrical energy, it consumes it to drive non-spontaneous redox reactions. In an electrolysis cell, electrical energy is transformed into chemical energy. This is particularly useful for chemical processes that do not occur naturally or need to happen faster.

An electrolysis cell features two main components: electrodes and an electrolyte solution. The electrodes, typically made of inert materials like platinum or graphite, facilitate the transfer of electrons. The electrolyte solution helps conduct ions, making the transfer of electrical energy possible.

• At the anode, oxidation takes place.
• At the cathode, reduction occurs.

Electrolysis is widely used in industries, ranging from the production of aluminum to electroplating metals like gold.

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, are core processes in electrochemistry involving the transfer of electrons between substances. These reactions are fundamental to the working of both galvanic and electrolytic cells.
A redox reaction has two half-reactions:

• One involving oxidation, where a substance loses electrons.
• The other is reduction, where a substance gains electrons.

In electrochemical cells, the balance between these reactions is carefully maintained to ensure the desired energy transformation. Redox reactions are not limited to electrochemistry. They also occur in biological systems, combustion processes, and many other scenarios.

Daniell Cell

The Daniell cell is an iconic example of a galvanic cell, named after John Frederic Daniell, a British chemist. It elegantly demonstrates the concepts of redox reactions as it consists of two half-cells:

• A zinc anode in a zinc sulfate solution.
• A copper cathode in a copper(II) sulfate solution.

These half-cells are connected by a salt bridge and an external wire.

In a Daniell cell, zinc undergoes oxidation and loses electrons creating zinc ions, while copper ions undergo reduction, gaining electrons and depositing solid copper. The resulting reaction can be represented as:
\[Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)\]
The Daniell cell exemplifies how chemical reactions can effectively be used to generate electrical energy, a principle fundamental to battery technology.

Electrolysis of Water

Electrolysis of water is a classic application of an electrolysis cell. It involves splitting water molecules into hydrogen and oxygen gases using an electric current. The process requires inert electrodes, such as those made of platinum or graphite, and an electrolyte to conduct ions effectively, such as sulfuric acid or sodium chloride dissolved in water.

In the electrolysis of water, the following half-reactions occur:

• At the cathode: Water gains electrons to produce hydrogen gas and hydroxide ions: \(2H_2O(l) + 2e^- \rightarrow H_2(g) + 2OH^-(aq)\)
• At the anode: Water loses electrons, forming oxygen gas and protons: \(2H_2O(l) \rightarrow O_2(g) + 4H^+(aq) + 4e^-\)

The overall reaction converts electrical energy into chemical energy, breaking down water molecules and releasing them as gas. Electrolysis of water is a sustainable method for producing hydrogen fuel, with water being the only byproduct.