Galvanic cell and EMF of the cell

Electrochemistry is a fascinating field that deals with the relationship between electrical energy and chemical reactions. One of the important components in the study of electrochemistry is the galvanic cell, also known as the voltaic cell. In this lesson, we will explore the concept of galvanic cells, their components, how they work, the concept of electromotive force (EMF) of a cell, and provide several examples to enhance your understanding.

Galvanic cells: An overview

A galvanic cell is a type of electrochemical cell that converts chemical energy into electrical energy through a spontaneous redox (reduction-oxidation) reaction. These cells are widely used in everyday applications, such as the batteries that power our devices.

A galvanic cell usually consists of two half-cells, each containing an electrode and an electrolyte. The half-cells are connected by a salt bridge or a porous membrane that allows the flow of ions between them. When the electrodes are connected via a wire, an electric current flows through the circuit as a result of the redox reaction.

Basic components of a galvanic cell

Let us understand the basic components of a galvanic cell:

1. Anode: The anode is the electrode where oxidation occurs. In a galvanic cell, the anode is negatively charged because it releases electrons during the oxidation process.
2. Cathode: The cathode is the electrode where reduction takes place. It is positively charged in a galvanic cell because it gains electrons during the reduction process.
3. Electrolyte: An electrolyte is a substance, often a solution, that contains ions and can conduct electricity. The electrolyte facilitates the movement of ions to maintain charge balance in the half-cells.
4. Salt bridge: Salt bridge is a device used to connect two half-cells and allow the flow of ions, thereby maintaining electrical neutrality in the system.
5. External circuit: The external circuit consists of wires and other components that connect the anode to the cathode, forming a complete electrical circuit.

How a galvanic cell works

To understand how a galvanic cell works, let's look at a common example: the zinc-copper cell.

In this cell:

• The anode is made of zinc metal (Zn).
• The cathode is made of copper metal (Cu).
• The zinc half-cell contains a solution of zinc sulphate (_ZnSO4).
• The copper half-cell contains a solution of copper sulphate (_CuSO4).
• The two half-cells are connected by a salt bridge or a porous membrane.

The reactions at each electrode can be represented as follows:

Anode (Oxidation): Zn (s) → Zn 2+ (aq) + 2e - Cathode (Reduction): Cu 2+ (aq) + 2e - → Cu (s)

In the zinc half-cell, zinc metal (Zn) is oxidized to zinc ions (Zn 2+), releasing electrons. These electrons travel through an external circuit to the copper cathode.

In the copper half-cell, copper ions (Cu 2+) gain electrons and are reduced to copper metal, which is deposited on the copper electrode.

These simultaneous oxidation-reduction reactions are what generate electricity. The flow of electrons through an external circuit from the anode to the cathode creates electric current. Meanwhile, the salt bridge helps balance the charge by allowing ions to flow between the two half-cells.

Visualization of a galvanic cell

Consider the following schematic representation of a galvanic cell, where zinc and copper are used as electrodes:

    ----- Wire -----
    Zn(s)Cu(s)
    ZnSO 4 (aq) CuSO 4 (aq)
    [salt bridge]

This setup summarizes the function and flow of a galvanic cell, emphasizing the movement of electrons from the zinc anode to the copper cathode and the role of the salt bridge in maintaining the equilibrium of the system.

Electromotive force (EMF) of the cell

The electromotive force (EMF) of a cell, also known as cell potential or cell voltage, is a measure of the energy provided by the cell to drive electrons from the anode to the cathode through an external circuit. The EMF is what powers electrical devices connected to the cell.

Calculating the EMF of a cell

The EMF of a galvanic cell can be calculated using the standard electrode potentials of the two half-cells involved in the reaction. The standard electrode potential (E 0) is the potential difference between the electrode and its solution at standard conditions (298K, 1M concentration, and 1 atm pressure).

The overall cell emf can be determined by subtracting the standard electrode potential of the anode from the potential of the cathode:

E cell 0 = E cathode 0 - E anode 0

For the zinc-copper cell example, if the standard electrode potentials are:

E Cu 2+/Cu 0 = +0.34 V E Zn 2+/Zn 0 = -0.76 V

Then, the EMF of the cell is calculated as:

E cell 0 = 0.34 V - (-0.76 V) = 1.10 V

This emf of 1.10 volts represents the ability of the cell to conduct electric current.

Factors affecting EMF

Many factors can affect the EMF of a galvanic cell, including:

• Concentration: Changing the concentration of electrolytes can affect the cell potential. According to Le Chatelier's principle, increasing the concentration of reactants generally increases the cell potential, while increasing the concentration of products decreases it.
• Temperature: While standard conditions assume a temperature of 25°C (298 K), any deviation from this can affect the EMF due to changes in reaction kinetics and thermodynamics.
• Pressure: For gases involved in an electrochemical reaction, changes in pressure can affect the cell potential. This is particularly relevant in cells with gaseous electrodes, such as hydrogen fuel cells.

Measuring EMF: Potentiometer

A potentiometer is a device used to accurately measure the EMF of a cell. Unlike a simple voltmeter, the potentiometer does not draw current from the cell being measured, resulting in a more accurate measurement.

This setup involves adjusting the known voltage source to balance the EMF of the cell until no current flows through the circuit. The voltage of the known source is then equal to the EMF of the cell being measured.

Example problem

Let's apply our understanding of galvanic cells and EMF with an example problem:

Consider a galvanic cell with the following half-reactions:

Anode: Mg (s) → Mg 2+ (aq) + 2e - (E 0 = -2.37 V) Cathode: Ag + (aq) + e - → Ag (s) (E 0 = +0.80 V)

Calculate the standard EMF of the cell.

Solution:

The EMF of the cell can be calculated using the following formula:

E cell 0 = E cathode 0 - E anode 0

Substitute the values:

E cell 0 = 0.80 V - (-2.37 V) = 3.17 V

Thus, the standard EMF of the cell is 3.17 volts.

Applications of galvanic cells

Galvanic cells are of great importance in a variety of applications, and serve as a vital component of the modern world:

• Batteries: Galvanic cells are the building blocks of batteries, which provide power to countless devices, from small electronics like smartphones to large electric vehicles.
• Fuel cells: While slightly different in design, fuel cells rely on galvanic cell principles to convert chemical energy from hydrogen or other fuels into electrical energy, providing a clean and efficient energy source.
• Corrosion prevention: Understanding the redox reactions in galvanic cells helps in designing methods to prevent corrosion, which is essentially a galvanic process.

Conclusion

Galvanic cells are a cornerstone of electrochemical understanding, demonstrating how chemical reactions can be used to generate electrical energy. Through concepts such as EMF, components such as anode and cathode, and the application of these cells in batteries and other technologies, galvanic cells remain an important area of study and innovation in chemistry. Mastering these topics paves the way for advances in sustainable energy, electronic devices, and industrial processes.

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