Electrochemical Cell

Definition of Electrochemical Cell

Electrochemical cells are the main sites for electrochemical reactions to take place. So what are they exactly?

The main concept of electrochemical cells is that they allow a redox reaction to be split, in order to control it. In redox reactions, the processes of oxidation and reduction occur simultaneously. This means that splitting these two processes apart can be used for specific chemical needs. By splitting the reaction, you can control the flow of electrons between the oxidation and reduction half-reactions.

Electrochemical cells are complete circuits, which means that even though they take apart the given redox reaction, they are connected at two points: the external circuit and the salt bridge. The external circuit will allow for the transfer of electrons, while the salt bridge will facilitate the mobility of the anions of the reaction. This way the reaction forms a circuit and is able to proceed.

Below we will talk about the definition of each electrochemical cell.

Types of Electrochemical Cells

Here we will discuss the two types of general electrochemical cells: voltaic and electrolytic. More importantly, we will focus on how they differ from each other, and how they were developed for different purposes and chemical needs.

Voltaic (Galvanic) Cell

A voltaic cell is an electrochemical cell which facilitates a spontaneous reaction.

These types of reactions will cause energy to be released. By creating a voltaic cell, you can monitor the energy through a voltmeter, as electrons will flow in the external circuit creating electricity.

Voltaic cells have two half-cells, each of which facilitate a single half-reaction, be it oxidation or reduction. You will see a diagram of this in the next section. The two half cells are connected by a salt bridge to promote the mobility of anions through the solutions, completing the circuit. This is opposed to how electrolytic cells are comprised.

Electrolytic Cell

Electrolytic cells usually don't have half-cells, but rather are performed in a single reaction compartment, such as a single beaker. This is because the reaction we are trying to perform is non-spontaneous, so mixing the reagents will not result in a reaction.

In electrolysis, the reaction you are trying to perform is non-spontaneous, meaning that you have to put energy into the system to make the reaction proceed. Applying a current to the reaction through electrodes will cause oxidation at one electrode and reduction at the other electrode. A current, being a stream of electrons, will supply energy and shift the equilibrium of the reaction to one side.

Electrochemical Cell Diagram

Here we will go through and represent a diagram of an electrochemical cell, specifically a voltaic (galvanic) cell.

A voltaic electrochemical cell diagram depicting a Copper-Zinc cell. Source: glossary.periodni.com

Above you will see the diagram of an electrochemical cell depicting the reaction of Zinc with Copper Sulphate. Here Zinc is oxidised, since it loses electrons, while Copper is reduced, since it gains electrons.

Below the diagram, you will notice the half-reactions for each half-cell. These will tell you what reaction is occurring at each electrode. The electrodes are labelled above to state which one is the anode and which one if the cathode. Additionally, below the diagram you will find the cell notation for the reaction.

In the next section, we will talk about the different features of electrochemical cells.

Features of Electrochemical Cells

Electrochemical cells are composed of many different components, as seen in the diagram above. Here, we will discuss the main components of electrochemical cells. These are often the same in voltaic and electrolytic cells.

Electrodes

There will be two electrodes in each electrochemical reaction. Each electrode will function to carry out a specific half-reaction. These will be either oxidation or reduction.

The electrodes can be either made from a metal that will participate in the reaction, or they can be inert. Inert electrodes do not participate chemically in the reaction, as they will be made of either graphite or platinum.

Connectors: External Circuit and Salt Bridge

In an electrochemical cell, be it voltaic or electrolytic, there will be a need for an external circuit where the current can be carried. The flow of electrons generated in the reaction will be transferred through this circuit.

On the other hand, the flow of anions will be facilitated by the salt bridge. This is usually an external piece of equipment that is porous and can support the flow of anions but not the cations. Additionally, this can be either an external tube or a thin membrane separating different parts of the same beaker.

Additional Accessories

Other components which are crucial for electrochemical cells include the storage compartment of the reaction, which can be either a single beaker, as seen in electrolysis reactions, or two beakers, which act as half cells for the redox reaction.

Additionally, another component which deals with the flow of electrons is necessary. This regulatory equipment for electricity can be either in the form of a voltage meter, for voltaic cells, or a battery (or other power source) for electrolytic cells. These components directly monitor and modulate the flow of electrons in the electrochemical cell.

Applications of Electrochemical Cells

Here we will discuss the main applications of electrochemical cells, which concern calculating the standard cell potential and other thermodynamic values of electrochemical reactions.

Standard Cell Potential

So what is cell potential?

So there are two ways which can be used to find the standard cell potential for a reaction.

1. Perform the reaction and read the voltage meter.
2. Calculate the standard cell potential from standard electrode potentials.

Here, we will focus on how to calculate the standard cell potential by combining two standard electrode potentials.

You can find information on different standard electrode potentials in a table such as this one (you can probably find one at the end of your chemistry textbook):

Table of standard electrode potentials. Source: inspiritvr.com

Above you can see all the different voltage values for different electrode potentials.

To calculate the standard electrode potential of an electrochemical cell, use the following formula:

E⁰_cell = E⁰_red - E⁰_oxid

Other Thermodynamic Factors

The standard cell potential is related to other thermodynamic aspects, such as Gibbs free energy, the equilibrium constant, and entropy.

Gibbs Free Energy

Take a look at the following equation:

ΔG = −nFE_cell

This equation joins the parameters of Gibbs free energy with the cell potential. Through this equation, you can calculate Gibbs free energy just by knowing the cell potential of a reaction, or the other way round.

This is because the other components in the equation are constants: "n" is the number of electrons involved in the reaction, while "F" is Faraday's Constant (= 96,485 C/mol).

Equilibrium Constant

We can use the equilibrium constant to monitor the redox reaction.

Take a look at the following equation:

ΔG⁰ = −RTlnK

Thus:

−nFE⁰_cell = −RTlnK

Hence:

E⁰_cell = (RTlnK) / (nF)

Through this equation, you can find the equilibrium constant by knowing the standard cell potential or the other way round. This is because the other values in the formula are constants: "T" being the temperature in K, and "R" being the gas constant (= 8.314 joule kelvin⁻¹ mole⁻¹).

Predicting Redox Direction

Predicting the redox direction, or in other words the direction of electron flow, depends on the electrodes and the reactions they facilitate.

So, we know that in a redox reaction there will be oxidation and reduction occurring simultaneously. The electrodes split those two processes apart.

The reduction will occur at the cathode, while oxidation will occur at the anode. One you have figured out which electrode is the cathode and which one is the anode, predicting the direction of electron flow should be easy.

Electrons will flow from the anode to the cathode.