Lesson Explainer: Primary Galvanic Cells Chemistry

In this explainer, we will learn how to describe primary cells and explain how they produce electrical energy.

A galvanic cell is a system that produces electrical energy from spontaneous chemical reactions. One example of a galvanic cell is a common household battery. The chemical reactions within the battery provide energy to power electronic devices. When all of the reactants have been used up, the battery can no longer provide energy. Sometimes, we distinguish between “primary” galvanic cells and “secondary” galvanic cells. A primary galvanic cell is used once and discarded, while a secondary galvanic cell can be recharged and reused.

To understand how primary galvanic cells generate electrical energy, we will look at the chemical reactions involved. Electronic devices rely on the coordinated movement of electrons. We call the flow of charge created by the moving electrons an electric current. Electronic devices are powered by an electric current that they use to do work.

The redox reactions that take place inside a primary galvanic cell create the flow of electrons that makes up an electric current. Redox reactions are chemical reactions where electrons are transferred from one species to another. To understand a galvanic cell, we need to understand how the electrons flow within it.

An oxidation half-reaction will take place at the anode, while a reduction half-reaction will take place at the cathode. Electrons will leave the anode, flow through the circuit, and arrive at the cathode. Along the way, the flow of charge can power devices such as toys or digital clocks.

Let’s take a more detailed look at a specific kind of primary galvanic cell, a mercury battery. These batteries are not commonly produced anymore due to the toxicity of mercury, but their chemistry gives a good example of how a primary galvanic cell operates.

A mercury cell is made up of a mercury oxide cathode and a zinc anode with a potassium hydroxide electrolyte. The cathode is connected to the base of the battery, while the anode is connected to the cap. Devices, like photography equipment, that use mercury batteries are designed so that electrons can flow from the anode, through the cap, through any circuitry in the device, and then back to the base and the cathode. The separator in the middle of the battery allows for the charge to be carried across it to complete the circuit but prevents any of the chemicals on opposite sides from mixing and reacting.

Let’s take a look at the individual chemical reactions in a mercury battery. At the anode, zinc combines with hydroxide ions to form zinc oxide, water, and two electrons:

At the cathode, mercury oxide reacts with water and two electrons to form mercury and hydroxide ions:

Therefore, the overall reaction for the cell will be

This battery has a limited life span because the reactions cannot easily be reversed. We begin with a supply of mercury oxide in the cathode and zinc metal in the anode. These chemicals gradually transform into liquid mercury and zinc oxide as power is drawn from the cell. Once we produce mercury and zinc oxide, it is not feasible to turn those products back into the initial reactants within the battery. Instead, once the reactants are used, the battery is discarded and a new one is used.

A fuel cell is a particular type of galvanic cell whose reactants can be refilled to allow continuous power generation. The most common type of fuel cell, a hydrogen fuel cell, is supplied with hydrogen gas and oxygen gas to be oxidized and reduced, respectively, to generate electrical energy. Fuel cells differ from other galvanic cells as they do not store energy. Their operation requires a continuous fuel supply and continuous products removal.

Fuel cells are of great interest in the field of transportation because they are more efficient than a car’s typical internal combustion engine. A combustion engine ultimately converts the chemical energy in gasoline to kinetic energy to drive the engine’s pistons; however, during combustion, some of the energy is lost in other forms such as heat or sound. Approximately, of the chemical energy in gasoline is successfully captured and powers the car.

In comparison, a fuel cell more efficiently converts the chemical energy with up to of the energy being successfully converted. The only direct waste products of a hydrogen fuel cell are water and heat, while combustion engines produce many harmful pollutants that may impact climate change and respiratory health.

Spacecrafts also use fuel cells supplied with hydrogen gas and oxygen gas. As the only direct waste products of a fuel cell are heat and water vapor which can be condensed and reused as drinkable water for astronauts.

Let’s take a closer look at the chemistry of a hydrogen fuel cell.

Galvanic cells rely on redox reactions and the flow of electrons they produce to power devices. In the case of a hydrogen fuel cell, hydrogen gas is oxidized and oxygen gas is reduced.

In a typical hydrogen fuel cell, both the anode and the cathode are present in the form of a hollow container that is lined with a porous carbon that allows connection between the internal room and the electrolyte that may be an aqueous solution of potassium hydroxide or an acidic solution of sulfuric acid.

In the case of an electrolyte solution of , at the anode, the hydrogen gas is fed into the system and loses electrons to form hydrogen ions that react with hydroxide ions in the electrolyte, forming water molecules:

At the cathode, oxygen gas most often from the air is reduced by the products of the previous reaction to form hydroxide ions. In some cases, a pressurized cylinder of pure oxygen is used instead:

These reactions combine to send electrons through the circuit, providing electrical energy. By combining the two half-equations, the overall reaction for this kind of hydrogen cell is

The only waste products are heat and water, and additional hydrogen can be added to provide continuous energy generation.

Example 3: Using the Direction of Current in a Hydrogen Fuel Cell to Identify the Electrode Where Oxidation Occurs

The following diagram is a representation of a hydrogen fuel cell.

Considering the direction of electrons through the top circuit, shown by the red arrow, what gas is introduced through pipe A?

1. Hydrogen
2. Oxygen
3. Steam

Answer

This question is asking us to identify which gas enters the left side of the hydrogen fuel cell. In the diagram, we can see that electrons travel from the left side of the cell to the right side.

The gas fed into the left side of this cell will lose electrons. Those electrons will travel along the circuit as indicated by the red arrow. On the right side, they will be accepted by the other gas.

Choice C is incorrect, as steam can be a product of this reaction but is not one of the initial reactants. The two gases fed into a hydrogen fuel cell are hydrogen gas and oxygen gas. The only thing left to determine is which gas is oxidized and which gas is reduced.

Hydrogen is the gas that is oxidized and loses electrons, forming ions, while oxygen is the gas that is reduced and accepts electrons, forming ions. Electrons will move from the side of the fuel cell with hydrogen gas toward the side of the fuel cell with oxygen gas.

The correct answer is choice A, hydrogen.

The voltage or electric potential difference found in a battery is also called the “cell potential.” We can calculate the cell potential by finding the difference between the two individual electrode potentials, also called reduction potentials.

Electrons will be given up at any anode and accepted at any cathode, but not all anodes and cathodes will do so with the same intensity. An oxidation reaction that happens quite easily at the anode will have a very negative reduction potential, meaning that electrons will readily separate. On the other hand, a reduction reaction that happens quite easily at the cathode will have a very positive reduction potential.

All things being equal, a large difference in reduction potentials means a large voltage with many high-energy electrons leaving the battery at any given moment. A small difference in reduction potential means a small voltage and weaker electric current. Think of how a ball rolls from a high potential energy state at the top of a hill to the low potential energy state at the bottom; electrons will move from their unstable, high-energy position at the anode to the more stable, lower-energy position at the cathode. It is worth noting that the conductivity of the wire also affects the strength of the current, but a discussion of conductivity is beyond the scope of this explainer.

In the case of an electrolyte solution of sulfuric acid, the following equation shows how the reduction potential of the cathode in a hydrogen fuel cell under acidic conditions is V:

Meanwhile, the oxidation of hydrogen has a reduction potential of 0 V:

This reduction potential is 0 V, but that does not mean there is no electrical force acting on the electrons. The oxidation of hydrogen has been arbitrarily selected to be the reference point of zero on the scale of relative electrode potentials.

The overall voltage for the battery, also known as the standard cell potential, can be calculated as a difference between the two reduction potentials:

Electrons will flow from the electrode with a lower reduction potential (the anode) to the electrode with a higher reduction potential (the cathode). The greater the difference in reduction potential, the higher the cell potential, and the more energy the electrons will carry along the circuit.

Let’s summarize what has been learned in this explainer.