Examples of Galvanic Cells in Everyday Life

Electricity is one of the most revolutionary discoveries in the history of mankind. While today we have several methods to harvest electrical energy on a tremendous scale, galvanic cell, one of the earliest methods to generate electricity, is still of great value to our technology. A galvanic cell or voltaic cell is an electrochemical cell named after two scientific pioneers of the 18th century, Luigi Galvani, a professor of anatomy in Bologna, Italy, and Alessandro Volta, an Italian physicist, chemist, and pioneer of electricity and power. 

Definition

A galvanic cell is an electrochemical device that generates electric energy by chemical reactions between dissimilar conductors coupled by an electrolyte and a salt bridge. The spontaneous oxidation-reduction reactions among the components power a galvanic cell. In a redox reaction, an electron transfer generates electrical energy, which is channeled through a galvanic cell.

Discovery

In 1780, while demonstrating the nervous system using the lower half of a frog, Luigi Galvani found that the muscles in the frog’s legs contracted when a copper wire attached to the nerve and a zinc wire attached to the muscle made contact with each other. He presumed that the electricity was coming from the animal tissue of the frog and named it “animal electricity.” Later, Alessandro Volta established that the two metals created an electrical current with electrons exchanging among them; thus, the metals alone were responsible for the twitches. Volta published his theory of bi-metal electricity in 1799, along with his magnificent innovation based on it: the Voltaic pile, which has now evolved into the battery. For this reason, batteries are often designated as galvanic cells as well as voltaic cells. Even though Volta had no understanding of how a battery or a galvanic cell worked, these discoveries are considered milestones in the history of scientific developments. Forty years later, Faraday concluded that the flow of electrons results from a redox reaction caused among two metals. Let’s try to understand the mechanism of galvanic cells in a detailed manner. 

Mechanism

The basic principle of the voltaic cell is a simultaneous oxidation and reduction reaction called a redox reaction. The chemical energy released in a redox reaction, during the transfer of electrons from a higher to lower potential, is harnessed as electric energy using electrodes. To understand the complex interplay of chemicals in a galvanic cell, let’s take an example of a specific galvanic cell called the “Daniel cell.” A Daniell cell is a type of electrochemical cell that is composed of copper and zinc electrodes. When a strip of zinc metal (Zn) is immersed in an aqueous solution of copper sulfate ({CuSO}_{4}), dark-colored solid deposits form on the zinc metal’s surface, and the blue color of the {Cu}^{2+} ion vanishes. Copper metal depositions on the zinc metal’s surface have resulted in zinc ions being added to the solution. This reaction is chemically represented by:

In the above reaction, zinc would lose two electrons as a result of the chemical process and copper will absorb those electrons to become elemental copper. Because these two metals will be placed in separate containers and connected by a conducting wire, an electric current will form, transferring all electrons from one metal to the other. To simplify balancing the overall equation easier and to highlight the actual chemical transformations, it’s common to split the oxidation-reduction processes into half-reactions. The half-reactions of oxidation and reduction are separated and connected by a wire, which forces electrons to flow through it. A negative charge rises in the reduction half-cell and a positive charge builds in the oxidation half-cell when electrons are moved from the oxidation half-cell to the reduction half-cell. If the cell didn’t have a way for ions to flow between the two solutions, the charge accumulation would act to oppose the current from anode to cathode, effectively preventing electron flow.

Batteries

A battery is a power source consisting of one or more electrochemical cells with external connections for powering electrical devices such as flashlights, mobile phones, and electric cars. There are two types of batteries: disposable (primary) batteries, which have irreversible electrode reactions and cannot be recharged, and rechargeable (secondary) batteries, which produce an insoluble substance that sticks to the electrodes. By providing an electrical potential in the opposite direction, these batteries can be recharged. A rechargeable battery is briefly converted from a galvanic cell to an electrolytic cell during the recharging process. Batteries are intelligently designed devices that follow the same basic principles as galvanic cells. The main difference between commercial batteries and the galvanic cells we discussed earlier is that commercial batteries use solids or pastes as reactants rather than liquids to maximize electrical output per unit mass. Another advantage of using highly concentrated or solid reactants is that the concentrations of the reactants and products do not change significantly when the battery is discharged, resulting in a relatively constant output voltage throughout the discharge process. The output of the Zn/Cu cell, on the other hand, drops logarithmically as the reaction progresses.

Leclanché Dry Cell

The dry cell is by far the most common type of battery used in several electronic devices, such as TV remotes, watches,  and many other devices. Although the dry cell was patented in 1866 by the French chemist Georges Leclanché and more than 5 billion such cells are sold every year, the details of its electrode chemistry are still not completely understood. Despite its name, the Leclanché dry cell is actually a “wet cell.” The chemical process which produces electricity in a Leclanché cell begins when zinc atoms on the surface of the anode oxidize, i.e. they give up both their valence electrons to become positively charged ions. As the Zn ions move away from the anode, leaving their electrons on its surface, the anode becomes more negatively charged than the cathode. When the cell is connected to an external electrical circuit, the excess electrons on the zinc anode flow through the circuit to the carbon rod, the movement of electrons forming an electric current. As the current travels around the circuit, when the electrons enter the cathode (carbon rod), they combine with manganese dioxide (MnO2) and water (H2O), which react with each other to produce manganese oxide and negatively charged hydroxide ions. This is accompanied by a secondary acid-base reaction in which the hydroxide ions (OH–) accept a proton (H+) from the ammonium ions present in the ammonium chloride electrolyte to produce molecules of ammonia and water

Lithium–Iodine Battery

Lithium batteries feature a metallic lithium anode and are used as primary batteries. Lithium-metal batteries are another name for these types of batteries. Their high charge density and expensive cost per unit set them apart from other batteries. Lithium cells can produce voltages ranging from 1.5 V (similar to a zinc-carbon or alkaline battery) to around 3.7 V, depending on the design and chemical compounds utilized. They all have minor amounts of liquid water in them, which increases the mass and poses a risk of corrosion. As a result, a lot of effort has gone into developing water-free batteries. The lithium–iodine battery is one of the few commercially effective water-free batteries. The cathode is a solid complex of I2 and the anode is lithium metal. A layer of solid LiI separates them, acting as an electrolyte by allowing Li+ ions to diffuse.

Nickel–Cadmium (Ni-Cd) Battery

Several electrical appliances and gadgets such as drills, portable vacuum cleaners, and AM/FM digital tuners employ the nickel-cadmium, or NiCad, battery. It’s a water-based cell with a cadmium anode and a highly oxidized nickel cathode known as nickel(III) oxo-hydroxide, or NiO. (OH). The design enhances the electrodes’ surface area while minimizing the distance between them, lowering internal resistance and allowing for a relatively high discharge current. A metal box with a sealing plate and a self-sealing safety valve are typical of Ni-Cd batteries. This is known as the jelly-roll design and allows a Ni-Cd cell to deliver a much higher maximum current than an equivalent size alkaline cell. Alkaline cells have a bobbin construction where the cell casing is filled with electrolyte and contains a graphite rod which acts as the positive electrode. As a relatively small area of the electrode is in contact with the electrolyte (as opposed to the jelly-roll design), the internal resistance for an equivalent sized alkaline cell is higher which limits the maximum current that can be delivered.

Lead–Acid (Lead Storage) Battery

A lead-acid battery is a type of battery that converts chemical energy into electrical energy using sponge lead and lead peroxide. Because of its higher cell voltage and lower cost, lead-acid batteries are most typically employed in power plants and substations. When sulfuric acid dissolves, its molecules split into two types of ions: positive hydrogen ions (2H+) and sulfate negative ions ({SO}_{4}^{-}) that can freely move around. If the two electrodes are immersed in solutions and linked to a DC supply, positively charged hydrogen ions will migrate towards the electrodes and attach to the supply’s negative terminal. Negatively charged {SO}_{4}^{-} ions gravitated toward the electrodes attached to the supply main’s positive terminal (i.e., anode). Each hydrogen ion takes one electron from the cathode, and each sulfate ion absorbs two negative ions from the anodes, forming sulfuric and hydrogen acid when they combine with water. The oxygen created by the equation above reacts with the lead oxide to make lead peroxide ({PbO}_{2}) As a result, while the lead cathode remains lead during charging, the lead anode is transformed into lead peroxide, which has a chocolate color.

Fuel Cells

A fuel cell is an electrochemical cell that uses an electrochemical reaction to create electrical energy from fuel. To keep the processes that generate electricity going, these cells need a constant supply of fuel and an oxidizing agent (usually oxygen). As a result, until the supply of fuel and oxygen is shut off, these cells can continue to generate power. Despite being conceived in 1838, fuel cells did not enter commercial usage until a century later, when NASA utilized them to power space capsules and satellites. Many facilities, including industrial and commercial buildings, now use these devices as a primary or supplementary source of electricity. The membrane electrode assembly (MEA) in fuel cells is made up of gas diffusion layers, electrodes, and a polymer electrolyte membrane. Inside the MEA, electrochemical reactions generate power. Hydrogen flows to the anode side of a single PEM fuel cell, where it is divided into protons and electrons through reactions in the presence of a catalyst. The electrons are conducted through a network of carbon nanoparticles in the electrode, which provides current output to power a device before they reach the cathode on the opposite side. Meanwhile, protons move through a proton exchange membrane to reach the cathode, while oxygen from the air diffuses through a gas diffusion layer (GDL) in the MEA. All of these processes take place within a cell stack. Fuel, water, and air management, coolant control, hardware, and software are all part of a broader system that comprises cell stacks. The systems range in size and function depending on their intended usage, which ranges from transportation to industrial gear to backup power to supplement the electric grid.