Electrochemical Cell - Definition, Examples, Types, Uses, FAQs

What is an Electrochemical Cell?

An energy cell is a machine that can generate electricity from the chemical reactions that occur in it, or use the electrical energy given to it to make chemical reactions in it. The given devices are capable of converting a chemical energy into an electrical energy, or even vice versa. A common example of an electrochemical cell is a standard 1.5-volt cell used to power many electronic devices such as TV remote and clocks. Such cells are capable of producing energy from the chemical reactions that occur in them and maintain the so-called Galvanic cells or Voltaic cells. In other words, the cells that cause the chemical reaction to take place when they pass through electrical energy are called electrical cells.

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Electrochemical Cell Example

A typical example of an electrochemical cell is a standard 1.5 volt cell intended for consumer use. A battery consists of one or more cells, usually connected in parallel, series or even series pattern and parallel. We interact with electrical cells in all aspects of our daily lives from disposable AA batteries in our remote controls and lithium-ion batteries on our iPhones to the nerve cells scattered throughout our bodies. There are basically two types of electrical cells namely galvanic cells, also called Voltaic cells and electrolytic cells. Important examples of electrolysis are the decomposition of water into hydrogen and oxygen, and bauxite into aluminum and other chemicals. Electroplating (e.g., copper, silver, nickel or chromium) is performed using an electrolytic cell. Electrolysis is a process that uses direct electric current (DC).

Cell Representation

Keep in mind that normal cell power can be calculated from the E0cell probability in both the connection and the reduction response. The positive energy of the cell indicates that the reaction continues spontaneously in the direction in which the response is recorded. On the other hand, a malignant cell reaction is found automatically on the reverse side. Cell recognition is a brief description of voltaic or galvanic (automatic) cells. Response conditions (pressure, temperature, concentration, etc.), anode, cathode, and electrode components are all described in this unique shortcut. Remember that oxidation occurs in the anode and reduction occurs in the cathode.

When the anode and cathode are connected by wire, electrons flow from the anode to the cathode. Normal galvanic cell A typical arrangement of half cells connected to form a galvanic cell. Using genetic engineering, let's assemble a cell. One beaker contains 0.15 M Cd (NO₃) ₂ and Cd metal electrodes. The other beaker even contains 0.20 M AgNO₃ and usually the Ag metal electrode. In response, the silver ion is reduced by receiving an electron, and the solid Ag is the cathode. Cadmium is connected by loss of electrons, and a strong Cd is anode.

Electrochemical Cell Diagram

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Half-Cells and Cell Strength

Electrochemical cells are made up of two and a half cells, each containing an electrode embedded in the electrolyte. The same electrolyte can also be used for both the half cells.

These half cells are connected by a salt bridge that provides an ionic communication platform between them without allowing them to combine. An example of a salt bridge is filter paper coated with potassium nitrate or sodium chloride solution.

One half part of an electrochemical cell loses electrons due to oxidation and the other gains electrons in the degradation process. It can be known that the equilibrium reaction occurs in both half cells, and when equilibrium is reached, the network voltage becomes 0 and the cell stops producing electricity.

The tendency of the electrode to contact the electrolyte for the loss or acquisition of electrons is explained by its electrode strength. The values of this energy can be used to predict the overall strength of a cell. Typically, the electrode strength is measured with the help of a standard hydrogen electrode as a reference electrode (known electrode).

Lower and Second Cells

The primary cells actually use and dislodge galvanic cells. The electrochemical reactions which take place in these cells are irreversible. Therefore, reactants are used for the production of electrical energy and the cell stops producing electricity when the reactants are completely depleted. Second cells (also known as rechargeable batteries) are electrical cells in which the cell has a flexible response, i.e. the cell can function as a Galvanic cell and an Electrolytic cell. Most primary batteries (many cells connected to a series, similarities, or combinations of the two) are considered to be harmful and environmentally damaging devices. This is because they require about 50 percent of the energy they contain in their production process. They also contain a lot of toxic metals and are considered hazardous waste.

Types of Electric Cells

There are two main types of electrochemical cells that are:

1. Galvanic cells (that are also called as Voltaic cells)

2. Electrical cells

The three main components of electrolytic cells are:

Cathode (negatively charged by electrolytic cells)

Anode (charged well with electrolytic cells)

Electrolyte

The electrolyte provides the electron exchange between the cathode and the anode. The most widely used electrolytes in electrolytic cells include water (containing dissolved ions) and molten sodium chloride.

Electrolytic Cell Function

Melted sodium chloride (NaCl) can be subjected to electrolysis. Here, two inert electrodes are placed in molten sodium chloride (containing Na + cations separated from Cl^- anions). When electrical energy is transferred to a region, the cathode becomes rich in electrons and has a negative charge. The well-charged sodium cation is now attracted to the badly charged cathode. This causes the formation of metallic sodium in the cathode. At the same time, chlorine atoms are attracted to a well-charged cathode. This causes the formation of chlorine gas (Cl₂) in the anode (corresponding to the release of two electrons, terminating the circuit). The corresponding chemical estimates and overall cell response are given below.

Cathode reaction: [Na⁺ + e^– → Na] x 2

Anode reaction: 2Cl^- → Cl₂ + 2e^–

Cell Response: 2NaCl → 2Na + Cl₂

Therefore, the molten sodium chloride can be subjected to electrolysis in an electrolytic cell to produce metallic sodium and chlorine gas as products.

Electrolytic Cell Applications

The primary use of electrolytic cells is to produce oxygen gas and hydrogen gas in water. They are used to extract aluminum from bauxite. Another notable use of electrolytic cells is electroplating, which is the process of forming a thin protective layer of a particular metal on the surface of another metal. Electrorefining of most non-ferrous metals is performed with the help of electrolytic cells. Such electrical cells are also used in electrowinning processes. It can be noted that the industrial production of very pure copper, pure zinc, and high-grade aluminum is regularly performed by electrolytic cells.

Simple Electric Cells

In Simple Electric Cells, the electricity is used to produce chemical changes. Similarly, we say that electrical energy is converted into chemical energy. Reversal occurs in Simple Cells (also known as Electric Cells) where chemical changes are used to generate electricity which means that chemical energy is converted into electrical energy.

Simple Electric Cell Parts:

A simple cell consists of two solid electrodes embedded in an electrolyte connected together by a cable-like conductor.

The two electrodes should be two separate metals.

The electrolyte can be an acid solution, an alkaline solution, a salt solution or a fruit like orange or lemon.

The electrical power generated depends on the positions of the two metals in the reactivity series.

"The farther the metal is in the regenerative series, the greater will be the electric current".

For example, a magnesium/ electrode pair will provide greater strength than a zinc/ copper electrode pair.

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How Do Simple Electric Cells Work?

Let's use a simple cell example to discuss further.

Metal electrode 1: Zinc

Steel electrode 2: Copper

Electrolyte: Reduce NaCl

The highest metal in the recycled steel chain will be selected and oxidized. It will easily remove electrons and a non-electrode (known as Anode). In the example above, zinc is more effective than copper, so the zinc electrode will be a non-electrode. Electrons leave the zinc electrode passing through the connecting wires to the copper electrode.

The lower metal at the bottom of the reconstituted metal chain will be a fine electrode (known as Cathode). Thus, the copper will be a fine electrode. The electrolyte contains good ions (sodium cations) and hydrogen ions. Hydrogen ions have a greater release rate than sodium ions and will be specially released (reduced) by receiving electrons from a non-wireless electrode. This is a redox reaction where oxidation and decomposition occur simultaneously. Here, electrons are transferred from zinc atoms to hydrogen ions.

Simple Cell Uses

Light cells are also commonly referred to as batteries that provide a portable form of electrical energy. They supply electricity to watches, lights, etc. that do not need to be connected to a large power supply.

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