Electrolytic Cell vs Galvanic Cell & FAQs

Jul 9, 2022 | Turito Team

electronic cell

Electrolytic Cell

A cell is a device capable of producing electrical energy from chemical reactions or employing electrical energy to bring about a chemical reaction. So, cells can be grouped into two major categories: one that produces electrical energy from chemical reactions and another that uses electrical energy to bring about a chemical reaction. While the former is called a galvanic or voltaic cell, the latter is an electrolytic cell. 

Electrolytic Cell vs Galvanic Cell

Both electrolytic and galvanic cells operate differently. The following table enumerates the key differences between electrolytic cells vs galvanic cells.

Electrolytic Cells Galvanic Cells
It favours non-spontaneous reactions wherein electrical energy from an external source is converted to chemical energy.  It favours spontaneous reactions wherein chemical energy gets converted to electrical energy. 
Both electrodes lie submerged in the same electrolyte solution.  The half-cells are connected by a salt bridge.
The cathode is the negative electrode, and the anode is the positive electrode. Driven by an external power source, the electrons flow from anode to cathode. The cathode is the positive electrode, and the anode is the negative electrode. The electrons still flow from the anode to the cathode.

However, both cells contain two half-cells for a net-redox reaction with reduction and oxidation. In both the cells, oxidation occurs at the anode and reduction at the cathode.

What is an Electrolytic Cell?

An electrolytic cell is a device designed to utilise electrical energy and facilitate a non-spontaneous redox reaction. Thus, electrical energy is converted to chemical energy via the process of electrolysis. 

So, what is electrolysis? 

The process involving the passage of electric current from an external source into a solution of electrolyte is called electrolysis. 

An electrolytic cell is suitable for the electrolysis of certain compounds such as water when subjected to electrolysis forms gaseous hydrogen and oxygen.  

Components of an Electrolytic Cell 

The following electrolytic cell diagram shows the primary components of an electrolytic cell (see figure 1). 

An electrolytic cell      

An electrolytic cell consists of the following main parts: 

  • Tank: An electrolytic cell has an electrolytic tank made of a non-conducting material such as bakelite or glass. 
  • Electrolyte: The solution to be electrolysed (electrolyte) is filled in the tank. The electrolyte can be in the form of a solution or a fused state. 
  • Electrodes: The cell has two metallic/graphite rods dipped in the electrolyte and connected to a battery. The rod connected to the positive terminal acts as an anode, and that connected to the negative terminal acts as a cathode.  
Points to Remember

  • The anode has positive polarity, so anions move towards it.
  • The cathode has negative polarity, so cations move towards it.

Understanding Electrolytic Cell Function With Examples

Decomposition of sodium chloride is the simplest example of understanding how an electrolytic cell functions. If you melt sodium chloride to obtain a liquid and pass an electric current through the molten salt, it decomposes. 

The molten NaCl ionises to give solid sodium and chlorine gas. When electric current flows through the molten sodium chloride, the positive sodium ions move towards the negative cathode, and the negative chloride ions get attracted by the positive anode. Since chloride atoms cannot exist independently, they combine to form chlorine molecules. The half-reactions can be expressed as follows:

At the cathode: 

Na+ + e → Na 

At the anode:

Cl – e→ Cl (Primary change)

Cl + Cl → Cl2 (Secondary change)

The net reaction for the process can be written as shown below:

2NaCl (l) → 2Na (s) + Cl2 (g)

Faraday’s Laws of Electrolysis

Michael Faraday extensively studied electrolysis and established a relationship between the amount of substance deposited at the electrode and the quantity of electricity passed through the electrolyte. He formulated two quantitative laws that simplify the calculations based on electrolysis. 

Faraday’s 1st law of Electrolysis

According to the law, the weight of metal deposited at the cathode is directly proportional to the quantity of electricity used. 

W ∝ Q

W= ZQ 

Since Z = Electrochemical equivalent = Equivalent weight/ Faraday’s constant 

Faraday’s Constant= 96500

Q= it

The formula can be written as follows:

W = Zit 

W= (Eq wt × i × t)/ 96500

Faraday’s 2nd law of Electrolysis

According to the second law, if the same electricity (in terms of quantity) is passed through different electrolytes that stay connected in series, then the weight of metal deposited at the cathode is directly proportional to its equivalent weight. 

W ∝ Eq wt

w1 / w2= E1 / E2

Solved Example

Example 1: If a 0.2 Ampere current is passed through an aqueous solution of copper sulphate for an hour, what will be the copper mass accumulated at the cathode?

Solution: t= 1 hour = 60×60 seconds = 3600 s

Current = i = 0.2 A

The atomic mass of copper = 63.5

Valency of Copper in copper sulphate = 2

So equivalent mass of copper = atomic mass / valency 

Equivalent mass = 63.5/2 

Equivalent mass = 31.75

The electrochemical equivalent of copper can be calculated as follows:

Z = Eq. mass / 96500

Z = 31.75 / 96500

Z = 3.29 × 10-4g/C

According to Faraday’s first law of electrolysis, the mass of copper will be

W = Zit

W = 3.29 × 10 -4 × 0.2 × 3600

W = 0.2369 g

Answer: 0.2369 g mass of copper will be accumulated at the cathode.

Check your Knowledge:

Question 1: How long will you pass a current of 4 Ampere through a solution of silver nitrate to coat a metal surface of area 50 cm2 with a 0.005 mm thick layer? Given the density of silver is 10.5 g/cm3

Real-Life Applications of Electrolytic Cell

An electrolytic cell finds application in various useful procedures, such as coating certain metals to prevent their corrosion. Following are some useful applications of an electrolytic cell: 

1. Extraction and Refining of Metals

Refining metals involves impure metal extraction from its ore.

Anode: The ore serves as the anode.

Electrolyte: The salt of the metal becomes the electrolyte. 

Cathode: Pure metal deposits on the cathode.

Creating an electrolytic cell for metal extraction and refining is used for several metals such as copper, zinc, and more. As the pure metal deposits at the cathode, it is stripped from the cathode. The anode keeps shrinking in the process, and the impure anode is replaced to continue the process. 

 2. Chemical Production

Electrolysis produces chemicals like caustic soda, potassium permanganate, chlorine, ammonium per-sulphate, oxygen, hydrogen, and more on a large scale. 

Some processes like the electrolysis of brine give off several useful products. The major product collected at the cathode is caustic soda, while chlorine is liberated at one pole (at the anode) and hydrogen at the other (at the cathode). 

3. Electrolytic reduction of Metals from Compounds 

Aluminium is obtained from bauxite using an electrolytic cell. High-grade bauxite comprises up to 70% aluminium oxide, silica and iron oxide. Firstly, aluminium oxide is collected at the cathode and later, a suitable treatment helps obtain pure aluminium by reducing aluminium oxide. The process gives pure aluminium up to 99.5%.

4. Electroplating 

The process of covering the metal with another, usually for the following purposes, is called electroplating: 

  • To add more value to it
  • Decoration
  • Protection against corrosion 
  • Repair 
  • Intermediate manufacturing process

In electroplating, the object to be electroplated is made the cathode, and the metal to be deposited over it is made the anode. The electrolyte also contains the metal to be deposited. Often the cathode is surrounded by a set of anodes or rotated at uniform speed around the anode to give uniform coverage. Gold and silver plating are done following the same procedure.

Frequently Asked Questions

Q1. What are the criteria for product formation during electrolysis?

The two criteria form the base of electrolysis:

  • The substance with higher standard reduction potential is preferentially reduced at the cathode. 
  • The substance with lower standard reduction potential is preferentially oxidised at the anode.

Q2. What is sacrificial protection?

Sacrificial protection is the process of protecting a metal at the expense of some other more active or electropositive metal coated on its surface. It is often done via electrolysis; a common example is the galvanisation of iron. In this process, a thin layer of zinc is deposited on an iron substance to protect it from corrosion. 

Q3. Where are electrolytic cells used?

Electrolytic cells have a wide range of uses, such as electroplating metals, separating a pure form of a metal from its metallic compounds, recharging a battery, and more. The process of separating chemical compounds via electrolytic cells is known as electrolysis. Electroplating jewellery is one of the most common uses of electrolysis.


An electrolytic cell has a multitude of benefits linked to it. From extraction of pure metals to concealing an impure one with a pure one – an electrolytic cell makes it possible to carry out the desired processes. Faraday’s law of electrolysis helps in quantifying the amount of product that can be collected at the cathode of an electrolytic cell.