Electrochemistry: Meaning, Important Terms, Electrolysis and Redox Reaction

Electrochemistry is the branch of chemistry that studies the relationship between electricity and chemical reactions. It involves the movement of electrons between substances, which can be harnessed to drive chemical reactions or generate electrical energy. Electrochemistry includes various key concepts such as Electrochemical Cells, Redox Reactions, Galvanic Cells, Electrolytic Cells, Nernst Equation,

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This Story also Contains

1. Electrochemistry
2. Electrolysis
3. Some Solved Examples
4. Summary

Electrochemistry emerged through several key scientists' work over time, but the foundation can be traced back to the late 18th and early 19th centuries. The Italian scientist Alessandro Volta is credited with discovering the first chemical battery, known as the "Voltaic Pile," in 1800. This device could produce a steady electric current from a chemical reaction, demonstrating the direct relationship between chemical reactions and electricity. This discovery was pivotal in developing electrochemistry because it provided the first evidence that chemical processes could produce electrical energy. The British chemist Humphry Davy further advanced electrochemistry by using the voltaic pile to perform electrolysis. In 1807, he isolated several chemical elements, including sodium and potassium, through electrolysis. His work provided key insights into the nature of chemical reactions and the role of electricity in driving them. Michael Faraday's contributions in the 1830s were fundamental to the development of electrochemistry. Faraday established the laws of electrolysis, which quantify the relationship between the amount of substance transformed in an electrochemical reaction and the quantity of electricity passed through the system. His laws and principles laid the groundwork for understanding how electrical energy can be used to drive chemical reactions. Together, these scientists developed the key principles of electrochemistry, demonstrating how chemical reactions can produce electrical energy and how electricity can drive chemical changes. Their work provided the basis for many modern applications in chemistry and engineering.

Electrochemistry

Electrochemistry is the branch of science that deals with transforming chemical energy into electrical energy and vice versa or the relationship between electrical and chemical energy produced in a redox reaction.

Galvanic Cell (or Voltanic Cell)
Consider the following redox reaction:

$\mathrm{Zn}(\mathrm{s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \longrightarrow \mathrm{Cu}(\mathrm{s})+\mathrm{Zn}^{2+}(\mathrm{aq})$

Zn displaces copper ions (Cu2+) from aqueous solution in the above reaction. This reaction can be achieved very easily in practice. Put a Zn rod into a solution of CuSO₄ (containing Cu²⁺ ions). It is observed that the blue color of the CuSO₄ solution disappears after some time. In this situation, Zn loses 2 electrons per atom, and Cu²⁺ ions in the solution accept them. In this manner, cu2+ ions from the solution are deposited in the form of solid Cu and Zn goes into the solution as Zn²⁺ (colorless). The reaction can well be understood in terms of two half-reactions:

Oxidation : $\quad \mathrm{Zn}(\mathrm{s}) \longrightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-}$
Reduction: $\mathrm{Cu}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{Cu}(\mathrm{s})$

We can make the same reaction occur even if the copper ions and zinc rod are not in direct contact. If we put the Cu²⁺ ions and Zn rod in two separate containers connect the two by a conducting metallic wire and introduce an inverted U-shaped instrument (called a salt-bridge). Electrons will still be transferred through the connecting wires. The electrons from the Zn rod travel to Cu²⁺ ions through the connecting wires and the same reaction occurs. This flow of electrons through the wire generates electricity.

Electrolysis

It is a process by which an electric current is passed through a substance to effect a chemical change. A chemical change is when the substance loses or gains an electron (oxidation or reduction). The process is carried out in an electrolytic cell, an apparatus of positive and negative electrodes held apart and dipped into a solution containing positively and negatively charged ions. The substance to be transformed may form the electrode, constitute the solution, or be dissolved in the solution. Electric current enters through the negatively charged electrode (cathode); positively charged components of the solution travel to this electrode, combine with the electrons, and are transformed into neutral elements or molecules. The negatively charged components of the solution travel to the other electrode (anode), give up their electrons and are transformed into neutral elements or molecules. If the substance to be converted is the electrode, the reaction is generally in which the electrode dissolves by giving up electrons.

Recommended topic video on(Electrochemistry)

Some Solved Examples

Example.1

1. Consider the reaction

$2 \mathrm{Ag}^{+}+\mathrm{Cd} \rightarrow 2 \mathrm{Ag}+\mathrm{Cd}^{2+}$

The standard potential for

$\mathrm{Ag}^{+} \rightarrow \mathrm{Ag}$ and $\mathrm{Cd}^{2+} \rightarrow \mathrm{Cd}$ couples are 0.80 V and -0.40 V respectively Now,

a) Standard potential $E^0$ for the cell is 1.20V

b) Cd electrode is the negative electrode

Which of the following statements is correct?

1)a

2)b

3) (correct)a,b

4)none

Solution

$E^0=E_{\text {cathode }}^0-E_{\text {anod }}^0$

$=0.80-(-0.04)=1.20 \mathrm{~V}$

The negative electrode is always the electrode that has a lower value of reduction potential

$\therefore C d$ is negative electrode

Hence, the answer is the option (3).

Example.2

2. Choose the correct option :

a) During recharging Gibbs's free energy is positive

b) In a concentration cell, the reduction will take place in the cell's compartment where the concentration is higher

c) Photosynthesis is an electrochemical process

1) (correct)a,b,c

2)a,b

3)b,c

4)a

Solution

Gibbs's free energy is positive as the mark is done on the cell during recharging. Le Chatiler's

principle indicates that reaction is more favorable to reaction when concentration is higher

Photosynthesis is an electrochemical process where water/arsenite is used as an electron.
Hence, the answer is the option (1).

Example.3

3. Which of the following statements is correct?

1)$\mathrm{E}_{\text {cell }}$ and $\Delta \mathrm{G}$ of all reactions are both extensive properties

2)$\mathrm{E}_{\text {cell }}$ and $\Delta \mathrm{G}$ of all reactions are both intensive properties

3) (correct)$E_{\text {cell }}$ is an intensive property while $\Delta G$ extensive property

4)$E_{\text {cell }}$ is an extensive property while $\Delta G$ intensive property

Solution

$E_{\text {cell }}$ is independent of the number of moles and so is intensive

But $\Delta G=-n F E$ it depends on n. Hence, it is an extensive property

Hence, the answer is the option (3).

Example.4

4. Given :

$\begin{aligned} &(i) C(\text { graphite })+\mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g}) ; \\ & \Delta r H^{\ominus}=x \mathrm{kJmol}^{-1}\end{aligned}$

$\begin{aligned} & \quad(\text { ii }) C(\text { graphite })+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g}) \\ & \Delta r H^{\Theta}=y \mathrm{kJmol}^{-1}\end{aligned}$

$\begin{aligned} & \quad(i i i) \mathrm{CO}(\mathrm{g})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g}) \\ & \Delta r H^{\Theta}=z k \mathrm{kmol}^{-1}\end{aligned}$

Based on the above thermochemical equations, find out which one of the following algebraic relationships is correct.

1)$y=2 z-x$

2)$x=y-z$

3)$z=x+y$

4) (correct)$x=y+z$

Solution

Introduction to Electrochemistry -

ELECTROCHEMISTRY
Electrochemistry is the branch of science that deals with transforming chemical energy into electrical energy and vice versa or the relationship between electrical and chemical energy produced in a redox reaction.

Electrolytic Cell
Consider the following redox reaction:

$\mathrm{Zn}(\mathrm{s})+\mathrm{Cu}^{2+}(\mathrm{aq}) \longrightarrow \mathrm{Cu}(\mathrm{s})+\mathrm{Zn}^{2+}(\mathrm{aq})$

Zn displaces copper ions (Cu2+) from aqueous solution in the above reaction. This reaction can be achieved very easily in practice. Put a Zn rod into a solution of CuSO₄ (containing Cu²⁺ ions). It is observed that the blue color of the CuSO₄ solution disappears after some time. In this situation, Zn loses 2 electrons per atom, and Cu²⁺ ions in the solution accept them. Cu²⁺ ions from the solution in this manner are deposited in the form of solid Cu and Zn goes into the solution as Zn²⁺ (colorless). The reaction can well be understood in terms of two half-reactions:

Oxidation: $\quad \mathrm{Zn}(\mathrm{s}) \longrightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-}$
Reduction: $\mathrm{Cu}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{Cu}(\mathrm{s})$

Now, we can make the same reaction occur even if the copper ions and zinc rod are not in direct contact. Suppose we put the Cu2+ ions and Zn rod in two separate containers connect the two by a conducting metallic wire and introduce an inverted U-shaped instrument (called a salt-bridge). In that case, electrons will still be transferred through the connecting cables. The electrons from the Zn rod travel to Cu²⁺ ions through the connecting wires and the same reaction occurs.

-

$C_{(\text {graphite })}+\mathrm{O}_2(g) \rightarrow \mathrm{CO}_2(g) \Delta_r H^0=x \mathrm{KJ} / \mathrm{mol}---(1)$

$C_{(\text {graphite })}+\frac{1}{2} \mathrm{O}_2(g) \rightarrow \mathrm{CO}_2(\mathrm{~g}) \Delta_r H^0=y \mathrm{KJ} / \mathrm{mol}----(2)$

$\mathrm{CO}(\mathrm{g})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g}) \Delta_r H^0=z \mathrm{KJ} / \mathrm{mol}-----(3)$

$(1)=(2)+(3)$

$X=Y+Z$
Therefore,option(4) is correct

Example.5

5. Which of the following statements is incorrect about the electrolytic cell?

1)It converts electrical energy into chemical energy.

2)The non-spontaneous reaction is made spontaneous by the help of electricity.

3) (correct)$\Delta \mathrm{G}<\mathrm{O}$ for the reaction in electrolytic cell

4)$\Delta \mathrm{G}>\mathrm{O}$ for the reaction in electrolytic cell

Solution

In an electrolytic cell, a non-spontaneous reaction is made spontaneous by the help of electrical energy.

$\therefore \Delta \mathrm{G}>\mathrm{O}$ for the reaction in electrolytic cel

Hence, the answer is the option (3).

Summary

Electrochemistry is a concept that allows for the design and optimization of various electrochemical processes and devices, impacting industries ranging from energy storage to material processing. Electrochemistry gives us so many applications such as batteries, electroplating, corrosion, and fuel cells. Batteries: Electrochemical cells that store and release electrical energy. Common types include lead-acid, nickel-cadmium, and lithium-ion batteries. Electroplating: A process that uses electrolysis to deposit a layer of metal onto a surface. It is used for corrosion protection and decorative purposes. Corrosion: The degradation of metals due to electrochemical reactions with their environment. Understanding electrochemistry helps in developing methods to prevent or slow down corrosion. Fuel Cells: Devices that convert chemical energy directly into electrical energy through redox reactions. They are used in applications like powering vehicles and space missions. Electrochemistry has its fundamentals such as redox reactions and electrochemical cells, which include galvanic and electrochemical cells. In Redox Reactions electrochemistry primarily deals with redox reactions, where the oxidation state of substances changes. Oxidation involves the loss of electrons, and reduction consists of the gain of electrons. Electrochemical cells are devices that convert chemical energy into electrical energy or vice versa. There are two main types: Galvanic Cells (Voltaic Cells): Generate electricity from spontaneous redox reactions. Example: Daniel's cell. Electrolytic Cells use electrical energy to drive non-spontaneous reactions—for example: The electrolysis of water.