Electrochemical Series: Definition and Its Applications

The electrochemical series is a key concept in Chemistry, listing elements and ions based on their standard electrode potentials. This series helps predict the behaviour of elements in redox reactions, determine reactivity, and understand the direction of electron flow in electrochemical cells. This page aims to simplify the electrochemical series and explain its importance for students, making it easier to grasp its real-world applications, such as in batteries, corrosion prevention, and electroplating. By exploring this series, you’ll learn how chemists utilise it to make predictions about chemical reactions and design efficient systems.

Introduction

The electrochemical series ranks elements based on their ability to either lose or gain electrons. Elements that tend to lose electrons easily are called electropositive, while elements that easily accept electrons are called electronegative. This categorisation helps in predicting the behavior of metals and non-metals in various reactions, particularly in redox (oxidation-reduction) processes.

• Electropositive Elements: Electropositive elements are those that readily lose electrons, such as alkali metals like sodium (Na) and potassium (K). These metals are good reducing agents.

• Electronegative Elements: Electronegative elements tend to gain electrons in reactions, such as halogens like fluorine (F) and chlorine (Cl). They are strong oxidising agents.

Electrochemical Series Chart

The electrochemical series chart provides a clear visual representation of how different elements and ions compare in terms of their tendency to undergo reduction or oxidation.

Elements at the top of the series are more electropositive and are easily oxidised, losing electrons. Elements lower down are more electronegative and are easily reduced, gaining electrons. This chart helps understand corrosion and compatibility between different metals. For instance, metals that are far apart in the series, such as zinc and copper, will corrode more easily when in contact with each other.

Tricks to Remember the Electrochemical Series:

The Electrochemical Series lists elements and their ions based on their electrode potential. A trick to remember the series is to use a mnemonic or phrase where the first letter of each word corresponds to the symbols of elements in the series.

For example, to recall the standard electrode potential series, you could use this simple mnemonic for some of the key elements:

This phrase includes some common elements in the electrochemical series in the order of their reactivity (from most reactive to least reactive).

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Here’s how the mnemonic breaks down:

• P: Potassium (K)

• S: Sodium (Na)

• C: Calcium (Ca)

• M: Magnesium (Mg)

• A: Aluminum (Al)

• C: Carbon (C)

• Z: Zinc (Zn)

• I: Iron (Fe)

• S: Silver (Ag)

• M: Mercury (Hg)

Applications of Electrochemical Series

1. Oxidising and Reducing Strengths

The electrochemical series helps identify strong oxidising agents, which are at the top of the series, and strong reducing agents, which are at the bottom.

Calculation of Standard EMF of Electrochemical Cells

The standard EMF of a cell can be calculated using the formula:

\[E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}\]

For example, consider the reaction:

\[2Ag^+ + Cd \rightarrow 2Ag + Cd^{2+}\]

The standard reduction potentials are:

$Ag^+/Ag = +0.80 \, \text{V}, \quad Cd^{2+}/Cd = -0.40 \, \text{V}$.

Using the formula:

$E^\circ_{\text{cell}} = 0.80 - (-0.40) = 1.20 \, \text{V}$.

This positive EMF indicates a spontaneous reaction.

2. Predicting the Feasibility of Redox Reactions

The spontaneity of a reaction can be determined by the free energy change ($\Delta G^\circ$), which is related to the cell EMF by:

\[\Delta G^\circ = -nFE^\circ_{\text{cell}}\]

If $E^\circ_{\text{cell}}$ is positive, the reaction is spontaneous.

Example: To check if copper sulfate can be stored in a nickel vessel:

$Ni^{2+}/Ni = -0.25 \, \text{V}, \quad Cu^{2+}/Cu = +0.34 \, \text{V}$.

For the reaction:

$Ni + CuSO_4 \rightarrow NiSO_4 + Cu$

The EMF is:

$E^\circ_{\text{cell}} = 0.34 - (-0.25) = 0.59 \, \text{V}$.

Since the EMF is positive, copper sulfate cannot be stored in a nickel vessel.

3. Predicting the Product of Electrolysis

During electrolysis, ions with higher reduction potentials are discharged at the cathode. For example, when NaCl is electrolysed:

At the cathode, the $H^+$ ion is reduced (deposited) because its reduction potential ($0.00 \, \text{V}$) is higher than that of $Na^+$ ($-2.71 \, \text{V}$). At the anode, $OH^-$ is oxidised because it has a lower reduction potential ($0.40 \, \text{V}$) compared to $Cl^-$ ($1.36 \, \text{V}$).

Electrochemical Series Important Points

• The standard reduction potential compares the tendency of a substance to gain electrons against hydrogen ($0 \, \text{V}$).

• Strong reducing agents have low or negative reduction potentials and are found at the bottom of the series, while strong oxidising agents have high or positive reduction potentials and are found at the top.

• As one moves down the series, the electropositivity and activity of metals increase, while for non-metals, these decrease.

Solved Problems

Problem 1: Spontaneity of Reaction

For the reaction:

$Fe^{3+} + 2Cl^- \rightarrow Fe^{2+} + Cl_2$

With the standard reduction potentials:

$Fe^{2+}/Fe = -0.44 \, \text{V}, \quad Cl_2/Cl^- = 1.36 \, \text{V}$.

The EMF is:

$E^\circ_{\text{cell}} = 1.36 - (-0.44) = 1.80 \, \text{V}$.

A positive EMF indicates the reaction occurs spontaneously.

Problem 2: Identifying the Strongest Reducing Agent

Given the following half-reactions:

$Zn^{2+} + 2e^- \rightarrow Zn \, (E^\circ = -0.76 \, \text{V})$,

$Cr^{3+} + 3e^- \rightarrow Cr \, (E^\circ = -0.74 \, \text{V})$,

$Cu^{2+} + 2e^- \rightarrow Cu \, (E^\circ = +0.34 \, \text{V})$,

$Fe^{3+} + e^- \rightarrow Fe^{2+} \, (E^\circ = +0.77 \, \text{V})$.

Zinc, with the lowest reduction potential ($-0.76 \, \text{V}$), is the strongest reducing agent.

Problem 3: Calculating EMF of the Cell

For the cell:

$Cu^{2+} + Fe \rightarrow Cu + Fe^{2+}$

With the following standard oxidation potentials:

$Cu \rightarrow Cu^{2+} + 2e^- \, (E^\circ = -0.34 \, \text{V})$,

$Fe \rightarrow Fe^{2+} + 2e^- \, (E^\circ = 0.41 \, \text{V})$.

The EMF is:

$E^\circ_{\text{cell}} = -0.34 - (-0.41) = 0.07 \, \text{V}$.

Conclusion

The electrochemical series is a valuable tool for understanding redox reactions, predicting the feasibility of reactions, and determining the behavior of metals and non-metals in various chemical processes. Whether calculating cell EMFs, determining spontaneity, or identifying strong oxidizers and reducers, the electrochemical series is essential for a wide range of applications, including corrosion control, electroplating, and electrochemical cells.