The Electrochemical Series Revision Notes for Leaving Cert Chemistry

The Electrochemical Series

What is the electrochemical series?

The electrochemical series is a powerful tool in chemistry that ranks elements based on how easily they give up electrons. Think of it as a "league table" of metals, arranged from the most reactive at the top to the least reactive at the bottom.

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Key definition: The electrochemical series lists elements in order of their standard electrode potentials. This means it shows the decreasing tendency for each element to be oxidised (lose electrons) under standard conditions.

Elements higher up in the series have a stronger tendency to lose electrons and become positive ions. Elements lower down prefer to stay as neutral atoms or gain electrons.

Understanding the hydrogen half-cell

To create this series, chemists needed a reference point - something to compare all other elements against. They chose the hydrogen half-cell as their standard.

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The standard hydrogen half-cell consists of:

A platinum electrode (which doesn't react)
A 1M solution of H⁺ ions
Hydrogen gas at atmospheric pressure bubbling over the platinum

This setup is given a standard electrode potential of exactly 0.00 volts. Every other element is then compared to this hydrogen standard to determine where it sits in the electrochemical series.

Order of elements in the electrochemical series

The electrochemical series follows this general pattern (from most reactive to least reactive):

Most reactive (top of series):

Metals like potassium, sodium, calcium, magnesium, aluminium
These metals lose electrons very easily
They are easily oxidised under standard conditions

Middle section:

Zinc, iron, lead, hydrogen

Least reactive (bottom of series):

Copper, mercury, silver, gold
These metals have a very small tendency to lose electrons
They may even occur in nature as pure metals

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Key Principle: Metals near the top of the series have a greater tendency to lose their outer electrons than metals lower down. This explains why some metals corrode easily while others like gold remain unreactive.

Applications of the electrochemical series

Predicting displacement reactions

The electrochemical series allows us to predict which displacement reactions will occur. A metal higher in the series will always displace a metal lower down from a solution of its ions.

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Rule: More reactive metal displaces less reactive metal

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Worked Example: Copper and Silver Displacement

When copper metal is placed in silver nitrate solution:

Copper is higher than silver in the electrochemical series
Copper displaces silver from the solution
Silver metal precipitates out as distinctive crystals
The copper goes into solution as Cu²⁺ ions

Other examples of displacement reactions:

Magnesium displaces copper: $\text{Mg} + \text{Cu}^{2+} \rightarrow \text{Mg}^{2+} + \text{Cu}$
Iron displaces copper: $\text{Fe} + \text{Cu}^{2+} \rightarrow \text{Fe}^{2+} + \text{Cu}$
Copper displaces silver: $\text{Cu} + 2\text{Ag}^+ \rightarrow \text{Cu}^{2+} + 2\text{Ag}$

Galvanising and corrosion protection

Understanding the electrochemical series helps explain how we can protect metals from corrosion.

Galvanising involves coating iron or steel with zinc. Since zinc is higher in the electrochemical series than iron:

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Zinc has a greater tendency to lose electrons than iron
Zinc corrodes preferentially, protecting the iron underneath
Even if the zinc coating is scratched, it continues to protect the iron
This is why galvanised items like buckets and wheelbarrows are coated with zinc

Coating with unreactive metals: Sometimes iron or steel is coated with metals lower in the electrochemical series, such as tin or chromium. However, if this coating is damaged and the base metal is exposed to air and water, corrosion will occur rapidly.

Cathodic protection

This method uses the electrochemical series principle to prevent corrosion on a larger scale. A more reactive metal (higher in the series) is placed in contact with the metal that needs protection.

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Sacrificial anodes: In this method, blocks of zinc are attached to steel structures like ship hulls or underground pipes. The zinc acts as the anode and corrodes instead of the steel. The zinc needs to be replaced periodically, but this protects the much more expensive steel structure.

This technique is called cathodic protection because the steel becomes the cathode in the electrochemical cell and doesn't corrode.

Key examples and reactions

Most common corrosion example: The rusting of iron and steel occurs when these metals are exposed to oxygen and water. Since iron is relatively high in the electrochemical series, it readily loses electrons to form iron oxide (rust).

Preventing corrosion methods:

Painting and greasing - creates a barrier preventing oxygen and water from reaching the metal
Galvanising - coating with zinc (more reactive metal)
Using sacrificial anodes - attaching more reactive metals like zinc blocks

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Industrial Application: Copper Extraction

Scrap iron can be used to extract copper from copper sulphate solutions because iron is higher than copper in the electrochemical series, so iron displaces copper from solution.

Summary

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Key Points to Remember:

The electrochemical series ranks elements by their tendency to lose electrons - most reactive at the top, least reactive at the bottom
More reactive metals displace less reactive metals from solutions of their ions - this is the key to predicting displacement reactions
The hydrogen half-cell is the standard reference point with a potential of 0.00 volts that all other elements are compared against
Galvanising works because zinc is more reactive than iron - zinc corrodes preferentially to protect the iron underneath
Cathodic protection uses sacrificial anodes - more reactive metals like zinc protect less reactive metals like steel by corroding instead of them

The Electrochemical Series (Leaving Cert Chemistry): Revision Notes