R3.2.3 Reactivity of elements based on periodic trends

Relative Ease of Oxidation and Reduction, and Displacement Reactions

Imagine you're a chemist holding two shiny metal strips—zinc and copper.
You dip the zinc strip into a blue copper sulfate solution and observe as the solution fades to colorless, while a reddish-brown solid forms on the zinc.
What just happened?

This is a displacement reaction, a striking demonstration of the relative ease of oxidation and reduction.

Predicting Reactivity Trends in Metals and Halogens

Metals: The Ease of Oxidation

Metals are excellent reducing agents because they tend to lose electrons and form positive ions (cations) during chemical reactions.

But why do some metals lose electrons more readily than others?
This depends on their position in the periodic table and their intrinsic reactivity.
- Group 1 Metals (Alkali Metals):
  - These metals are highly reactive because they have only one valence electron, which is easily lost.
  - Reactivity increases as you move down the group because the outermost electron is farther from the nucleus, experiencing weaker electrostatic attraction.
- Other Metals:
  - For metals outside Group 1, their reactivity can be tested experimentally using displacement reactions.
  - A more reactive metal will displace a less reactive metal from its ionic solution.

Example

Zinc and Copper Displacement

When zinc metal is placed in a solution of copper(II) sulfate ${CuSO}_{4}$ , the following reaction takes place:
$Zn (s) + {Cu}^{2 +} (a q) \to {Zn}^{2 +} (a q) + Cu (s)$
Here, zinc is oxidized to ${Zn}^{2 +}$ (it loses electrons), and copper(II) ions are reduced to copper metal (they gain electrons). This demonstrates that zinc is more reactive than copper.

Common Mistake

Many students forget that a displacement reaction only occurs if the metal in solid form is more reactive than the metal in the ionic solution. If the metal in solution is more reactive, no reaction will occur.

Halogens: The Ease of Reduction

Halogens, as non-metals, act as oxidizing agents because they readily gain electrons to form negative ions (anions).
However, their reactivity decreases as you move down Group 17 of the periodic table. Why does this trend occur?
- Fluorine:
  - Fluorine is the most reactive halogen due to its small atomic radius and high electronegativity, which allow it to attract electrons very effectively.
- Chlorine, Bromine, and Iodine:
  - As you move down the group, the atomic radius increases, and the outer electrons are farther from the nucleus.
  - This makes it harder for the atom to attract additional electrons, reducing reactivity.

Example

Chlorine and Bromide Ions

When chlorine gas ${Cl}_{2}$ is bubbled into a solution of potassium bromide $KBr$ , the following reaction occurs:
${Cl}_{2} (g) + 2 {Br}^{-} (a q) \to 2 {Cl}^{-} (a q) + {Br}_{2} (a q)$
In this reaction, chlorine is reduced to ${Cl}^{-}$ , and bromide ions are oxidized to bromine. This reaction happens because chlorine is more reactive than bromine.

Tip

To remember the reactivity trend of halogens, think:

Fluorine > Chlorine > Bromine > Iodine.

Reactivity decreases as you move down Group 17.

Displacement Reactions

Displacement reactions are a practical tool for comparing the reactivity of metals and halogens.
These reactions involve the transfer of electrons, making them redox reactions.

Metal Displacement Reactions

In a metal displacement reaction, a more reactive metal displaces a less reactive metal from its compound.

Note

This is a direct application of the reactivity of metals.

General Reaction: ${Metal}_{1} (s) + {Metal}_{2}^{n +} (a q) \to {Metal}_{1}^{n +} (a q) + {Metal}_{2} (s)$

Self review

Can magnesium displace zinc from zinc sulfate? Why or why not?

Halogen Displacement Reactions

In a halogen displacement reaction, a more reactive halogen displaces a less reactive halogen from its ionic compound.

Note

These reactions highlight the decreasing reactivity of halogens down Group 17.

General Reaction: ${Halogen}_{1} (g) + 2 {Halide}_{2}^{-} (a q) \to 2 {Halide}_{1}^{-} (a q) + {Halogen}_{2} (a q)$

Example

Example: Chlorine and Bromide Ions ${Cl}_{2} (g) + 2 {Br}^{-} (a q) \to 2 {Cl}^{-} (a q) + {Br}_{2} (a q)$
Chlorine displaces bromine because it is more reactive. The solution changes color as bromine is formed.

Common Mistake

Students often confuse the roles of halogens and halides. Remember, halogens (e.g., ${Cl}_{2}$ ) are reduced, while halides (e.g., ${Br}^{-}$ ) are oxidized.

Reflection

Self review

Can you predict the products of the reaction between magnesium and copper sulfate? What about iodine and potassium chloride?

Theory of Knowledge

How does the concept of reactivity relate to the ethical use of natural resources, such as metals and halogens, in industrial applications?
Consider the balance between technological advancement and environmental sustainability.

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R3.2.3 Reactivity of elements based on periodic trends

Relative Ease of Oxidation and Reduction, and Displacement Reactions

Predicting Reactivity Trends in Metals and Halogens

Metals: The Ease of Oxidation

Zinc and Copper Displacement

Halogens: The Ease of Reduction

Chlorine and Bromide Ions

Displacement Reactions

Metal Displacement Reactions

Halogen Displacement Reactions

Reflection

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Reactivity and Displacement Reactions