Reactivity 3.2.3 — Predicting the Relative Ease of Oxidation and Reduction Using Periodic Trends

We study about how periodic trends allow us to predict how easily an element is oxidized or reduced. These trends are particularly evident when comparing:

• Metals across the periodic table — their tendency to lose electrons (oxidation).
• Non-metals, especially halogens — their tendency to gain electrons (reduction).
• The results of displacement reactions between metals and metal ions, or halogens and halide ions.

Oxidation and Reduction: Periodic Trends

Metals:

• Metals tend to lose electrons and form positive ions — they undergo oxidation.
• As you go down Group 1 or Group 2, it becomes easier for metals to lose electrons because:
  • Atomic radius increases
  • Ionization energy decreases
  • Shielding increases
• Thus, lower group metals are more reactive (more easily oxidized).

Halogens (Group 17):

• Halogens tend to gain electrons — they undergo reduction.
• As you go up Group 17:
  • Atomic radius decreases
  • Electron affinity increases
  • Effective nuclear attraction is stronger
• Thus, fluorine is the strongest oxidizing agent in the group (gains electrons most readily).

Displacement Reactions

Metal–Metal Ion Displacement:

A more reactive metal (stronger reducing agent) displaces a less reactive metal from its aqueous salt solution:

\( \text{Zn(s)} + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu(s)} \)

Here, zinc is oxidized and copper is reduced. This means zinc is more easily oxidized than copper.

Halogen–Halide Displacement:

A more reactive halogen (stronger oxidizing agent) displaces a less reactive halide ion from solution:

\( \text{Cl}_2(aq) + 2\text{Br}^-(aq) \rightarrow 2\text{Cl}^-(aq) + \text{Br}_2(aq) \)

This shows that chlorine is a stronger oxidizing agent than bromine.

Interpreting Reactivity Using Supplied Data

You are not expected to memorize the full reactivity series. Instead, you will be provided data tables, typically showing the results of metal displacement reactions or halogen reactivities.

Your job will be to interpret these results to:

• Rank elements in order of oxidizing/reducing strength
• Predict whether a given redox reaction will occur spontaneously

Oxidation Half-Equations for Metals

• These show the loss of electrons:

\( \text{Mg(s)} \rightarrow \text{Mg}^{2+}(aq) + 2e^- \)

Reduction Half-Equations for Halogens

\( \text{Br}_2(aq) + 2e^- \rightarrow 2\text{Br}^-(aq) \)

Why Metals Differ in Ease of Oxidation

• Stronger metallic bonding in transition metals can make them less reactive (e.g. Fe, Cu, Zn).
• Alkali metals have only one electron to lose, so are very easily oxidized.

Why Halogens Differ in Ease of Reduction

• Electronegativity and small size make F₂ and Cl₂ excellent oxidizing agents.
• Down the group, reduced attraction for incoming electrons makes I₂ the weakest oxidizer.