Redox Reactions in term of Electron Exchange, Oxidation Number | Chemistry for Grade 11 (IGCSE) PDF Download

Redox & Electron Transfer

 Oxidation & Reduction 

• Redox reactions involve electron transfer between substances
• Oxidation refers to a reaction where an element, ion, or compound loses electrons, leading to an increase in its oxidation number. For instance, when silver reacts with chlorine, silver is oxidized to silver ions
  Ag → Ag⁺ + e^-

• Reduction is a chemical reaction where an element, ion, or compound gains electrons, leading to a decrease in its oxidation number. This process is often represented in a half equation.
  • Example: When oxygen reacts with magnesium, oxygen is reduced to oxide ions.
    O2 + 4e^- → 2O^2-

Example: Identifying Redox Reactions

• In the reaction between zinc and copper sulfate, zinc displaces copper from the compound, forming zinc sulfate and elemental copper.
• This reaction can be represented as: Zn + CuSO₄ → ZnSO₄ + Cu.
• The ions present in the equation include: Zn(s), Cu²⁺(aq), SO₄^2-(aq), Zn²⁺(aq), SO₄^2-(aq), Cu(s).
• The spectator ions, which do not participate in the reaction, are represented by SO₄^2- (aq).
• These spectator ions can be removed, simplifying the equation to focus on the reactive species.
• Upon analyzing the ionic equation, it can be separated into two half equations by accounting for the transfer of electrons.
• Zn(s) → Zn²⁺(aq) + 2e^-
• Cu²⁺(aq) + 2e^- → Cu(s)
• Observing the half equations reveals that zinc undergoes oxidation (loses electrons), while copper ions experience reduction (gain electrons).

Identifying Redox Reactions

• The oxidation number (also known as oxidation state) represents the level of oxidation (or reduction) of an atom or ion within a compound.
• It signifies the number of electrons that an atom has either gained, lost, or shared during the formation of a compound.
• Oxidation numbers are crucial for monitoring electron movement in redox processes. They are denoted by a positive or negative sign followed by a number (unlike charges, which are represented by a number followed by a positive or negative sign).
• For instance, aluminum typically exhibits an oxidation state of +3 in compounds.
• There are straightforward rules to assist in determining the oxidation number of any element.
• Table of Rules for Assigning Oxidation Numbers:
  
• Redox reactions are recognized by changes in the oxidation numbers as reactants transform into products.

Testing Redox Reactions

• The examination of redox reactions involves observing a color alteration in the solution under analysis.
• Two prevalent examples include acidified potassium manganate(VII) and potassium iodide.
• Potassium manganate(VII) (KMnO₄) functions as an oxidizing agent, commonly utilized to detect reducing agents.
• When acidified potassium manganate(VII) interacts with a reducing agent, its purple hue fades to colorless.
• Diagram illustrating the color shift when potassium manganate(VII) reacts with a reducing agent.
  
• Potassium iodide (KI) is a reducing agent commonly employed to identify oxidizing agents.
• When added to an acidic solution of an oxidizing agent like aqueous chlorine or hydrogen peroxide (H₂O₂), the solution adopts a red-brown color due to iodine (I₂) formation.
  2KI(aq) + H₂SO₄(aq) + H₂O₂(aq) → I₂(aq) + K₂SO₄(aq) + 2H₂O(l)
• Potassium iodide acts as a reducing agent because it undergoes oxidation by losing electrons while hydrogen peroxide is reduced.
  2I^- →  I₂ + 2e^-
• Diagram to show the colour change when potassium iodide is added to an oxidising agent:
  

Oxidizing & Reducing Agents

Oxidizing Agent

• Definition: An oxidizing agent is a substance that causes another substance to be oxidized while itself getting reduced.
• Key terms: Oxidizes, Reduced
• Examples: Oxidizing agents include hydrogen peroxide, fluorine, and chlorine.

Reduction and Oxidation in Chemistry

• A reducing agent is a substance that reduces another substance and becomes oxidized in the process.
• Common examples of reducing agents include carbon and hydrogen.
• Reduction is crucial in the chemical industry for extracting metals from ores.

Example

• CuO + H₂ → Cu + H₂O
• In the given reaction, hydrogen is serving as the reducing agent, causing the reduction of CuO, while undergoing oxidation itself by losing electrons.
  H₂ → 2H⁺ + 2e^-
• Copper oxide (CuO) acts as the oxidizing agent, causing the oxidation of hydrogen, while itself being reduced to copper (Cu) by gaining electrons.
  Cu2⁺ +2e^- →  Cu