Lesson Note on Redox Reactions: Agents, IUPAC Names and Balancing Equations for SSS 2

Class: Senior Secondary School 2 (SS2 / SSS 2)
Term: 3rd Term
Week: 4
Age: 16 years
Duration: 45 minutes
Subject: Chemistry
Curriculum Theme: Chemical Reactions
Previous Lesson: Redox Reactions: Meaning, Definitions and Oxidation Numbers.
Topic: OXIDATION-REDUCTION (REDOX) REACTIONS
Subject Matter: Connection of oxidation numbers with IUPAC name, oxidizing and reducing agents, redox equation.

Specific Objectives

By the end of the lesson, pupils should be able to:

Cognitive Domain:

• Define oxidation, reduction, and redox reactions.
• State the rules for assigning oxidation numbers to elements in compounds and ions.
• Identify oxidizing and reducing agents in a given redox reaction.
• Explain the connection between oxidation numbers and the IUPAC naming of inorganic compounds.

Affective Domain:

• Appreciate the importance of redox reactions in everyday life and industrial processes.
• Show interest in systematically balancing redox equations.

Psychomotor Domain:

• Assign oxidation numbers to elements in various inorganic compounds and ions.
• Write the IUPAC names of compounds given their formulas, and write formulas given their IUPAC names, using oxidation numbers.
• Write half-reactions (oxidation and reduction) for a given redox process.
• Balance simple redox equations using the ion-electron method.

Social Domain:

• Participate actively in class discussions on redox reactions.
• Collaborate with peers to solve problems involving oxidation numbers and balancing redox equations.

Reference Materials

The following resources were used in planning this lesson:

• Senior Secondary Schools Education Curriculum
• State Unified Scheme of Work
• Osei Yaw Ababio. (2009). New School Chemistry for Senior Secondary Schools. Africana First Publishers Plc.
• https://www.chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/20%3A_Electrochemistry/20.1%3A_Oxidation-Reduction_Reactions
• https://www.iupac.org/what-we-do/books/redbook/redbook-chapter-ir-9/

Instructional Materials

The teacher will teach this lesson with the aid of:

• Whiteboard or chalkboard
• Markers or chalk
• Periodic table charts
• Samples of Iron(II) oxide and potassium permanganate (KMnO4)

Rationale for the Lesson

This lesson helps pupils understand fundamental chemical processes involving electron transfer, which are essential for explaining phenomena like corrosion, batteries, and many industrial chemical reactions. It enables pupils to systematically name inorganic compounds and balance complex chemical equations.

Prerequisite/Previous Knowledge

Pupils should have prior knowledge of atoms, ions, chemical symbols, formulas, and basic concepts of chemical reactions and valency.

Lesson Content/Board Summary

OXIDATION-REDUCTION (REDOX) REACTIONS

Definition of Oxidation and Reduction

Oxidation is a chemical process involving the loss of electrons, an increase in oxidation number, or the gain of oxygen/loss of hydrogen.

Reduction is a chemical process involving the gain of electrons, a decrease in oxidation number, or the loss of oxygen/gain of hydrogen.

A Redox reaction is a chemical reaction in which both oxidation and reduction occur simultaneously.

Oxidation Numbers

Oxidation number (or oxidation state) is the charge an atom would have if electrons were transferred completely.

The following are rules for assigning oxidation numbers:

• The oxidation number of an element in its free or uncombined state is zero (e.g., O2, Na, Fe).
• The oxidation number of a monatomic ion is equal to its charge (e.g., Na+ is +1, Cl- is -1).
• The oxidation number of Group 1 metals is always +1 in compounds.
• The oxidation number of Group 2 metals is always +2 in compounds.
• The oxidation number of hydrogen is usually +1 in compounds, except in metal hydrides where it is -1 (e.g., NaH).
• The oxidation number of oxygen is usually -2 in compounds, except in peroxides where it is -1 (e.g., H2O2) and in OF2 where it is +2.
• The sum of the oxidation numbers of all atoms in a neutral compound is zero.
• The sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion.

Examples of Assigning Oxidation Numbers:

Example 1: Determine the oxidation number of N in Pb(NO3)2.

• Step 1: Assign known oxidation numbers. Pb is Group 14, but often +2 in compounds. O is -2.
• Step 2: Let x be the oxidation number of N.
• Step 3: Sum of oxidation numbers = 0.
  (+2) + 2(x + 3(-2)) = 0
  2 + 2(x – 6) = 0
  2 + 2x – 12 = 0
  2x – 10 = 0
  2x = 10
  x = +5
• Conclusion: The oxidation number of N in Pb(NO3)2 is +5.

Example 2: Determine the oxidation number of Fe in FeSO4·7H2O.

• Step 1: Assign known oxidation numbers. In SO4^2-, S is +6 and O is -2. In H2O, H is +1 and O is -2. The water of crystallisation does not affect the oxidation state of Fe.
• Step 2: The SO4 group has a charge of -2. For the compound to be neutral, Fe must have a charge of +2.
• Conclusion: The oxidation number of Fe in FeSO4·7H2O is +2.

Example 3: Determine the oxidation number of Cu in [Cu(NH3)4]2+.

• Step 1: Assign known oxidation numbers. NH3 is a neutral ligand, so its overall oxidation number is 0.
• Step 2: Let x be the oxidation number of Cu.
• Step 3: Sum of oxidation numbers = charge of ion.
  x + 4(0) = +2
  x = +2
• Conclusion: The oxidation number of Cu in [Cu(NH3)4]2+ is +2.

Connection of Oxidation Numbers with IUPAC Name

The IUPAC system uses Roman numerals in parentheses to indicate the oxidation number of the metal ion, especially for transition metals that can have multiple oxidation states.

Examples:

• Pb(NO3)2: Lead(II) nitrate (Pb has an oxidation number of +2)
• FeSO4·7H2O: Iron(II) tetraoxosulphate(VI) heptahydrate (Fe has an oxidation number of +2)
• Iron(II) trioxosulphate(IV) corresponds to FeSO3 (Fe has an oxidation number of +2, S in SO3^2- has +4).

Oxidizing and Reducing Agents

An oxidizing agent (oxidant) is a substance that causes another substance to be oxidized, and in the process, it gets reduced. Oxidizing agents gain electrons and have their oxidation number decreased.

A reducing agent (reductant) is a substance that causes another substance to be reduced, and in the process, it gets oxidized. Reducing agents lose electrons and have their oxidation number increased.

Examples:

• Common Oxidizing Agents: KMnO4, K2Cr2O7, HNO3, H2SO4 (conc.), O2, Cl2.
• Common Reducing Agents: H2S, SO2, H2, Na, Fe2+, Sn2+.

Half-Reactions and Balancing Redox Equations

A redox reaction can be split into two half-reactions: an oxidation half-reaction and a reduction half-reaction.

Steps for Balancing Redox Equations (Ion-Electron Method in Acidic Medium):

1. Separate the overall reaction into two half-reactions (oxidation and reduction).
2. Balance all atoms except O and H in each half-reaction.
3. Balance oxygen atoms by adding H2O molecules to the side deficient in oxygen.
4. Balance hydrogen atoms by adding H+ ions to the side deficient in hydrogen.
5. Balance the charge in each half-reaction by adding electrons (e-) to the more positive side.
6. Multiply each half-reaction by an integer to make the number of electrons lost equal to the number of electrons gained.
7. Add the two balanced half-reactions together and cancel out common species (e.g., electrons, H+, H2O).

Worked Example 1: Balance the reaction: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

• Step 1: Separate into half-reactions.
  Oxidation: Zn → Zn2+
  Reduction: Cu2+ → Cu
• Step 2: Balance atoms (already balanced).
• Step 3: Balance O (no O).
• Step 4: Balance H (no H).
• Step 5: Balance charge by adding electrons.
  Oxidation: Zn → Zn2+ + 2e-
  Reduction: Cu2+ + 2e- → Cu
• Step 6: Electrons are already equal.
• Step 7: Add half-reactions.
  Zn + Cu2+ + 2e- → Zn2+ + 2e- + Cu
  Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) (Balanced)

Worked Example 2: Balance the reaction: MnO4-(aq) + Fe2+(aq) → Mn2+(aq) + Fe3+(aq) (in acidic medium)

• Step 1: Separate into half-reactions.
  Oxidation: Fe2+ → Fe3+
  Reduction: MnO4- → Mn2+
• Step 2: Balance atoms (Fe and Mn are balanced).
• Step 3: Balance O in reduction half-reaction.
  MnO4- → Mn2+ + 4H2O
• Step 4: Balance H in reduction half-reaction.
  8H+ + MnO4- → Mn2+ + 4H2O
• Step 5: Balance charge by adding electrons.
  Oxidation: Fe2+ → Fe3+ + e- (Charge: +2 on left, +3-1=+2 on right)
  Reduction: 8H+ + MnO4- + 5e- → Mn2+ + 4H2O (Charge: +8-1-5=+2 on left, +2 on right)
• Step 6: Multiply oxidation half-reaction by 5 to equal electrons.
  5(Fe2+ → Fe3+ + e-) => 5Fe2+ → 5Fe3+ + 5e-
• Step 7: Add the two balanced half-reactions.
  5Fe2+ + 8H+ + MnO4- + 5e- → 5Fe3+ + 5e- + Mn2+ + 4H2O
  5Fe2+(aq) + 8H+(aq) + MnO4-(aq) → 5Fe3+(aq) + Mn2+(aq) + 4H2O(l) (Balanced)

Teaching Methods/Instructional Techniques

Discussion, Lecture, Demonstration, Question and Answer, Visual Aids

Instructional Procedures

Step 1: Introduction

Time: 5 minutes
Teaching Skill: Set Induction
Teacher’s Activity: The teacher greets the pupils and asks them to recall what they know about valency and how elements combine. The teacher then introduces the topic of oxidation-reduction reactions, stating that it involves electron transfer.
Pupils’ Activity: Pupils respond to questions about valency and listen attentively to the introduction of the new topic.
Learning Point: Pupils recall prior knowledge and are introduced to redox reactions.

Step 2: Explanation of Oxidation and Reduction

Time: 7 minutes
Teaching Skill: Explanation/Definition
Teacher’s Activity: The teacher defines oxidation, reduction, and redox reactions in terms of electron transfer and changes in oxidation numbers, writing the definitions on the board.
Pupils’ Activity: Pupils listen, ask questions for clarity, and copy the definitions into their notebooks.
Learning Point: Pupils understand the fundamental definitions of oxidation, reduction, and redox reactions.

Step 3: Assigning Oxidation Numbers

Time: 8 minutes
Teaching Skill: Explanation/Demonstration
Teacher’s Activity: The teacher explains the rules for assigning oxidation numbers with examples on the board, demonstrating how to calculate oxidation numbers for elements in compounds like Pb(NO3)2, FeSO4·7H2O, [Cu(NH3)4]2+, and KMnO4.
Pupils’ Activity: Pupils pay attention to the rules, follow the examples, and attempt to calculate oxidation numbers for other simple compounds provided by the teacher.
Learning Point: Pupils learn how to correctly assign oxidation numbers to elements in various chemical species.

Step 4: Connection to IUPAC Naming

Time: 6 minutes
Teaching Skill: Explanation/Application
Teacher’s Activity: The teacher explains how oxidation numbers are used in the IUPAC naming of inorganic compounds, especially for metals with variable oxidation states, giving examples like Iron(II) trioxosulphate (IV) and Lead(II) nitrate.
Pupils’ Activity: Pupils understand the role of Roman numerals in IUPAC nomenclature and practice writing names or formulas based on oxidation numbers.
Learning Point: Pupils connect oxidation numbers to the systematic naming of inorganic compounds.

Step 5: Oxidizing and Reducing Agents

Time: 5 minutes
Teaching Skill: Explanation/Identification
Teacher’s Activity: The teacher defines oxidizing and reducing agents and provides examples of common oxidants and reductants. The teacher emphasizes that the oxidizing agent is reduced and the reducing agent is oxidized.
Pupils’ Activity: Pupils identify oxidizing and reducing agents in simple reactions provided by the teacher.
Learning Point: Pupils can differentiate between oxidizing and reducing agents.

Step 6: Writing Half-Reactions and Balancing Equations

Time: 9 minutes
Teaching Skill: Demonstration/Problem-Solving
Teacher’s Activity: The teacher explains the steps for balancing redox equations using the ion-electron method in acidic medium and works through the two examples on the board (Zn + Cu2+ → Zn2+ + Cu and MnO4- + Fe2+ → Mn2+ + Fe3+).
Pupils’ Activity: Pupils follow the step-by-step balancing process, ask questions, and attempt to balance similar equations.
Learning Point: Pupils learn the procedure for writing half-reactions and balancing redox equations.

Step 7: Evaluation/Review

Time: 5 minutes
Teaching Skill: Questioning/Assessment
Teacher’s Activity: The teacher evaluates the learning by asking the following questions:

1. Define oxidation and reduction in terms of electron transfer.
2. State any three rules for assigning oxidation numbers.
3. Determine the oxidation number of Mn in KMnO4.
4. In the reaction: Fe + CuSO4 → FeSO4 + Cu, identify the oxidizing and reducing agents.

Pupils’ Activity: Pupils answer orally and in writing.
Learning Point: Pupils demonstrate understanding of the lesson.

Step 8: Conclusion

Time: 2 minutes
Teaching Skill: Summarization
Teacher’s Activity: The teacher summarizes the key learning points of the lesson, emphasizing the importance of redox reactions in chemistry. The teacher assigns homework for pupils to balance more redox equations.
Pupils’ Activity: Pupils listen to the summary and copy down the assigned homework.
Learning Point: Pupils consolidate their understanding and prepare for further practice.

Lesson Keywords

• Oxidation – Loss of electrons or increase in oxidation number.
• Reduction – Gain of electrons or decrease in oxidation number.
• Redox Reaction – A chemical reaction involving simultaneous oxidation and reduction.
• Oxidation Number – The charge an atom would have if electrons were completely transferred.
• Oxidizing Agent – A substance that causes oxidation and is itself reduced.
• Reducing Agent – A substance that causes reduction and is itself oxidized.
• IUPAC Name – Systematic name for chemical compounds.

Differentiation

For struggling learners, the teacher will provide simpler examples for assigning oxidation numbers and balancing redox equations, offering more one-on-one guidance. Advanced learners will be given more complex redox reactions to balance and asked to research real-life applications of redox chemistry, such as in batteries or industrial electrolysis.

Note for teachers using this lesson plan

Ensure pupils have a solid grasp of basic chemical formulas and ion charges before delving into oxidation numbers. Emphasize the systematic approach to balancing redox equations, as this can be challenging. Use visual aids like a periodic table to help pupils with oxidation number rules. Encourage active participation and collaborative problem-solving.