Chapter 6: Advanced Oxidation-Redox Applications and Calculations

Overview

Building on the fundamental concepts of electrochemistry, this chapter delves into advanced applications of oxidation-reduction reactions in analytical chemistry, industrial processes, and environmental technology. We explore quantitative redox methods including titrations, redox indicators, and stoichiometric calculations. Understanding these advanced redox applications is crucial for chemical analysis, water treatment, battery technology, and many industrial processes that rely on controlled electron transfer reactions.

Redox Fundamentals Review

Learning Objectives

After studying this chapter, you should be able to:

• Perform redox titrations using common oxidizing and reducing agents
• Understand and use redox indicators for endpoint detection
• Calculate concentrations and quantities in redox reactions
• Apply redox principles in quantitative analysis
• Understand industrial redox processes and their applications
• Analyze environmental redox applications in water treatment and pollution control
• Solve complex redox stoichiometry problems

6.1 Redox Titrations

What are Redox Titrations?

Redox titrations are volumetric analysis methods where a redox reaction is used to determine the concentration of an unknown solution. One solution of known concentration (titrant) is added to another solution of unknown concentration until the reaction is complete (equivalence point).

Common Redox Titrations

1. Acidified Potassium Manganate(VII) Titrations

Principle: $MnO_4^-$ oxidizes $Fe^{2+}$ to $Fe^{3+}$ in acidic medium Reaction: $MnO_4^-(aq) + 5Fe^{2+}(aq) + 8H^+(aq) \rightarrow Mn^{2+}(aq) + 5Fe^{3+}(aq) + 4H_2O(l)$

Indicator: Self-indicator (purple $MnO_4^-$ becomes colorless $Mn^{2+}$) Conditions: Must be carried out in acidic medium (dilute $H_2SO_4$)

Procedure:

1. Fill burette with standardized $KMnO_4$ solution
2. Pipette known volume of $Fe^{2+}$ solution into conical flask
3. Add dilute $H_2SO_4$ to acidify
4. Titrate until permanent pale pink color appears

2. Potassium Dichromate(VI) Titrations

Principle: $Cr_2O_7^{2-}$ oxidizes $Fe^{2+}$ to $Fe^{3+}$ in acidic medium Reaction: $Cr_2O_7^{2-}(aq) + 6Fe^{2+}(aq) + 14H^+(aq) \rightarrow 2Cr^{3+}(aq) + 6Fe^{3+}(aq) + 7H_2O(l)$

Indicator: Diphenylamine sulfonate (blue → purple) Conditions: Must be carried out in acidic medium

3. Iodine-Thiosulfate Titrations

Principle: $I_2$ oxidizes $S_2O_3^{2-}$ to $S_4O_6^{2-}$ Reaction: $I_2(aq) + 2S_2O_3^{2-}(aq) \rightarrow 2I^-(aq) + S_4O_6^{2-}(aq)$

Indicator: Starch (blue-black → colorless) Conditions: Carried out in neutral or slightly alkaline medium

Procedure:

1. Pipette known volume of $I_2$ solution into flask
2. Add starch indicator
3. Titrate with standardized $Na_2S_2O_3$ solution until colorless

---

Did You Know?

Redox titrations are among the most accurate analytical methods in chemistry. They can determine concentrations with precisions of up to 0.1%, making them essential for quality control in pharmaceuticals, environmental monitoring, and industrial processes.

6.2 Redox Indicators

Types of Redox Indicators

1. Self-Indicators

Change: Color change of the titrant itself Example: $KMnO_4$ (purple → colorless)

2. Specific Indicators

Change: Color change based on specific redox conditions Example: Starch + $I_2$ (colorless → blue-black)

3. Oxidation-Reduction Indicators

Change: Color change due to oxidation or reduction of the indicator Examples:

• Diphenylamine sulfonate: Colorless → purple
• Methylene blue: Blue → colorless
• Potassium ferricyanide: Yellow → brown

Selection of Redox Indicators

Criteria for selection:

• Sharp color change at equivalence point
• Color change should be easily detectable
• Indicator should not interfere with the reaction
• Should work under the conditions of the titration

Key Properties

Redox potential: Indicator should have oxidation potential close to the system Reversibility: Should be reversible for accurate endpoint detection Stability: Should be stable under reaction conditions

---

SPM Exam Tips

• For $KMnO_4$ titrations, remember to use dilute $H_2SO_4$, NOT $HCl$ ($Cl^-$ would be oxidized)
• Always use a white tile behind the flask for better color detection
• Starch should be added near the endpoint of iodine titrations to prevent adsorption
• Record the burette reading to 2 decimal places for accuracy
• Know the stoichiometry of each redox reaction for calculations

6.3 Redox Calculations

Mole Ratio Method

Step-by-step approach:

1. Write balanced redox equation
2. Determine mole ratio from coefficients
3. Calculate moles of known substance
4. Use mole ratio to find moles of unknown substance
5. Calculate concentration or quantity of unknown

Example Calculations

Example 1: $KMnO_4$ vs $Fe^{2+}$

Problem: 25.0 $cm^3$ of $Fe^{2+}$ solution requires 18.5 $cm^3$ of 0.02 mol $dm^{-3}$ $KMnO_4$ for complete oxidation. Calculate the concentration of $Fe^{2+}$.

Solution:

1. Balanced equation: $MnO_4^- + 5Fe^{2+} + 8H^+ \rightarrow Mn^{2+} + 5Fe^{3+} + 4H_2O$
2. Mole ratio: 1 mol $MnO_4^-$ : 5 mol $Fe^{2+}$
3. Moles $MnO_4^-$ = 0.0185 $dm^3$ × 0.02 mol $dm^{-3}$ = $3.70 \times 10^{-4}$ mol
4. Moles $Fe^{2+}$ = $5 \times 3.70 \times 10^{-4} = 1.85 \times 10^{-3}$ mol
5. [$Fe^{2+}$] = $1.85 \times 10^{-3}$ mol / 0.025 $dm^3$ = 0.074 mol $dm^{-3}$

Example 2: $I_2$ vs $S_2O_3^{2-}$

Problem: 10.0 $cm^3$ of $I_2$ solution requires 22.4 $cm^3$ of 0.1 mol $dm^{-3}$ $Na_2S_2O_3$ solution. Calculate the concentration of $I_2$.

Solution:

1. Balanced equation: $I_2 + 2S_2O_3^{2-} \rightarrow 2I^- + S_4O_6^{2-}$
2. Mole ratio: 1 mol $I_2$ : 2 mol $S_2O_3^{2-}$
3. Moles $S_2O_3^{2-}$ = 0.0224 $dm^3$ × 0.1 mol $dm^{-3}$ = $2.24 \times 10^{-3}$ mol
4. Moles $I_2$ = $2.24 \times 10^{-3} / 2 = 1.12 \times 10^{-3}$ mol
5. [$I_2$] = $1.12 \times 10^{-3}$ mol / 0.010 $dm^3$ = 0.112 mol $dm^{-3}$

Example 3: Percentage Purity

Problem: 1.245 g of impure iron ore requires 24.6 $cm^3$ of 0.015 mol $dm^{-3}$ $KMnO_4$ for complete oxidation. Calculate the percentage of iron in the ore.

Solution:

1. Balanced equation: $MnO_4^- + 5Fe^{2+} + 8H^+ \rightarrow Mn^{2+} + 5Fe^{3+} + 4H_2O$
2. Moles $MnO_4^-$ = 0.0246 $dm^3$ × 0.015 mol $dm^{-3}$ = $3.69 \times 10^{-4}$ mol
3. Moles $Fe^{2+}$ = $5 \times 3.69 \times 10^{-4} = 1.845 \times 10^{-3}$ mol
4. Mass Fe = $1.845 \times 10^{-3}$ mol × 56 g $mol^{-1}$ = 0.1033 g
5. Percentage Fe = (0.1033 / 1.245) × 100% = 8.30%

---

Safety Reminder

When performing redox titrations:

• Always wear eye protection and lab coats
• Handle acids and oxidizing agents with care
• $KMnO_4$ and $K_2Cr_2O_7$ are strong oxidizers - avoid contact with skin
• Iodine solutions can stain - handle carefully
• Properly dispose of chemical waste
• Use appropriate concentration to avoid violent reactions

6.4 Industrial Redox Applications

1. Water Treatment

Disinfection

Chlorination: $Cl_2 + H_2O \rightarrow HCl + HOCl$

• HOCl oxidizes organic matter and kills bacteria
• Effective against most pathogens

Ozonation: $O_3 + \text{organic matter} \rightarrow \text{oxidized products} + CO_2$

• Strong oxidant, no residual taste/odor
• Effective against chlorine-resistant pathogens

Removal of Heavy Metals

Reduction: $Cr^{6+} + 3Fe^{2+} \rightarrow Cr^{3+} + 3Fe^{3+}$

• Toxic $Cr^{6+}$ converted to less toxic $Cr^{3+}$
• Precipitated as $Cr(OH)_3$ and removed

2. Chemical Manufacturing

Contact Process (Sulfuric Acid)

Catalytic oxidation: $2SO_2 + O_2 \rightleftharpoons 2SO_3$

• Vanadium(V) oxide catalyst
• Exothermic reaction

Haber Process (Ammonia)

Oxidation: $4NH_3 + 5O_2 \rightarrow 4NO + 6H_2O$ (first step)

• Iron catalyst
• Nitric acid production

3. Battery Technology

Lead-Acid Battery

Discharge: $Pb + PbO_2 + 2H_2SO_4 \rightarrow 2PbSO_4 + 2H_2O$ Recharge: Reverse reaction

• Lead and lead dioxide electrodes
• Sulfuric acid electrolyte

Lithium-Ion Battery

Redox reactions: $Li^+$ intercalation in electrodes

• High energy density
• Rechargeable

4. Metal Extraction and Refining

Extraction of Aluminum

Electrolytic reduction: $Al^{3+} + 3e^- \rightarrow Al$

• Molten cryolite electrolyte
• Carbon electrodes

Electrorefining of Copper

Anode oxidation: $Cu \rightarrow Cu^{2+} + 2e^-$ Cathode reduction: $Cu^{2+} + 2e^- \rightarrow Cu$

• Purifies copper to 99.99%

6.5 Environmental Redox Applications

1. Air Pollution Control

Catalytic Converters

Oxidation: $2CO + O_2 \rightarrow 2CO_2$ Reduction: $2NO + 2CO \rightarrow N_2 + 2CO_2$

• Platinum/rhodium catalysts
• Converts harmful gases to less harmful products

2. Wastewater Treatment

Chemical Oxygen Demand (COD)

Oxidation of organic matter: $\text{Organic matter} + Cr_2O_7^{2-} + H^+ \rightarrow CO_2 + Cr^{3+}$

• Measures pollution level
• Uses potassium dichromate

Biological Oxygen Demand (BOD)

Microbial oxidation: $\text{Organic matter} + O_2 \rightarrow CO_2 + H_2O + \text{biomass}$

• Measures biodegradable organic matter
• Important for water quality assessment

3. Soil Remediation

Reductive dechlorination: $Cl-R + 2e^- + 2H^+ \rightarrow R-H + Cl^-$

• Removes chlorinated pollutants
• Uses reducing agents like zero-valent iron

---

Did You Know?

• The world's largest redox process is the production of sulfuric acid, with annual production exceeding 240 million tons
• Redox indicators like methylene blue were discovered in the 19th century and revolutionized analytical chemistry
• Modern fuel cells are essentially redox devices that convert chemical energy directly to electrical energy with high efficiency

6.6 Advanced Redox Stoichiometry

Complex Calculations

Example 4: Back Titration

Problem: Excess KI is added to 25.0 $cm^3$ of $Cu^{2+}$ solution. The liberated $I_2$ requires 18.3 $cm^3$ of 0.1 mol $dm^{-3}$ $Na_2S_2O_3$. Calculate [$Cu^{2+}$].

Solution:

1. $Cu^{2+}$ oxidizes $I^-$: $2Cu^{2+} + 4I^- \rightarrow 2CuI + I_2$
2. $I_2$ reacts with $S_2O_3^{2-}$: $I_2 + 2S_2O_3^{2-} \rightarrow 2I^- + S_4O_6^{2-}$
3. Moles $S_2O_3^{2-}$ = 0.0183 × 0.1 = $1.83 \times 10^{-3}$ mol
4. Moles $I_2$ = $1.83 \times 10^{-3} / 2 = 9.15 \times 10^{-4}$ mol
5. From reaction 1: 1 mol $I_2$ ≡ 2 mol $Cu^{2+}$
6. Moles $Cu^{2+}$ = $2 \times 9.15 \times 10^{-4} = 1.83 \times 10^{-3}$ mol
7. [$Cu^{2+}$] = $1.83 \times 10^{-3}$ / 0.025 = 0.0732 mol $dm^{-3}$

Example 5: Multiple Redox Reactions

Problem: A sample contains both $Fe^{2+}$ and $C_2O_4^{2-}$. When titrated with $KMnO_4$, $Fe^{2+}$ reacts first, then $C_2O_4^{2-}$ is oxidized after acidification. Calculate amounts of each ion.

Solution:

1. First endpoint ($Fe^{2+}$): $MnO_4^- + 5Fe^{2+} \rightarrow Mn^{2+} + 5Fe^{3+}$
2. Second endpoint ($C_2O_4^{2-}$): $2MnO_4^- + 5C_2O_4^{2-} + 16H^+ \rightarrow 2Mn^{2+} + 10CO_2 + 8H_2O$
3. Use first titration volume for $Fe^{2+}$ calculation
4. Use second titration volume for $C_2O_4^{2-}$ calculation

---

Summary

Key Concepts

1. Redox titrations are quantitative analysis methods using redox reactions
2. Redox indicators detect endpoints through color changes
3. Stoichiometric calculations use mole ratios from balanced equations
4. Industrial applications include water treatment, chemical manufacturing, and battery technology
5. Environmental applications include pollution control and remediation
6. Advanced calculations include back titrations and multiple redox systems

Analytical Methods

• $KMnO_4$ titrations: Self-indicator for $Fe^{2+}$ determination
• $K_2Cr_2O_7$ titrations: Specific indicators for $Fe^{2+}$ determination
• $I_2$-$S_2O_3$ titrations: Starch indicator for iodine determination
• Back titrations: For substances that react slowly or multiple components

Industrial Processes

• Water treatment: Chlorination, ozonation, heavy metal removal
• Chemical manufacturing: Contact process, Haber process
• Battery technology: Lead-acid, lithium-ion
• Metal extraction: Aluminum electrolysis, copper refining

Environmental Applications

• Air pollution: Catalytic converters
• Water treatment: COD, BOD measurement
• Soil remediation: Reductive dechlorination

Problem-Solving Strategy

1. Write balanced redox equations
2. Determine stoichiometric ratios
3. Calculate moles of known substance
4. Use ratios to find unknown quantities
5. Consider special cases (back titration, multiple components)

Practice Questions

1. Explain why $KMnO_4$ titrations must be carried out in acidic medium.

2. A 0.25 g sample of iron ore requires 24.8 $cm^3$ of 0.015 mol $dm^{-3}$ $KMnO_4$ for complete oxidation. Calculate the percentage of iron in the ore.

3. Describe how redox reactions are used in the treatment of wastewater and explain the significance of COD and BOD measurements.

4. Calculate the concentration of $Na_2S_2O_3$ solution if 25.0 $cm^3$ of it liberates $I_2$ from excess KI, and the liberated $I_2$ requires 18.5 $cm^3$ of 0.1 mol $dm^{-3}$ $KMnO_4$ for titration.

---

Related Topics:

• Chapter 1: Chemical Equilibrium
• Chapter 2: Rate of Reaction
• Chapter 3: Thermochemistry
• Chapter 4: Electrochemistry
• Chapter 5: Carbon Compounds
• Chapter 7: Manufactured Substances