10.03 Oxidation Numbers
1.1 Introduction to Oxidation Numbers
Definition:
- Oxidation Number (Oxidation State): A numerical value assigned to an element in a compound to indicate the degree of oxidation (loss of electrons) or reduction (gain of electrons) of that element.
Purpose:
- Helps in determining the electron transfer in redox reactions.
- Essential for naming compounds, especially those involving transition metals with variable valencies.
Key Concepts:
- Transition Metals: Elements from the central region of the Periodic Table (Groups 3-12). They are typically hard, strong, dense metals that form colored compounds and exhibit multiple oxidation states.
- Variable Valency: Transition metals can form ions with different charges (oxidation states) in different compounds.
1.2 Understanding Oxidation Numbers through Examples
Examples of Compounds with Variable Oxidation Numbers:
| Compound | Oxidation Number of Metal |
|---|---|
| Copper(II) Sulfate (CuSO4 ) | +2 |
| Iron(III) Oxide (Fe2O3) | +3 |
| Cobalt(II) Chloride (CoCl2) | +2 |
| Iron(II) Chloride (FeCl2) | +2 |
| Iron(III) Chloride (FeCl3) | +3 |
| Manganese(II) Oxide (MnO) | +2 |
| Manganese(IV) Oxide (MnO2) | +4 |
| Manganese(VII) Oxide (Mn2O7) | +7 |
Explanation:
- Copper in CuSO4 : Present as Cu2+ ion (oxidation number +2).
- Iron in Fe2O3: Present as Fe3+ ion (oxidation number +3).
- Cobalt in CoCl2: Present as Co2+ ion (oxidation number +2).
Table 10.1: Variable Oxidation Numbers of Transition Metals
| Oxidation Number | Example Compound | Ion Form |
|---|---|---|
| +1 | Copper(I) Oxide (Cu2O) | Cu+ |
| +2 | Copper(II) Oxide (CuO) | Cu2+ |
| +2 | Iron(II) Chloride (FeCl2 ) | Fe2+ |
| +3 | Iron(III) Chloride (FeCl3) | Fe3+ |
| +2 | Manganese(II) Oxide (MnO) | Mn2+ |
| +4 | Manganese(IV) Oxide (MnO2) | Mn4+ |
| +7 | Manganese(VII) Oxide (Mn2O7 ) | Mn7+ |
1.3 Rules for Determining Oxidation Numbers
Table: Rules for Assigning Oxidation Numbers
| Rule | Explanation |
|---|---|
| 1. Free Elements | The oxidation number of an uncombined element is 0. |
| Examples: | H2 Cl2 O2 Zn |
| 2. Monoatomic Ions | The oxidation number of a monoatomic ion is equal to the charge on the ion. |
| Examples: | Zn2+ is +2, O2− is -2 |
| 3. Hydrogen (H) | +1 when bonded to non-metals; -1 when bonded to metals. |
| 4. Oxygen (O) | -2 in most compounds; exceptions include peroxides (e.g., H2O2) where it is -1. |
| 5. Fluorine (F) | Always -1 in its compounds. |
| 6. Other Halogens (Cl, Br, I) | Usually -1, unless they are combined with oxygen or other halogens in which case their oxidation numbers can vary. |
| 7. The Sum of Oxidation Numbers | The sum must equal the overall charge of the compound or ion. |
Detailed Explanation with Examples:
- Free Elements:
- In H2, each hydrogen has an oxidation number of 0.
- In Cl2, each chlorine atom has an oxidation number of 0.
- Monoatomic Ions:
- In Zn2+, zinc has an oxidation number of +2.
- In O2−, oxygen has an oxidation number of -2.
- Hydrogen in Compounds:
- In HCl, hydrogen has an oxidation number of +1.
- In metal hydrides like NaH, hydrogen has an oxidation number of -1.
- Oxygen in Compounds:
- In H2O, oxygen has an oxidation number of -2.
- In H2O2 (hydrogen peroxide), each oxygen has an oxidation number of -1.
- Fluorine in Compounds:
- In NaF\text{NaF}NaF, fluorine has an oxidation number of -1.
- Other Halogens:
- In ClO−, chlorine has an oxidation number of +1.
- In Cl2O7, chlorine has an oxidation number of +7.
- Sum of Oxidation Numbers:
- In H2O, 2(+1)+(−2)=0
- In SO42− = S + 4(-2) = −2, so S=+6
1.4 Using Oxidation Numbers to Identify Redox Reactions
Key Principle:
- Change in Oxidation Numbers: If the oxidation number of an element increases, it is oxidized; if it decreases, it is reduced.
Example 1: Metal Displacement Reaction
Zn(s)+CuSO4 (aq)→ZnSO4 (aq)+Cu(s)
Ionic Equation:
Zn(s)+Cu2+(aq)→Zn2+(aq)+Cu(s)
Oxidation Numbers:
| Element | Reactants Ox. No. | Products Ox. No. |
|---|---|---|
| Zn | 0 | +2 |
| Cu | +2 | 0 |
Analysis:
- Zinc (Zn): Oxidation number increases from 0 to +2 (oxidized).
- Copper (Cu): Oxidation number decreases from +2 to 0 (reduced).
- Conclusion: This is a redox reaction as zinc is oxidized and copper is reduced.
- Example 2: Halogen Displacement Reaction
Cl2(aq) + 2KI(aq) → 2KCl(aq) + I2(aq) - Ionic Equation:
Cl2(aq) + 2I−(aq) → 2Cl−(aq) + I2(aq)
Oxidation Numbers:
| Element | Reactants Ox. No. | Products Ox. No. |
|---|---|---|
| Cl | 0 | -1 |
| I | -1 | 0 |
Analysis:
- Chlorine (Cl): Oxidation number decreases from 0 to -1 (reduced).
- Iodine (I): Oxidation number increases from -1 to 0 (oxidized).
- Conclusion: This is a redox reaction as chlorine is reduced and iodide ions are oxidized.
1.5 Summary of Oxidation Numbers and Redox Reactions
- Oxidation Number: Indicates the oxidation state of an element in a compound.
- Determining Redox Reactions: Look for changes in oxidation numbers of elements between reactants and products.
- Redox Reactions: Always involve both oxidation and reduction processes.
2. Key Vocabulary
- Oxidation Number: A number assigned to an element in a compound indicating its oxidation state.
- Redox Reaction: A chemical reaction involving the transfer of electrons, resulting in the oxidation of one substance and the reduction of another.
- Oxidizing Agent: A substance that causes another to oxidize by accepting electrons; it is itself reduced.
- Reducing Agent: A substance that causes another to reduce by donating electrons; it is itself oxidized.
- Transition Metals: Elements in the central block of the Periodic Table (Groups 3-12) that exhibit variable oxidation states.
- Displacement Reaction: A reaction where a more reactive element displaces a less reactive element from its compound.
- Monoatomic Ion: An ion consisting of a single atom with a positive or negative charge.
- Half-Reaction: A part of a redox reaction that shows either oxidation or reduction separately.
4. Additional Key Concepts
4.1 Balancing Redox Equations
Methods:
- Half-Reaction Method:
- Step 1: Write the unbalanced equation.
- Step 2: Separate into oxidation and reduction half-reactions.
- Step 3: Balance all atoms except hydrogen and oxygen.
- Step 4: Balance oxygen atoms by adding H2O.
- Step 5: Balance hydrogen atoms by adding H+ (in acidic solutions) or OH− (in basic solutions).
- Step 6: Balance the charge by adding electrons (e−).
- Step 7: Multiply the half-reactions by appropriate coefficients to equalize electrons.
- Step 8: Add the half-reactions together and simplify.
- Ion-Electron Method:
- Focus on balancing the number of electrons lost and gained in the reaction.
Example:
- Balancing the reaction between Zn and CuSO4
- Zn(s) + CuSO4 (aq) → ZnSO4 (aq) + Cu(s)
Half-Reactions:
- Oxidation: Zn → Zn2++2e−
- Reduction: Cu2++2e− → Cu
Balanced Equation:Zn(s) + Cu2+(aq)→Zn2+(aq) + Cu(s)
4.2 Oxidizing and Reducing Agents in Detail
Reducing Agents:
- Function: Donate electrons to another substance, causing it to reduce.
- Common Examples:
- Hydrogen (H2): Used in hydrogenation reactions.
- Carbon (C): Used in the reduction of metal ores.
- Carbon Monoxide (CO): Acts as a reducing agent in the blast furnace.
Oxidizing Agents:
- Function: Accept electrons from another substance, causing it to oxidize.
- Common Examples:
- Oxygen (O2): In combustion and respiration.
- Hydrogen Peroxide (H2O2): Used in disinfection and bleaching.
- Potassium Permanganate (KMnO4 ): Used as an oxidizing agent in titrations.
- Potassium Dichromate (K2Cr2O7): Used in cleaning glassware.
Behavior in Reactions:
- Reducing Agent: Gets oxidized (loses electrons).
- Oxidizing Agent: Gets reduced (gains electrons).
4.3 Applications of Redox Reactions
- Metallurgy:
- Extraction of Metals: Reduction of metal oxides to obtain pure metals.
- Example: Extraction of iron from hematite (Fe2O3) in a blast furnace.
- Energy Production:
- Combustion Engines: Release energy through redox reactions.
- Batteries: Store and release energy via redox processes.
- Fuel Cells: Generate electricity through redox reactions (e.g., hydrogen-oxygen fuel cells).
- Biological Systems:
- Cellular Respiration: Combustion of glucose to release energy.
- Photosynthesis: Conversion of carbon dioxide and water into glucose and oxygen.
- Environmental Processes:
- Decomposition of Pollutants: Redox reactions break down harmful substances.
- Corrosion Control: Preventing oxidation of metals.
4.4 Electrolysis and Fuel Cells
Electrolysis:
- Definition: A chemical process that uses electricity to drive a non-spontaneous reaction, resulting in the decomposition of compounds.
- Key Concepts:
- Anode: Positive electrode where oxidation occurs.
- Cathode: Negative electrode where reduction occurs.
- Ion Movement:
- Negative Ions (Anions): Move to the anode to lose electrons (oxidation).
- Positive Ions (Cations): Move to the cathode to gain electrons (reduction).
- Example: Electrolysis of Concentrated Sodium Chloride Solution
- At the Cathode (Reduction): 2H+(aq) + 2e− → H2(g)
- At the Anode (Oxidation): 2Cl−(aq) → Cl2(g) + 2e−2
Fuel Cells:
- Definition: Devices that convert chemical energy from a fuel into electricity through a redox reaction.
- Common Fuel Cell: Hydrogen-Oxygen Fuel Cell.
- Reactions:
- At the Anode (Oxidation):
2H2(g)→4H+(aq)+4e−2 - At the Cathode (Reduction):
2O2(g) + 4H+(aq) + 4e− → 2H2O(l) - Overall Reaction:
2H2(g)+O2(g)→2H2O(l)+Energy
- At the Anode (Oxidation):
- Advantages:
- High Efficiency: More efficient than traditional combustion.
- Clean Energy: Produces only water as a by-product.
4.5 Summary of Oxidation Numbers and Redox Reactions
- Oxidation Numbers: Assigning oxidation states helps in identifying and balancing redox reactions.
- Redox Reactions: Involve simultaneous oxidation and reduction processes, essential in various chemical and biological systems.
- Applications: Wide-ranging, from industrial metal extraction to clean energy production in fuel cells.
Examples
Define the following terms, giving one example of each.
a. An Oxidising Agent
b. A Reducing Agent
a. An Oxidising Agent
Definition:
- A substance that oxidizes another substance by accepting electrons; it is itself reduced in the process.
Example:
- Chlorine Gas (Cl2):
- In the reaction Cl2 + 2I−→2Cl−+ I2, chlorine acts as an oxidizing agent by accepting electrons from iodide ions.
b. A Reducing Agent
Definition:
- A substance that reduces another substance by donating electrons; it is itself oxidized in the process.
Example:
- Zinc (Zn):
- In the reaction Zn + Cu2+ → Zn2++ Cu, zinc acts as a reducing agent by donating electrons to copper ions.
Complete the following statement:
________ is the gain of electrons; _______ is the loss of electrons. During a redox reaction, the oxidising agent _______ electrons; the oxidising agent is itself _______ during the reaction.
Answer:
- Reduction is the gain of electrons; oxidation is the loss of electrons. During a redox reaction, the oxidising agent accepts electrons; the oxidising agent is itself reduced during the reaction.
Redox means reduction and oxidation. It can be defined by loss and gain of oxygen or by loss and gain of electrons.
a. Which definition is more useful?
b. Is it possible to have oxidation without reduction in a chemical reaction? Explain your answer.
Redox means reduction and oxidation. It can be defined by loss and gain of oxygen or by loss and gain of electrons.
a. Which definition is more useful?
Answer:
- The electron transfer definition (oxidation is the loss of electrons; reduction is the gain of electrons) is more useful. This broader definition allows for the identification and analysis of a wider range of redox reactions, including those that do not involve oxygen.
b. Is it possible to have oxidation without reduction in a chemical reaction? Explain your answer.
Answer:
No, it is not possible to have oxidation without reduction in a chemical reaction. Oxidation and reduction always occur together because the electrons lost by one substance during oxidation must be gained by another substance during reduction. This simultaneous process ensures that both oxidation and reduction happen concurrently in redox reactions.
Quizzes
Test 1
Quiz
Your score: 0 / 10