Redox reactions notes

Definition of Key Terms:
- Oxidation: The process involving the addition of oxygen to a substance.
- Example: When lead oxide is combined with carbon, lead and carbon monoxide are produced, indicating oxygen removal from lead oxide.
- Reduction: The process involving the removal of oxygen from a substance.
- Example: In the same reaction, carbon monoxide indicates the addition of oxygen.
- Redox Reactions: Named for the simultaneous occurrence of reduction and oxidation.

Nature of Redox Reactions:
- In any redox reaction, one substance is oxidized while another is reduced.
- The general reaction can be summarized:
- Example Reaction: Lead oxide + Carbon → Lead + Carbon monoxide + Oxygen

Reaction Analysis:
- Example 1: Carbon dioxide + Copper → Copper oxide + Carbon
- Oxidized: Carbon
- Reduced: Copper oxide
- Example 2: Iron oxide + Aluminium → Aluminium oxide + Iron
- Oxidized: Aluminium
- Reduced: Iron oxide
- Example 3: Carbon monoxide + Iron oxide → Carbon dioxide + Iron
- Oxidized: Carbon monoxide
- Reduced: Iron oxide
- Example 4: Zinc oxide + Hydrogen → Zinc + Water
- Oxidized: Hydrogen
- Reduced: Zinc oxide
- Example 5: Magnesium + Oxygen → Magnesium oxide
- Oxidized: Magnesium
- Reduced: Oxygen

Electron Involvement:
- Redox reactions can also be explained in terms of electron transfer.
- In the combustion of magnesium:
- $2Mg(s) + O_2(g) ightarrow 2MgO(s)$
- Reactivity of Magnesium:
 - Magnesium is oxidized, losing electrons to form $Mg^{2+}$ ions:
 - $Mg ightarrow Mg^{2+} + 2e^-$
 - Oxygen is reduced, gaining electrons to form $O^{2-}$ ions:
 - $O_2 + 4e^- ightarrow 2O^{2-}$

OILRIG: A technique to remember oxidation and reduction:
- Oxidation Is Loss of electrons.
- Reduction Is Gain of electrons.

Defining Half-Equations:
- Half-equations illustrate the processes of oxidation and reduction:
- Example of examples:
 - Oxidation Half-Equation:
 - $Mg ightarrow Mg^{2+} + 2e^-$
 - Reduction Half-Equation:
 - $O_2 + 4e^- ightarrow 2O^{2-}$

In a redox reaction:
- One substance is reduced (gains electrons) and one substance is oxidized (loses electrons).
- Oxidation always co-occurs with reduction.
Oxidation Defined:
- Two Aspects:
- Personification Redox: The addition of oxygen.
- Electrons: The loss of electrons.
Reduction Defined:
- Two Aspects:
- Personification Redox: The loss of oxygen.
- Electrons: The gain of electrons.

Usage of Oxidation Numbers:
- Oxidation numbers help determine whether an element has been oxidized or reduced.
- The oxidation number indicates the number of electrons gained or lost when an atom turns into an ion.
Rules for Assigning Oxidation Numbers:
- The oxidation number of an element in its elemental form is zero.
- The oxidation number of an ion equals its charge.
- Hydrogen typically has an oxidation number of +1 (except in hydrides, where it is -1).
- Oxygen typically has an oxidation number of -2 (except in peroxides, where it is -1).
- Fluorine always has an oxidation number of -1.
- Balancing Oxidation Numbers:
- Sum of oxidation numbers in a neutral compound must equal zero or the total charge of the ion in a charged species.

Compounds for Practice:
1. $KMnO_4$ (Potassium permanganate)
2. $NiO_2$ (Nickel(II) peroxide)
3. $P4O6$ (Tetraphosphorus hexoxide)
4. $Fe3O4$ (Iron(II,III) oxide, magnetite)
5. $SF_4$ (Sulfur tetrafluoride)
6. $XeOF_4$ (Xenon oxytetrafluoride)
7. $CO$ (Carbon monoxide)
8. $Na2C2O_4$ (Sodium oxalate)
9. $As2O3$ (Arsenic(III) oxide)
10. $NaBiO_3$ (Sodium bismuthate)
11. $Mg2P2O_7$ (Magnesium pyrophosphate)
12. $Hg2Cl2$ (Mercury(I) chloride)

Definitions:
- Oxidizing Agent: A substance that induces oxidation in another substance while being reduced itself.
- Reducing Agent: A substance that induces reduction in another substance while being oxidized itself.
Example of Reducing Agent:
- Reaction:
- $CuO (s) + H2 (g) ightarrow Cu (s) + H2O (l)$
- Analysis:
- Copper(II) oxide is reduced to copper by hydrogen, which acts as the reducing agent.
- Hydrogen is oxidized to water.

Potassium Manganate (VII):
- Acts as a strong oxidizing agent in acidic conditions, reducing from MnO4- (purple) to Mn2+ (colorless).
- Utilized to test for reducing agents due to its color change.
Potassium Dichromate (VI):
- A strong oxidizing agent reacting to form lower oxidation states.
- Used in breathalyzer tests to oxidize ethanol to ethanal, showcasing a color change from orange dichromate to green chromium(III).

Example:
- Balanced Equation:
- $Zn + CuCl2 ightarrow ZnCl2 + Cu$
- Ionic Equation:
- $Zn + Cu^{2+} ightarrow Zn^{2+} + Cu$
Note: Chloride ions are spectator ions.

Process:
- Use oxidation numbers to identify whether redox processes involve oxidation, reduction, both or neither.
Examples:
- (a) F_2 + 2e^- → 2F^-
- (b) MnO2 + 2H2O → MnO4^- + 4H^+ + 3e^-
- (c) H2C2O4 → 2CO2 + 2H^+ + 2e^-
- (d) 2CrO4^{2-} + 2H^+ → Cr2O7^{2-} + H2O
- (e) C2H4 + Cl2 → C2H4Cl2
- (f) Zn → Zn^2+ + 2e^-
- (g) Fe^2+ + 2e^- → Fe

Steps:
1. Identify oxidation numbers for all atoms.
2. Write half-equations for oxidation and reduction.
3. Balance the half-equations by multiplying electrons.
4. Combine both half-equations for a complete redox equation.
5. Adjust stoichiometric coefficients for changing oxidation numbers.
6. Balance atoms that do not change oxidation numbers.
7. For ionic redox equations, adjust coefficients according to ion charge.
Example Balancing Problem:
- $H_2O2 + HIO3 ightarrow O2 + H2O$
Final Balanced Equation Example:
- $6FeCl2 + K2Cr2O7 + 14HCl ightarrow 6FeCl3 + 2CrCl3 + 2KCl + 7H_2O$
Homework Assignment:
- Balance the following reactions:
1. $Al + AgNO3 ightarrow Al(NO3)3 + Ag$
2. $HI + H2SO4 ightarrow I2 + H2S + H2O$
3. $Fe2O3 + CO ightarrow Fe + CO2$
4. $PbS + O2 ightarrow PbO + SO2$
5. $P + HNO3 + H2O ightarrow H3PO4 + NO$