Exhaustive Review of Oxidation-Reduction Reactions

Oxidation and Reduction Principles

• Definitions and Redox Reactions: Oxidation and reduction are considered opposite chemical processes that always occur simultaneously within specific chemical reactions. These combined processes are termed redox reactions.

• Electron Transfer Perspective: Redox reactions involve the movement of electrons between chemical species. Definitions based on this transfer are as follows:     * Oxidation: The loss of electrons by an element, whether it exists in its free state or within a compound.     * Reduction: The gain of electrons by an element, whether in its free state or within a compound.     * OIL RIG Mnemonic: A common tool to remember these definitions is OIL RIG: Oxidation Is Loss, Reduction Is Gain.

• Example: Aluminium and Chlorine:     * When solid aluminium is heated in chlorine gas, aluminium chloride forms: $2Al(s) + 3Cl_{2}(g) \rightarrow 2AlCl_{3}(s)$.     * Aluminium chloride is an ionic compound containing $Al^{3+}$ and $Cl^{-}$ ions.     * Oxidation Half-Reaction: Each aluminium atom loses three electrons: $Al \rightarrow Al^{3+} + 3e^{-}$. Overall: $2Al(s) \rightarrow 2Al^{3+}(s) + 6e^{-}$. Aluminium is oxidised.     * Reduction Half-Reaction: Each chlorine atom in the molecule gains one electron: $Cl + e^{-} \rightarrow Cl^{-}$. Overall: $3Cl_{2}(g) + 6e^{-} \rightarrow 6Cl^{-}(s)$. Chlorine is reduced.

Oxidation Numbers and Identification Rules

• Definition of Oxidation Number: An oxidation number (or oxidation state) is assigned to every atom or ion in a substance to indicate the number of electrons lost, gained, or shared during chemical bonding. These values are positive, negative, or zero. A sign ($+$ or $-$) and the number (including $1$) must always be written unless the value is zero (e.g., $+1$, $-3$).

• Step-by-Step Rules for Determining Oxidation Numbers:     * Rule 1: Atoms of an element in their free, uncombined state have an oxidation number of $0$ (e.g., $Mg = 0$, $H$ in $H_{2} = 0$).     * Rule 2: For monatomic ions in ionic compounds, the oxidation number equals the ion's charge (e.g., in $MgCl_{2}$, $Mg^{2+} = +2$ and $Cl^{-} = -1$).     * Rule 3: Hydrogen in compounds or polyatomic ions is always $+1$, except in metal hydrides where it is $-1$ (e.g., $+1$ in $H_{2}O$ and $NH_{4}^{+}$; $-1$ in $MgH_{2}$).     * Rule 4: Oxygen in compounds or polyatomic ions is always $-2$, except in peroxides where it is $-1$ (e.g., $-2$ in $MgO$ and $SO_{4}^{2-}$; $-1$ in $H_{2}O_{2}$).     * Rule 5: For other elements in covalent compounds and polyatomic ions, oxidation numbers may vary and are often indicated by the compound name (e.g., Sulfur in sulfur(VI) oxide ($SO_{3}$) is $+6$; in sulfate(IV) ion ($SO_{3}^{2-}$) it is $+4$).     * Rule 6: The sum of all oxidation numbers in a neutral compound must equal $0$ (e.g., in $MgCl_{2}$, $(+2) + 2(-1) = 0$).     * Rule 7: The sum of oxidation numbers in a polyatomic ion must equal the net charge of the ion (e.g., in $OH^{-}$, $(-2) + (+1) = -1$).

• Naming Conventions: When naming polyatomic ions based on oxidation numbers, the name ends in "-ate." For example, the $NO_{2}^{-}$ ion (Nitrogen is $+3$) is called the nitrate(III) ion.

Identifying Oxidation and Reduction via Oxidation Numbers

• Definitions:     * Oxidation: An increase in the oxidation number of an element.     * Reduction: A decrease in the oxidation number of an element.

• Procedure for Identification:     1. Write the balanced chemical equation.     2. Calculate the oxidation number (placed in brackets) for every element below the equation.     3. If an oxidation number increases, that reactant is oxidised.     4. If an oxidation number decreases, that reactant is reduced.     5. Note: If no oxidation numbers change, the reaction is not a redox reaction.

• Worked Examples:     * Nitrogen in $NO_{2}$: $N + 2(-2) = 0 \rightarrow N = +4$. Name: Nitrogen(IV) oxide.     * Chromium in $Cr_{2}O_{7}^{2-}$: $2Cr + 7(-2) = -2 \rightarrow 2Cr - 14 = -2 \rightarrow 2Cr = +12 \rightarrow Cr = +6$. Name: Dichromate(VI) ion.     * Iron(III) oxide and Carbon monoxide: $Fe_{2}O_{3}(s) (+3) + 3CO(g) (+2) \rightarrow 2Fe(s) (0) + 3CO_{2}(g) (+4)$. Iron decreases from $+3$ to $0$ (reduced); Carbon increases from $+2$ to $+4$ (oxidised).     * Magnesium and Sulfuric acid: $Mg(s) (0) + H_{2}SO_{4}(aq) (+1) \rightarrow MgSO_{4}(aq) (+2) + H_{2}(g) (0)$. Magnesium increases from $0$ to $+2$ (oxidised); Hydrogen decreases from $+1$ to $0$ (reduced).

Oxidising and Reducing Agents

• Definitions:     * Oxidising Agent: The reactant that causes another substance to be oxidised. It achieves this by taking electrons (gaining them) or increasing the other element's oxidation number. The agent itself is reduced.     * Reducing Agent: The reactant that causes another substance to be reduced. It achieves this by donating electrons (losing them) or decreasing the other element's oxidation number. The agent itself is oxidised.

• Electron Transfer Examples:     * Calcium and Oxygen: $2Ca(s) + O_{2}(g) \rightarrow 2CaO(s)$. Calcium atoms (reducing agent) lose electrons to oxygen. Oxygen (oxidising agent) gains electrons from calcium.

• Oxidation Number Examples:     * Chlorine and Hydrogen sulfide: $Cl_{2}(g) (0) + H_{2}S(g) (-2) \rightarrow 2HCl(g) (-1) + S(s) (0)$. Sulfur increases from $-2$ to $0$; $H_{2}S$ is the reducing agent. Chlorine decreases from $0$ to $-1$; $Cl_{2}$ is the oxidising agent.     * Zinc and Copper(II) sulfate: $Zn(s) (0) + CuSO_{4}(aq) (+2) \rightarrow ZnSO_{4}(aq) (+2) + Cu(s) (0)$. Zinc increases from $0$ to $+2$; Zinc is the reducing agent. Copper decreases from $+2$ to $0$; $CuSO_{4}$ is the oxidising agent.

Common Chemical Agents and Tests

• Common Oxidising Agents (Table 10.1):     * Acidified potassium manganate(VII) ($H^{+}/KMnO_{4}$): Changes from purple to colourless ($MnO_{4}^{-} \rightarrow Mn^{2+}$).     * Acidified potassium dichromate(VI) ($H^{+}/K_{2}Cr_{2}O_{7}$): Changes from orange to green ($Cr_{2}O_{7}^{2-} \rightarrow Cr^{3+}$).     * Aqueous iron(III) salts ($Fe^{3+}$): Changes from yellow-brown to pale green ($Fe^{3+} \rightarrow Fe^{2+}$).     * Sodium chlorate(I) ($NaClO$): Turns many coloured dyes colourless.     * Hot concentrated sulfuric acid ($H_{2}SO_{4}$): Produces pungent colorless sulfur dioxide ($SO_{2}$).     * Nitric acid ($HNO_{3}$): Produces brown nitrogen dioxide gas ($NO_{2}$).     * Others: Oxygen ($O_{2}$), Chlorine ($Cl_{2}$), Manganese(IV) oxide ($MnO_{2}$).

• Common Reducing Agents (Table 10.2):     * Potassium iodide solution ($KI$): Changes from colourless to brown (Iodine forms).     * Aqueous iron(II) salts ($Fe^{2+}$): Changes from pale green to yellow-brown ($Fe^{2+} \rightarrow Fe^{3+}$).     * Hydrogen sulfide gas ($H_{2}S$): Forms a yellow precipitate of solid sulfur ($S$).     * Concentrated hydrochloric acid ($HCl$): Produces yellow-green chlorine gas ($Cl_{2}$).     * Others: Hydrogen ($H_{2}$), Carbon ($C$), Carbon monoxide ($CO$), Reactive metals.

• Amphoteric Behavior (Both Agents):     * Acidified Hydrogen Peroxide ($H^{+}/H_{2}O_{2}$): Usually an oxidising agent (e.g., with $KI$). Acts as a reducing agent when paired with stronger oxidisers like $KMnO_{4}$ or $K_{2}Cr_{2}O_{7}$.     * Sulfur Dioxide ($SO_{2}$): Usually a reducing agent. Acts as an oxidising agent with stronger reducing agents like $H_{2}S$ (oxidising it to yellow sulfur).

Redox in Everyday Activities

• Bleaches: Sodium chlorate(I) and hydrogen peroxide oxidise dyes in stains into colourless forms.
• Rusting: Iron is oxidised by oxygen and moisture to form hydrated iron(III) oxide ($Fe_{2}O_{3} \cdot xH_{2}O$).
• Browning of Fruit: Enzymes in apples, bananas, or potatoes react with oxygen to oxidise chemicals into brown melanins.
• Preserving Food: Sodium sulfite ($Na_{2}SO_{3}$) and sulfur dioxide ($SO_{2}$) act as reducing agents. They prevent oxidation (e.g., wine to vinegar, loss of Vitamin C) and reduce melanins back to colourless forms.

Questions & Discussion

• Revision Question 1: Define oxidation, reduction, oxidising agent, and reducing agent in terms of electrons.
• Revision Question 2: Identify if $I_{2}(aq) + 2e^{-} \rightarrow 2I^{-}(aq)$ is oxidation or reduction (Reason: Reduction, as electrons are gained).
• Revision Question 3: Identify if $Cu^{+}(aq) \rightarrow Cu^{2+}(aq) + e^{-}$ is oxidation or reduction (Reason: Oxidation, as an electron is lost).
• Revision Question 4: Identify if $2Br^{-}(aq) \rightarrow Br_{2}(aq) + 2e^{-}$ is oxidation or reduction (Reason: Oxidation, as electrons are lost).
• Revision Question 5: Determine oxidation numbers for bromine in $BrO_{2}$, $BrO_{3}^{-}$, and $BrO^{-}$; nitrogen in $NH_{3}$, $NO_{2}^{-}$, and $N_{2}O$; carbon in $CO$, $CO_{3}^{2-}$, $CH_{4}$, $C_{3}H_{6}$, and $HCO_{3}^{-}$.
• Revision Question 6: Determine oxidation numbers of sulfur in $SO_{3}^{2-}$ and $SO_{4}^{2-}$ and name the ions.
• Revision Question 7: Suggest an alternative name for nitrogen monoxide ($NO$).
• Revision Question 8: Define oxidation, reduction, oxidising agent, and reducing agent in terms of oxidation number.
• Revision Question 9: For the reaction $2FeCl_{3}(aq) + H_{2}S(g) \rightarrow 2FeCl_{2}(aq) + S(s) + 2HCl(aq)$, state which reactant is oxidised and which is reduced.
• Revision Question 10: For the reaction $C_{2}H_{4}(g) + 4CuO(s) \rightarrow 4Cu(s) + CO_{2}(g) + 2H_{2}O(g)$, identify the oxidised and reduced reactants using oxidation numbers.
• Revision Question 11: Identify the oxidising and reducing agents in $Fe_{2}O_{3}(s) + 3H_{2}(g) \rightarrow 2Fe(s) + 3H_{2}O(l)$.
• Revision Question 12: Name TWO reagents to prove a substance is a reducing agent (Expected answers: Acidified $KMnO_{4}$ or Acidified $K_{2}Cr_{2}O_{7}$) and state the color change and underlying chemical conversion.
• Revision Question 13: Name TWO substances that can behave as both oxidising and reducing agents (Expected answers: $H_{2}O_{2}$ and $SO_{2}$).