Comprehensive Study Notes on Redox Reactions

Introduction to Redox Reactions

• Matter and change are central to chemistry; transformations of matter between various kinds occur through a variety of reactions.
• Redox reactions represent an essential category found across physical and biological phenomena.
• Technical and industrial applications include:
  • Pharmaceuticals and biological processes.
  • Metallurgical extraction of reactive metals and non-metals.
  • Agricultural processes.
  • Combustion of fuels (wood, coal, gas) for domestic and commercial energy.
  • Electrochemical processes for manufacturing compounds like caustic soda ($NaOH$).
  • Operations of dry and wet batteries.
  • Prevention and study of metal corrosion.
• Emerging environmental issues categorized under redox phenomena include the Hydrogen Economy (utilizing liquid hydrogen as fuel) and the development of the "Ozone Hole."

Classical Concepts of Oxidation and Reduction

• Originally, oxidation was defined as the addition of oxygen to an element or compound. This occurs naturally as dioxygen ($O_2$) makes up approximately $20\%$ of the atmosphere.
  • Example: $2Mg(s) + O_2(g) \rightarrow 2MgO(s)$
  • Example: $S(s) + O_2(g) \rightarrow SO_2(g)$
• The definition of oxidation was later broadened to include the removal of hydrogen.
  • Example: $CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(l)$ (Hydrogen in methane is replaced by oxygen).
  • Example: $2H_2S(g) + O_2(g) \rightarrow 2S(s) + 2H_2O(l)$
• Further expansion defined oxidation as the addition of any electronegative element or the removal of any electropositive element.
  • Addition of electronegative elements (fluorine, chlorine, sulphur):
    • $Mg(s) + F_2(g) \rightarrow MgF_2(s)$
    • $Mg(s) + Cl_2(g) \rightarrow MgCl_2(s)$
    • $Mg(s) + S(s) \rightarrow MgS(s)$
  • Removal of electropositive element (potassium):
    • $2K_4[Fe(CN)_6](aq) + H_2O_2(aq) \rightarrow 2K_3[Fe(CN)_6](aq) + 2KOH(aq)$
• Reduction was initially defined as the removal of oxygen from a compound. Current definitions include the removal of oxygen or an electronegative element, or the addition of hydrogen or an electropositive element.
  • Removal of oxygen: $2HgO(s) \rightarrow 2Hg(l) + O_2(g)$
  • Removal of electronegative element (chlorine): $2FeCl_3(aq) + H_2(g) \rightarrow 2FeCl_2(aq) + 2HCl(aq)$
  • Addition of hydrogen: $CH_2=CH_2(g) + H_2(g) \rightarrow H_3C - CH_3(g)$
  • Addition of electropositive element (mercury): $2HgCl_2(aq) + SnCl_2(aq) \rightarrow Hg_2Cl_2(s) + SnCl_4(aq)$
• Simultaneity: Oxidation and reduction always occur together. The term "redox" was coined to describe this combined process.

Electronic Interpretation of Redox Reactions

• Redox reactions can be understood as processes involving electron transfer between reacting species.
• Creation of ionic compounds illustrates this:
  • $2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)$ can be written as $2Na(s) + Cl_2(g) \rightarrow 2Na^+Cl^-(s)$
• Half-reactions explicitly show the loss or gain of electrons:
  • Oxidation half-reaction (loss of electrons): $2Na(s) \rightarrow 2Na^+(g) + 2e^-$
  • Reduction half-reaction (gain of electrons): $Cl_2(g) + 2e^- \rightarrow 2Cl^-(g)$
• Summary of Electronic Definitions:
  • Oxidation: Loss of electron(s) by any species.
  • Reduction: Gain of electron(s) by any species.
  • Oxidising agent: Acceptor of electron(s).
  • Reducing agent: Donor of electron(s).

Competitive Electron Transfer and Activity Series

• When a zinc ($Zn$) strip is placed in a blue copper nitrate ($Cu(NO_3)_2$) solution, the blue color disappears, and a reddish copper coating forms on the zinc.
  • Reaction: $Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$
  • Zinc releases electrons (oxidized), and copper ions accept them (reduced).
• Reversing the experiment by placing a copper strip in a zinc sulphate solution results in no reaction, indicating copper cannot displace zinc from its ions.
• Experiments with copper and silver nitrate ($AgNO_3$) show copper displacing silver:
  • Reaction: $Cu(s) + 2Ag^+(aq) \rightarrow Cu^{2+}(aq) + 2Ag(s)$
• Comparing these results establishes an order of electron-releasing tendency, known as a metal activity series or electrochemical series: $Zn > Cu > Ag$.
• This competition is the basis for Galvanic cells, which convert chemical energy into electrical energy.

Oxidation Number Concept

• The concept of oxidation number was developed to track electron shifts in covalent compounds where charge transfer is only partial (an electron shift rather than complete transfer).
• Oxidation number denotes the oxidation state of an element in a compound based on a set of rules where shared electron pairs in a covalent bond belong entirely to the more electronegative element.
• Rule-based Calculation for Oxidation Numbers:
  1. Free/Uncombined State: Each atom bears an oxidation number of zero ($H_2, O_2, Cl_2, O_3, P_4, S_8, Na, Mg, Al$).
  2. Monoatomic Ions: Oxidation number equals the charge on the ion ($Na^+ = +1$, $Mg^{2+} = +2$, $Fe^{3+} = +3$, $Cl^- = -1$, $O^{2-} = -2$).
  3. Alkali Metals: Always $+1$ in their compounds.
  4. Alkaline Earth Metals: Always $+2$ in their compounds.
  5. Aluminium: Always $+3$ in all compounds.
  6. Oxygen: Usually $-2$. Exceptions include:
     • Peroxides ($H_2O_2, Na_2O_2$): $-1$ per oxygen atom.
     • Superoxides ($KO_2, RbO_2$): $-1/2$ per oxygen atom.
     • When bonded to fluorine: In oxygen difluoride ($OF_2$), oxygen is $+2$. In dioxygen difluoride ($O_2F_2$), oxygen is $+1$.
  7. Hydrogen: Usually $+1$. Exception: Binary metal hydrides ($LiH, NaH, CaH_2$) where it is $-1$.
  8. Fluorine: Always $-1$ in all compounds.
  9. Halogens ($Cl, Br, I$): $-1$ as halides; however, when combined with oxygen (oxoacids/oxoanions), they exhibit positive oxidation numbers.
  10. Algebraic Sum: Sum of oxidation numbers in a neutral compound must be zero ($0$). In a polyatomic ion, the sum must equal the charge on the ion.

Periodic Trends in Oxidation States

• Transition elements usually display several positive oxidation states.
• Highest oxidation number for representative elements:
  • Groups 1 and 2: Group number.
  • Groups 13-17: Group number minus 10.
• Example of third-period compounds showing max oxidation states:
  • Group 1 ($Na$): $NaCl$, $+1$
  • Group 2 ($Mg$): $MgSO_4$, $+2$
  • Group 13 ($Al$): $AlF_3$, $+3$
  • Group 14 ($Si$): $SiCl_4$, $+4$
  • Group 15 ($P$): $P_4O_{10}$, $+5$
  • Group 16 ($S$): $SF_6$, $+6$
  • Group 17 ($Cl$): $HClO_4$, $+7$

Stock Notation

• Proposed by Alfred Stock, this notation signifies the oxidation state of a metal in its formula using Roman numerals in parentheses.
• Examples:
  • Aurous chloride: $Au(I)Cl$; Auric chloride: $Au(III)Cl_3$.
  • Stannous chloride: $Sn(II)Cl_2$; Stannic chloride: $Sn(IV)Cl_4$.
  • Reduced and oxidized forms can be identified via Stock notation: $Hg_2(I)Cl_2$ is the reduced form of $Hg(II)Cl_2$.

Redefined Redox Terms via Oxidation Number

• Oxidation: An increase in the oxidation number of an element.
• Reduction: A decrease in the oxidation number of an element.
• Oxidising agent (Oxidant): A reagent that increases the oxidation number of an element in another substance.
• Reducing agent (Reductant): A reagent that lowers the oxidation number of an element in another substance.
• Redox reactions: Reactions involving changes in oxidation numbers.

Types of Redox Reactions

1. Combination Reactions: $A + B \rightarrow C$. At least one reactant must be in elemental form.
   • $C(s) + O_2(g) \rightarrow CO_2(g)$
2. Decomposition Reactions: Breakdown of a compound into multiple products, where at least one is in elemental state.
   • $2H_2O(l) \rightarrow 2H_2(g) + O_2(g)$
   • Note: Not all decomposition is redox (e.g., $CaCO_3 \rightarrow CaO + CO_2$ is not redox).
3. Displacement Reactions: $X + YZ \rightarrow XZ + Y$.
   • Metal Displacement: High activity metals displace lower activity ones (used in metallurgy).
     • Example: $TiCl_4(l) + 2Mg(s) \rightarrow Ti(s) + 2MgCl_2(s)$.
   • Non-metal Displacement: Mainly involves Hydrogen displacement.
     • Active metals ($Na, Ca$) displace hydrogen from cold water.
     • Less active metals ($Mg, Fe$) displace hydrogen from steam.
     • Many metals ($Zn, Sn$) displace hydrogen from acids.
   • Halogen Displacement: Oxidising power decreases down Group 17: $F_2 > Cl_2 > Br_2 > I_2$.
     • $Cl_2(g) + 2KBr(aq) \rightarrow 2KCl(aq) + Br_2(l)$.
     • Fluorine is so reactive it displaces oxygen from water: $2H_2O(l) + 2F_2(g) \rightarrow 4HF(aq) + O_2(g)$.
4. Disproportionation Reactions: A single element in a specific oxidation state is simultaneously oxidized and reduced.
   • The element must exist in at least three oxidation states.
   • Example: Decomposition of hydrogen peroxide: $2H_2O_2(aq) \rightarrow 2H_2O(l) + O_2(g)$.
   • Example: Bleach formation: $Cl_2(g) + 2OH^-(aq) \rightarrow ClO^-(aq) + Cl^-(aq) + H_2O(l)$.

The Paradox of Fractional Oxidation Numbers

• Fractional oxidation numbers ($C_3O_2$ as $+4/3$, $Br_3O_8$ as $+16/3$, $Na_2S_4O_6$ as $+2.5$) appear because they represent an average oxidation state.
• Structural analysis reveals discrete integer states:
  • Carbon suboxide ($C_3O_2$): Structural formula $O=C=C=C=O$. Outer carbons are $+2$, center carbon is $0$.
  • Tribromooctaoxide ($Br_3O_8$): Terminal bromines are $+6$, middle is $+4$.
  • Tetrathionate ion ($S_4O_6^{2-}$): Two extreme sulphurs are $+5$, two middle sulphurs are $0$.

Balancing Redox Reactions

• Oxidation Number Method:
  1. Identify atoms changing oxidation numbers.
  2. Balance the increase and decrease in oxidation number by multiplying by appropriate coefficients.
  3. Balance ionic charges by adding $H^+$ (acidic medium) or $OH^-$ (basic medium).
  4. Balance hydrogen and oxygen by adding $H_2O$.
• Half-Reaction Method (Ion-Electron Method):
  1. Split the skeletal equation into oxidation and reduction halves.
  2. Balance all atoms except $H$ and $O$.
  3. In acidic medium: add $H_2O$ for $O$ and $H^+$ for $H$. In basic medium: for every $H^+$ added, add equal $OH^-$ to both sides.
  4. Add electrons to balance charges.
  5. Multiply half-reactions so electron count matches, then add and cancel electrons.

Redox Titrations

• Indicators signify the equivalence point where reductant and oxidant stoichiometry are matched.
• Self-Indicator: Permanganate ($MnO_4^-$). A persistent pink tinge appears at concentrations as low as $10^{-6}\,mol\,dm^{-3}$.
• External Indicator: Diphenylamine is used for dichromate ($Cr_2O_7^{2-}$) titrations, producing an intense blue color.
• Iodometric Method: Starch acts as an indicator. Iodine produces a blue color with starch, which disappears when iodine is reduced by thiosulphate ($S_2O_3^{2-}$).

Electrode Processes and Daniell Cell

• A Daniell cell physically separates the zinc oxidation half-reaction from the copper reduction half-reaction.
• Components:
  • Zinc strip in zinc sulphate solution (Anode/Oxidation).
  • Copper strip in copper sulphate solution (Cathode/Reduction).
  • Salt Bridge: A U-tube with $KCl$ or $NH_4NO_3$ in agar-agar jelly provides electrical contact without mixing solutions.
• Current flows from cathode to anode; electron flow is from anode to cathode.
• Standard Electrode Potential ($E^0$):
  • Potential associated with an electrode at $298\,K$ and $1\,M$ concentration ($1\,atm$ for gases).
  • Standard Hydrogen Electrode is defined as $0.00\,V$.
  • Negative $E^0$: Redox couple is a stronger reducing agent than $H^+/H_2$. Example: $Li^+/Li$ ($-3.05\,V$).
  • Positive $E^0$: Redox couple is a weaker reducing agent than $H^+/H_2$. Example: $F_2/F^-$ ($+2.87\,V$).

Questions & Discussion

• Question: Why isalcoholic potassium permanganate used to manufacture benzoic acid from toluene?
• Answer: Toluene and potassium permanganate can be mixed more effectively in an alcoholic medium, facilitating the contact between the organic reactant and the oxidant.
• Question: Why does concentrated sulphuric acid reacting with chloride give $HCl$, but with bromide give $Br_2$?
• Answer: Sulphuric acid is an oxidant. It cannot oxidize $Cl^-$ but is strong enough to oxidize $Br^-$ to $Br_2$.
• Question: Is the reaction $H_2O(s) + F_2(g) \rightarrow HF(g) + HOF(g)$ a redox reaction?
• Answer: Yes. Fluorine is reduced from $0$ to $-1$, and oxygen in water is oxidized from $-2$ to $0$ in $HOF$.