Ch 07 - Redox Reactions11th Chemistry

Addition: Oxidation occurs when oxygen or any other electronegative element is added to a substance (e.g., $2Mg + O_2 \rightarrow 2MgO$).

Removal: Oxidation also occurs when hydrogen or any other electropositive element is removed from a substance (e.g., $2H_2S + O_2 \rightarrow 2S + 2H_2O$).

Removal: Reduction is the removal of oxygen or any other electronegative element from a substance (e.g., $2HgO \rightarrow 2Hg + O_2$).

Addition: Reduction also encompasses the addition of hydrogen or any other electropositive element to a substance (e.g., $CH_2=CH_2 + H_2 \rightarrow C_2H_6$).

Simultaneity: A redox reaction is a chemical process where oxidation and reduction occur simultaneously.

Whenever one interacting species is oxidized, another is inevitably reduced, representing an interconnected transfer of electrons or chemical nature.

Oxidation: The loss of electron(s) by any species during a chemical reaction.

Reduction: The gain of electron(s) by any species during a chemical reaction.

This can be visualized by splitting the overall redox reaction into two distinct half-reactions (the oxidation half and the reduction half).

Oxidising Agent (Oxidant): A species that accepts electron(s). By accepting electrons, it causes the oxidation of another species while it gets strictly reduced itself.

Reducing Agent (Reductant): A species that donates electron(s). By donating electrons, it causes the reduction of another species while it gets strictly oxidised itself.

The electron-releasing tendency (reducing activity) strictly follows the order: $Zn > Cu > Ag$.

Significance: Zinc can displace Copper from its salt solution ($Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu$), and Copper can displace Silver from its salt solution ($Cu + 2Ag^+ \rightarrow Cu^{2+} + 2Ag$).

Definition: It denotes the apparent oxidation state of an element in a compound, calculated based on a specific set of rules.

Core Assumption: It mathematically assumes that the shared electron pair in a covalent bond belongs entirely to the more electronegative element.

It serves as a highly practical bookkeeping method to track electron shifts during a reaction.

Standard: The oxidation number of O is mostly -2.

Exception 1 (Peroxides): In peroxides (e.g., $H_2O_2, Na_2O_2$), O is -1.

Exception 2 (Superoxides): In superoxides (e.g., $KO_2, RbO_2$), O is -1/2.

Exception 3 (Fluorides): When bonded to highly electronegative Fluorine, O is positive: +2 in $OF_2$ and +1 in $O_2F_2$.

Standard: The oxidation number of H is +1 when bonded to non-metals.

Exception (Metal Hydrides): When H is bonded to highly active, electropositive metals in binary compounds (e.g., $LiH, NaH, CaH_2$), its oxidation number is -1.

Fluorine ($F$): Because it is the most electronegative element in the periodic table, F has an oxidation number of strictly -1 in all its compounds.

Other Halogens ($Cl, Br, I$): They typically have an oxidation number of -1 (as halide ions).

Exception for Cl, Br, I: When combined with highly electronegative Oxygen (e.g., in oxoacids like $HClO_4$ and oxoanions like $ClO_3^-$), they exhibit positive oxidation numbers (+1, +3, +5, +7).

Alkali Metals (Group 1): Always possess an oxidation number of +1 (e.g., $Li, Na, K$).

Alkaline Earth Metals (Group 2): Always possess an oxidation number of +2 (e.g., $Mg, Ca, Ba$).

Aluminium ($Al$): Regarded to always have an oxidation number of +3 in all its compounds.

Method: The oxidation number is expressed by placing a Roman numeral in parentheses immediately after the symbol or name of the metal in the molecular formula.

Examples:

• Aurous chloride (Gold I): $Au(I)Cl$
• Auric chloride (Gold III): $Au(III)Cl_3$
• Stannous chloride (Tin II): $Sn(II)Cl_2$

Average Value: A fractional oxidation number is simply the mathematical average oxidation state of all atoms of that specific element within the compound.

Structural Reality: Structurally, the individual atoms of the element are present in different, whole-number oxidation states depending on their highly specific, localized bonding environments.

Terminal Bromines: The two terminal $Br$ atoms are bonded to three oxygen atoms each, giving them an oxidation state of +6.

Middle Bromine: The middle $Br$ atom is bonded to two oxygen atoms, giving it an oxidation state of +4.

Average Calculation: $(6 + 4 + 6) / 3 = +16/3$.

Terminal Sulphurs: The two extreme (terminal) $S$ atoms are heavily bonded to highly electronegative oxygens, giving each an oxidation state of +5.

Middle Sulphurs: The two central $S$ atoms are bonded solely to other Sulphur atoms ($S-S$ bonds), giving them an oxidation state of 0.

Average Calculation: $(5 + 0 + 0 + 5) / 4 = +2.5$.

Terminal Carbons: The two terminal carbon atoms are double-bonded to oxygen ($O=C=$), giving them an oxidation state of +2.

Middle Carbon: The central carbon atom is bonded only to other carbon atoms ($=C=$), giving it an oxidation state of 0.

Average Calculation: $(2 + 0 + 2) / 3 = +4/3$.

Nature of Mixed Oxides: They are not single homogeneous compounds but stoichiometric mixtures of two different oxides of the same metal existing in different, whole-number oxidation states.

Example ($Pb_3O_4$): Red lead is actually a mixture of $2PbO$ (where Lead is in the +2 state) and $PbO_2$ (where Lead is in the +4 state).

Average Calculation: $(2 + 2 + 4) / 3 = +8/3$. The fractional oxidation number is merely a mathematical average of these distinct whole-number states.

Combination: $A + B \rightarrow C$. At least one of the reactants ($A$ or $B$) must be in the elemental form (e.g., $C + O_2 \rightarrow CO_2$) for it to be a redox process.

Decomposition: $C \rightarrow A + B$. The breakdown of a compound into two or more components, where at least one product is in the elemental state (e.g., $2H_2O \rightarrow 2H_2 + O_2$).

Note: Not all decomposition reactions are redox (e.g., $CaCO_3 \rightarrow CaO + CO_2$ involves no change in oxidation numbers).

Metal Displacement: A metal in a compound is displaced by a more reactive (better reducing) metal in the uncombined state (e.g., $CuSO_4 + Zn \rightarrow Cu + ZnSO_4$).

Non-Metal Displacement: Mostly involves Hydrogen displacement (metals displacing $H_2$ from cold water, steam, or acids) or Halogen displacement (more reactive halogens displacing less reactive halide ions).

Variables & Units:

• H_2O_2 = Hydrogen peroxide (O.N. of Oxygen = -1)
• H_2O = Water (O.N. of Oxygen = -2)
• O_2 = Oxygen gas (O.N. of Oxygen = 0)

Conditions/Notes:

• The oxygen in peroxide (-1) splits into a lower state (-2 in water) and a higher state (0 in oxygen gas).

Very Active Metals ($Na, Ca, Ba$): Excellent reductants that can violently displace $H_2$ from cold water.

Less Active Metals ($Mg, Fe$): Displace $H_2$ only from high-energy steam.

Mildly Active Metals ($Cd, Sn$): Do not react with water/steam, but can displace $H_2$ from acids (e.g., $HCl$).

Least Active Metals ($Ag, Au$): Native state metals that cannot displace $H_2$ even from strong acids.

Variables & Units:

• Cl_2 = Chlorine gas (O.N. = 0)
• ClO^- = Hypochlorite ion (O.N. of Cl = +1)
• Cl^- = Chloride ion (O.N. of Cl = -1)

Conditions/Notes:

• Requires an alkaline/basic medium ($OH^-$).
• The hypochlorite ion acts as the active bleaching agent by oxidizing color-bearing stains.

Reactivity Order: $\mathbf{F_2 > Cl_2 > Br_2 > I_2}$.

Displacement Rule: A halogen higher in the group (stronger oxidant) can replace halide ions lower in the group from an aqueous solution.

Example: $Cl_2$ can displace $Br^-$ and $I^-$, but $I_2$ cannot displace $Cl^-$.

Principle: A more reactive halogen acts as a stronger oxidant and displaces a less reactive halide from its aqueous solution (e.g., $Cl_2$ displaces $Br^-$ and $I^-$).

Observation: The displaced $Br_2$ and $I_2$ molecules are non-polar and preferentially dissolve in an organic non-polar solvent layer (like $CCl_4$ or $CHCl_3$), where they exhibit distinct characteristic colors (Orange/Brown for $Br_2$, Violet/Purple for $I_2$) allowing for identification.

Equations: $Cl_2 + 2Br^- \rightarrow 2Cl^- + Br_2$ and $Cl_2 + 2I^- \rightarrow 2Cl^- + I_2$.

Definition: A highly specific type of redox reaction where an element in a given oxidation state is simultaneously oxidized and reduced.

Condition: The element undergoing disproportionation must be structurally able to exist in at least three different oxidation states (it starts in an intermediate state and converts into a higher state and a lower state simultaneously).

Variables & Units:

• H^+ = Hydrogen ion (Concentration = 1 M)
• e^- = Electron
• H_2 = Hydrogen gas (Pressure = 1 bar)

Conditions/Notes:

• The Standard Electrode Potential ($E^{\ominus}$) for the SHE is assigned a conventional value of 0.00 V at 298 K.
• It serves as the baseline to measure the standard electrode potentials of all other redox couples.

Disproportionation strictly requires an element to shift to both a lower and a higher (positive) oxidation state.

Since Fluorine is the most electronegative element in the entire periodic table, it can only exhibit the -1 oxidation state in its compounds.

Because it cannot be oxidized to a positive oxidation state (it can never lose electrons to another element), disproportionation is structurally impossible.

Variables & Units:

• F_2 = Fluorine gas (O.N. = 0)
• H_2O = Water (O.N. of O = -2)
• O_2 = Oxygen gas (O.N. = 0)
• HF = Hydrofluoric acid (O.N. of F = -1)

Conditions/Notes:

• Because Fluorine is the strongest oxidizing agent, it displaces oxygen from water.
• This reaction makes aqueous halogen displacement reactions impossible to carry out with $F_2$.

1. Oxidation Number Method: Balances equations by calculating the total increase and decrease in oxidation numbers of the reacting species and cross-multiplying the compounds by stoichiometric coefficients to make these electron transfers mathematically equal.

2. Half-Reaction (Ion-Electron) Method: Splits the overall skeleton reaction into two separate half-reactions (the oxidation half and the reduction half), balances them individually for mass and charge (by adding electrons, $H^+$, $OH^-$, or $H_2O$), and then algebraically recombines them.

1. Self-Indicators: Highly colored reagents like $KMnO_4$ act as their own indicator. The end-point is sharply signaled by the first lasting tinge of pale pink when all the colourless reductant is entirely consumed.

2. External/Internal Indicators: When reagents lack dramatic colour changes (e.g., $K_2Cr_2O_7$), specific indicators like diphenylamine are introduced. It turns intense blue immediately after the equivalence point is crossed.

3. Iodometric Titrations: Reagents that oxidise $I^-$ to $I_2$ rely on adding a starch indicator. Starch forms an intense blue complex with molecular $I_2$, which sharply disappears exactly when all $I_2$ is consumed by the standard thiosulphate titrant.

Variables & Units:

• P_4 = White Phosphorus (O.N. = 0)
• PH_3 = Phosphine gas (O.N. of P = -3)
• H_2PO_2^- = Hypophosphite ion (O.N. of P = +1)

Conditions/Notes:

• The reaction requires a basic/alkaline medium.
• Phosphorus is simultaneously oxidized to an oxidation state of +1 and reduced to an oxidation state of -3.

Variables & Units:

• S_8 = Elemental Sulphur (O.N. = 0)
• S^{2-} = Sulphide ion (O.N. of S = -2)
• S_2O_3^{2-} = Thiosulphate ion (O.N. of S = +2)

Conditions/Notes:

• The reaction requires a basic/alkaline medium.
• Sulphur is simultaneously oxidized to an oxidation state of +2 and reduced to an oxidation state of -2.

Redox Couple: A pair representing the oxidised and reduced forms of a substance participating in a half-reaction (e.g., $Zn^{2+}/Zn$).

Daniell Cell Setup: The redox reaction ($Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$) is executed indirectly.

Mechanism: A Zinc rod dips in $ZnSO_4$ solution (Anode/Oxidation) and a Copper rod dips in $CuSO_4$ solution (Cathode/Reduction). Electrons travel through an external metallic wire, generating a current, while a salt bridge completes the inner ionic circuit to maintain electrical neutrality.

Variables & Units:

• I_2 = Molecular Iodine (O.N. = 0)
• S_2O_3^{2-} = Thiosulphate ion (O.N. of S = +2)
• I^- = Iodide ion (O.N. = -1)
• S_4O_6^{2-} = Tetrathionate ion (O.N. of S = +2.5 avg)

Conditions/Notes:

• Starch is used as a specific indicator in this titration. It forms an intense blue complex with $I_2$ that sharply disappears exactly at the end-point when all $I_2$ is consumed.

Variables & Units:

• E^{\ominus}_{cell} = Standard cell potential (Volts, V)
• E^{\ominus}_{cathode} = Standard reduction potential of the right half-cell (V)
• E^{\ominus}_{anode} = Standard reduction potential of the left half-cell (V)

Conditions/Notes:

• By IUPAC convention, standard potentials are always written and used as Standard Reduction Potentials.
• Often written as $E^{\ominus}_{cell} = E^{\ominus}_{R} - E^{\ominus}_{L}$ (Right - Left).

Negative $E^{\ominus}$: Indicates that the redox couple is a stronger reducing agent than the $H^+/H_2$ couple. It readily loses electrons (undergoes oxidation) when coupled with the SHE.

Positive $E^{\ominus}$: Indicates that the redox couple is a weaker reducing agent than the $H^+/H_2$ couple. Instead, its oxidized form acts as a stronger oxidising agent that prefers to gain electrons (undergoes reduction) when coupled with the SHE.