Unit 5: Oxidation-Reduction and Principles of Metallurgy

Redox Equations and Electrode Potential

Redox Equations

Redox (Reduction-Oxidation) reactions involve the transfer of electrons.

• Oxidation: Loss of electrons (Oxidation Number increases).
• Reduction: Gain of electrons (Oxidation Number decreases).

Balancing redox equations is essential. The two common methods are:

1. Oxidation Number Method

1. Assign oxidation numbers (O.N.) to all atoms.
2. Identify atoms undergoing change in O.N.
3. Calculate the total increase in O.N. (oxidation) and total decrease in O.N. (reduction).
4. Use coefficients to equalize the total increase and decrease.
5. Balance all other atoms (except H and O).
6. Balance O by adding H₂O.
7. Balance H by adding H⁺ (in acidic medium) or OH^- (in basic medium).

2. Ion-Electron (Half-Reaction) Method

1. Split the reaction into two half-reactions (oxidation and reduction).
2. For each half-reaction:
   1. Balance atoms other than O and H.
   2. Balance O by adding H₂O.
   3. Balance H by adding H⁺ (for acidic medium).
   4. Balance charge by adding electrons (e^-).
3. Multiply the half-reactions by integers so that the number of electrons lost = number of electrons gained.
4. Add the two balanced half-reactions and cancel common terms.
5. (If in basic medium: add OH^- to both sides to neutralize any H⁺, forming H₂O).

Standard Electrode Potential (E^°)

This is a measure of the tendency of a species to be reduced (gain electrons). By convention, it is measured relative to the Standard Hydrogen Electrode (SHE), which is assigned a potential of E^° = 0.00 V.

Standard Conditions: 298 K (25°C), 1 M concentration for all ions, 1 atm pressure for all gases.

• High positive E^° (e.g., F₂/F^- = +2.87 V): Strong tendency to be reduced. Strong oxidizing agent.
• High negative E^° (e.g., Li⁺/Li = -3.05 V): Strong tendency to be oxidized. Strong reducing agent.

Application to Inorganic Reactions (Predicting Spontaneity)

We can predict if a redox reaction will be spontaneous by calculating the standard cell potential (E^°_cell).

Formula: E^°_cell = E^°_reduction (cathode) - E^°_reduction (anode)
(or E^°_cell = E^°_{species reduced} + E^°_{species oxidized})

The spontaneity is linked by the Gibbs free energy change (Δ G^°):

Formula: Δ G^° = -nFE^°_cell
Where: n = moles of electrons transferred, F = Faraday's constant (96,485 C/mol)

• If E^°_cell is positive → Δ G^° is negative → Reaction is spontaneous.
• If E^°_cell is negative → Δ G^° is positive → Reaction is non-spontaneous.

Example: Will Cu metal react with Zn²⁺ solution?
Reaction: Cu(s) + Zn²⁺(aq) → Cu²⁺(aq) + Zn(s)
(E^°_Cu²⁺/Cu = +0.34 V, E^°_Zn²⁺/Zn = -0.76 V)
Here, Cu is oxidized (anode) and Zn²⁺ is reduced (cathode).
E^°_cell = E^°_cathode - E^°_anode = (-0.76 V) - (+0.34 V) = -1.10 V.
Since E^°_cell is negative, the reaction is non-spontaneous. (The reverse reaction is spontaneous).

Principles of Volumetric Analysis (Redox Titrations)

This section covers the principles of two common redox titrations used to determine the concentration of an unknown solution.

Fe(II) and Oxalic Acid using standardized KMnO₄ solution

This is called Permanganometry. KMnO₄ (potassium permanganate) is a strong oxidizing agent and acts as its own self-indicator.

• Medium: Titration must be done in dilute H₂SO₄ (acidic medium).
  (Cannot use HCl: it gets oxidized to Cl₂. Cannot use HNO₃: it is an oxidizing agent itself.)
• Indicator: KMnO₄ is deep purple (MnO₄^-). After it is reduced, it forms Mn²⁺, which is colorless.
• Endpoint: The first persistent faint pink color in the flask (from one excess drop of KMnO₄).

Reactions:

Reduction: MnO₄^-(aq) + 8H⁺(aq) + 5e^- → Mn²⁺(aq) + 4H₂O(l) (x 2)
Oxidation (Fe II): Fe²⁺(aq) → Fe³⁺(aq) + e^- (x 10)
Overall: 2MnO₄^- + 10Fe²⁺ + 16H⁺ → 2Mn²⁺ + 10Fe³⁺ + 8H₂O
Molar Ratio: 2 moles KMnO₄ reacts with 10 moles Fe²⁺ (or 1:5).

Oxidation (Oxalic Acid): C₂O₄^2-(aq) → 2CO₂(g) + 2e^- (x 5)
Overall: 2MnO₄^- + 5C₂O₄^2- + 16H⁺ → 2Mn²⁺ + 10CO₂ + 8H₂O
Molar Ratio: 2 moles KMnO₄ reacts with 5 moles H₂C₂O₄ (2:5).
(Note: This titration must be heated to ~60°C as the reaction is slow to start).

Fe(II) with K₂Cr₂O₇ solution

This is called Dichrometry. K₂Cr₂O₇ (potassium dichromate) is a good primary standard (stable, pure).

• Medium: Also done in dilute H₂SO₄. (Here, HCl *can* be used, as K₂Cr₂O₇ is not strong enough to oxidize it).
• Indicator: K₂Cr₂O₇ is orange (Cr₂O₇^2-) and reduces to Cr³⁺ (green). This color change is not sharp, so a redox indicator like diphenylamine is used.
• Endpoint: Indicator changes from colorless to deep violet/blue.

Reaction:

Reduction: Cr₂O₇^2-(aq) + 14H⁺(aq) + 6e^- → 2Cr³⁺(aq) + 7H₂O(l) (x 1)
Oxidation (Fe II): Fe²⁺(aq) → Fe³⁺(aq) + e^- (x 6)
Overall: Cr₂O₇^2- + 6Fe²⁺ + 14H⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O
Molar Ratio: 1 mole K₂Cr₂O₇ reacts with 6 moles Fe²⁺ (1:6).

General Principles of Metallurgy

Metallurgy is the science of extracting metals from their ores and purifying them. The main steps are: (1) Concentration of ore, (2) Isolation of metal (reduction), (3) Purification (refining).

Chief Modes of Occurrence of Metals

The "native" or "combined" state of a metal depends on its reactivity, which is linked to its Standard Electrode Potential (E^°).

• Highly Negative E^° (Very Reactive): e.g., K, Na, Ca, Al.
  These are strong reducing agents, so they are easily oxidized. They are never found as native metals. They occur in stable, combined states (e.g., as chlorides, carbonates, silicates, oxides).
• Intermediate E^° (Moderately Reactive): e.g., Zn, Fe, Pb.
  These are found in combined states, typically as oxides or sulfides (e.g., ZnO, Fe₂O₃, PbS).
• Positive E^° (Least Reactive): e.g., Cu, Ag, Au, Pt.
  These are weak reducing agents. They are not easily oxidized and can be found in their native (elemental) state. They may also occur as less stable sulfides (e.g., Ag₂S).

Ellingham Diagrams (Thermodynamics of Reduction)

An Ellingham diagram plots the standard Gibbs free energy of formation of oxides (Δ G^°_f) versus Temperature (T).

Δ G^° = Δ H^° - TΔ S^°

• Key Features:
  1. Most lines for M(s) + O₂(g) → MO_x(s) have a positive slope.
     Reason: Δ S^° (entropy change) is negative (gas is consumed), so as T increases, the -TΔ S^° term becomes more positive, and Δ G^° becomes less negative (less stable).
  2. The line for C(s) + O₂(g) → CO₂(g) has a near-zero slope.
     Reason: Δ S^° ≈ 0 (1 mole gas → 1 mole gas).
  3. The line for C(s) + (1) / (2)O₂(g) → CO(g) has a negative slope.
     Reason: Δ S^° is positive (0.5 mole gas → 1 mole gas).
• Principle of Reduction: Any metal oxide MO can be reduced by a reducing agent (e.g., C) *if* the Δ G^° for the coupled reaction is negative.
  Rule of Thumb: On the diagram, a reducing agent (like C) can reduce a metal oxide (like FeO) at any temperature where the reducing agent's line is below the metal oxide's line.

Applications (using Carbon and CO)

• Reduction by Carbon (Smelting): The C → CO line slopes down and crosses below many metal oxide lines (like FeO, ZnO) at high temperatures. This is why C (coke) is used as a cheap reducing agent in blast furnaces.
  FeO(s) + C(s) → Fe(l) + CO(g) (at T > 1000 K)
• Reduction by Carbon Monoxide: The CO → CO₂ line is below the FeO line at lower temperatures. CO is the actual reducing agent in the upper, cooler part of the blast furnace.
  FeO(s) + CO(g) → Fe(s) + CO₂(g) (at T < 1000 K)

Electrolytic Reduction (Electrolysis)

This method is used for highly reactive metals (e.g., Al, Na, Mg) whose oxides cannot be reduced by Carbon at reasonable temperatures (their lines are very low on the Ellingham diagram).

The metal oxide or chloride is melted (often with a flux like Na₃AlF₆ for Al₂O₃ to lower the melting point) and a strong electric current is passed through it.
Example (Hall-Héroult for Al):
Cathode (Reduction): Al³⁺ + 3e^- → Al(l)
Anode (Oxidation): C(s) + 2O^2- → CO₂(g) + 4e^-

Hydrometallurgy

This involves using aqueous solutions (hydro) to extract the metal. It is common for low-grade ores and less reactive metals like Ag and Au.

Example (MacArthur-Forrest Cyanide Process):

1. Leaching: The ore (e.g., Ag) is crushed and treated with a NaCN solution, with air bubbled through. The Ag is oxidized and forms a soluble complex.
   4Ag(s) + 8CN^-(aq) + O₂(g) + 2H₂O(l) → 4[Ag(CN)₂]^-(aq) + 4OH^-(aq)
2. Precipitation (Reduction): A more reactive metal (a stronger reducing agent), like Zn, is added to the solution to displace the Ag.
   Zn(s) + 2[Ag(CN)₂]^-(aq) → [Zn(CN)₄]^2-(aq) + 2Ag(s) (Silver precipitates)

Purification (Refining) of Metals

The metal extracted (e.g., "blister copper," "pig iron") is still impure. Refining is the final step to achieve high purity.

Methods of Purification of Metals

1. Electrolytic Processes (Refining)

This is one of the most common methods for metals like Cu, Zn, Pb, Ag.

• Setup:
  • Anode: A thick block of the impure metal.
  • Cathode: A thin sheet of the pure metal.
  • Electrolyte: A solution of a salt of the metal (e.g., CuSO₄ for copper refining).
• Process:
  • At the Anode (Oxidation): The impure metal dissolves: Cu(impure) → Cu²⁺(aq) + 2e^-.
  • At the Cathode (Reduction): The pure metal plates out: Cu²⁺(aq) + 2e^- → Cu(pure).
• Impurities:
  • More reactive impurities (like Zn, Fe) also oxidize at the anode, but they are *too hard to reduce* and stay in the solution.
  • Less reactive impurities (like Ag, Au, Pt) do not oxidize. They fall off the anode and collect at the bottom as anode mud (a valuable byproduct).

2. Mond's Process

This process is specific for Nickel (Ni). It is a "vapor phase refining" method, based on the formation and decomposition of a volatile compound.

1. Formation: Impure Nickel is heated with Carbon Monoxide (CO) gas at a low temperature (50-80°C) to form volatile nickel tetracarbonyl (Ni(CO)₄) gas. Impurities are left behind.
   Ni(impure, s) + 4CO(g) \xrightarrow{50^° C} Ni(CO)₄(g)
2. Decomposition: The Ni(CO)₄ gas is then heated to a high temperature (180-220°C), causing it to decompose back into pure Ni metal and CO gas (which is recycled).
   Ni(CO)₄(g) \xrightarrow{200^° C} Ni(pure, s) + 4CO(g)

3. Zone Refining

This method is used to produce ultra-pure metals, especially for semiconductors (e.g., Silicon (Si), Germanium (Ge), Gallium (Ga)).

• Principle: Based on fractional crystallization. The principle is that impurities are more soluble in the molten (liquid) phase than in the solid phase of the metal.
• Process:
  1. An impure rod of the metal is taken.
  2. A circular, mobile heater is passed slowly from one end of the rod to the other.
  3. The heater creates a small molten zone. As the heater moves, the pure metal crystallizes (solidifies) *behind* it.
  4. The impurities "prefer" to stay in the molten zone and are thus "dragged" along with the heater to the far end of the rod.
  5. The process is repeated many times, concentrating all impurities at one end, which is then cut off.