Chapter 7: Redox Reactions 

Where redox reactions matter

• Energy and industry: Fuel burning, metal extraction, chemical manufacture, and the operation of dry and wet batteries involve redox processes.
• Other fields: Redox reactions are important in biological, pharmaceutical, metallurgical, and agricultural processes.
• Everyday and environmental relevance: Corrosion of metals and issues such as hydrogen as a fuel and the ozone hole are linked to redox phenomena.

How the definition broadened

• Initially: Oxidation meant the addition of oxygen to an element or compound.
• Later: Removal of hydrogen was also recognised as oxidation.
• Broader meaning: Addition of an electronegative element, or removal of an electropositive element, also counts as oxidation.

Reduction examples

• Oxygen removal: Mercuric oxide loses oxygen to form mercury.
• Electronegative element removal: Ferric chloride loses chlorine as it forms ferrous chloride.
• Hydrogen addition: Ethene gains hydrogen to form ethane.
• Electropositive element addition: In the mercuric chloride reaction, mercury is added to mercuric chloride.

Identifying oxidation and reduction

• Hydrogen sulphide with chlorine: Hydrogen sulphide is oxidised; chlorine is reduced by gaining hydrogen.
• Aluminothermic reaction: Aluminium is oxidised as it gains oxygen; iron oxide is reduced as it loses oxygen.
• Sodium and hydrogen: Sodium is oxidised and hydrogen is reduced.

Examples of redox reactions

• Sodium and chlorine: Sodium loses electrons; chlorine gains them.
• Sodium and oxygen or sulphur: Sodium is oxidised, while oxygen or sulphur is reduced.

Half-reactions

• Half-reaction: A redox process can be split into oxidation and reduction steps, each showing electron transfer.
• Overall reaction: Adding the half-reactions cancels electrons and gives the complete reaction.
• Ionic products: The products can be written as ions, such as \(\ce{Na+Cl-}\), \(\ce{(Na+)2O^{2-}}\) and \(\ce{(Na+)2S^{2-}}\).

Identifying agents in sodium reactions

• Sodium donates electrons, so it is oxidised and acts as the reducing agent.
• Chlorine, oxygen and sulphur accept electrons from sodium, so they are reduced and act as oxidising agents.

Sodium hydride formation

• Ionic product: Sodium hydride can be represented as \(\ce{Na+H-}\), indicating electron transfer from sodium to hydrogen.
• Conclusion: Sodium loses electrons and is oxidised; hydrogen gains electrons and is reduced. The reaction is a redox change.

Zinc and copper(II) ions

• Reaction: Zinc transfers electrons to copper(II) ions, producing zinc ions and copper metal.
• Evidence: Reddish copper deposits on zinc, while the blue colour due to \(\ce{Cu^{2+}}\) disappears as \(\ce{Zn^{2+}}\) forms.
• Electron change: Zinc is oxidised; copper(II) ions are reduced.

Copper and silver ions

• Reaction: Copper metal donates electrons to silver ions, forming copper(II) ions and silver metal.
• Observation: Silver deposits form and the solution turns blue due to \(\ce{Cu^{2+}}\).
• Equilibrium: Products are strongly favoured.

Cobalt and nickel ions

• Reaction: Cobalt transfers electrons to nickel(II) ions, forming cobalt(II) ions and nickel metal.
• Equilibrium: Both reactants and products remain at moderate concentrations; neither side is strongly favoured.

Comparing metals

• Electron competition: A metal that gives electrons to another metal's ions has a greater electron-releasing tendency.
• Observed order: Zinc transfers electrons to copper ions, and copper transfers electrons to silver ions; therefore, \(\mathrm{Zn > Cu > Ag}\) in electron-releasing tendency.
• Electrochemical series: Metals and their ions can be arranged by their tendency to release electrons; this arrangement is also called the metal activity series.
• Galvanic cells: Competitive electron-transfer reactions can produce electrical energy in galvanic cells.

Why oxidation numbers are useful

• Covalent reactions: Electron transfer may be partial in reality, but oxidation numbers track electron shifts by assuming complete transfer to the more electronegative atom.
• Bookkeeping model: The complete-transfer assumption is a useful convention, not necessarily a claim that full electrons actually move.
• Multiple atoms: If an element has two or more atoms in a molecule or ion, its stated oxidation number may be the average for those atoms.

Applying the sum rule

• Unknown value: Assign the known oxidation numbers, multiply each by the number of atoms, then set the total equal to \(0\) for a neutral compound or to the ion charge.
• Average value: When equivalent atoms of the same element have different oxidation states, the oxidation number assigned to that element is their average.

Patterns across elements

• Metals: Usually have positive oxidation numbers; non-metals may have positive or negative values.
• Transition elements: Commonly show several positive oxidation states.
• Highest oxidation number: For representative elements, it is generally the group number in Groups 1–2 and the group number minus \(10\) in Groups 13–18.
• Across a period: The highest oxidation number generally rises; in Period 3 it ranges from \(+1\) for sodium to \(+7\) for chlorine.
• Terminology: Oxidation number and oxidation state are used interchangeably.

Stock notation examples

• Gold chlorides: Aurous chloride is \(\ce{Au(I)Cl}\); auric chloride is \(\ce{Au(III)Cl3}\).
• Tin chlorides: Stannous chloride is \(\ce{Sn(II)Cl2}\); stannic chloride is \(\ce{Sn(IV)Cl4}\).
• Oxidation state comparison: \(\ce{Hg2(I)Cl2}\) is the reduced form relative to \(\ce{Hg(II)Cl2}\).

Agents in the copper reaction

• Oxidant: \(\ce{Cu2O}\) helps sulphur in \(\ce{Cu2S}\) increase its oxidation number; copper(I) in \(\ce{Cu2O}\) is reduced.
• Reductant: Sulphur in \(\ce{Cu2S}\) increases from \(-2\) to \(+4\) and helps copper in both reactants become metallic copper.

A decomposition that is not redox

• Calcium carbonate: It decomposes without any oxidation-state change: calcium stays at \(+2\), carbon at \(+4\), and oxygen at \(-2\).
• Check: Classify by checking oxidation-state changes, not by the reaction pattern alone.

Halogens in alkali

• Bleaching: The hypochlorite ion, \(\ce{ClO^-}\), oxidises coloured stains to colourless compounds.
• Bromine and iodine: They show a disproportionation trend similar to chlorine in alkali.
• Fluorine exception: Fluorine cannot have a positive oxidation state, so it does not disproportionate. It forms fluoride and oxygen difluoride with alkali; it also attacks water, producing some oxygen.

Further examples

• Mixed oxides: \(\ce{Fe3O4}\), \(\ce{Mn3O4}\), and \(\ce{Pb3O4}\) have fractional average oxidation numbers because the metal atoms occur in more than one whole-number state.
• Dioxygen ions: Oxygen has oxidation state \(+\frac{1}{2}\) in \(\ce{O2^+}\) and \(-\frac{1}{2}\) in \(\ce{O2^-}\).

Two approaches

• Oxidation number method: Balance using the increase and decrease in oxidation numbers.
• Half-reaction method: Separate the reaction into oxidation and reduction half-reactions, balance each, then add them.
• Choice of method: Either method may be used; choose the one that is more convenient for the reaction.

Procedure

• 1. Write formulas: Write correct formulas for all reactants and products.
• 2. Assign oxidation numbers: Identify the atoms whose oxidation numbers change.
• 3. Equalise electron change: Calculate the increase or decrease per atom and per species; use coefficients so total increase equals total decrease.
• 4. Balance charge: In aqueous reactions, use \(\ce{H+}\) in acidic solution or \(\ce{OH-}\) in basic solution to equalise ionic charges.
• 5. Balance H and O: Add \(\ce{H2O}\) to balance hydrogen, then verify oxygen, all atoms and total charge.

Example: dichromate and sulphite in acidic solution

• Oxidation-number changes: Cr changes from \(+6\) to \(+3\); S changes from \(+4\) to \(+6\).
• Electron-change match: Two Cr atoms gain \(6\) units in total; three S atoms lose \(6\) units in total.
• Agents: \(\ce{Cr2O7^2-}\) is the oxidising agent; \(\ce{SO3^2-}\) is the reducing agent.

Example: permanganate and bromide in basic solution

• Oxidation-number changes: Mn changes from \(+7\) to \(+4\) (gain of \(3\)); Br changes from \(-1\) to \(+5\) (loss of \(6\)).
• Electron-change match: Use \(2\) permanganate ions for each bromide ion.

Example: Fe²⁺ oxidised by dichromate in acidic solution

• Oxidation half-reaction: Iron loses one electron per ion.
• Reduction half-reaction: Dichromate gains six electrons; multiply the iron half-reaction by \(6\) to match.

Example: permanganate oxidises iodide in basic solution

• Half-reactions: Iodide is oxidised to iodine; permanganate is reduced to manganese dioxide.
• Electron match: Multiply the oxidation half-reaction by \(3\) and the reduction half-reaction by \(2\); both then exchange \(6\) electrons.

Important endpoint details

• Permanganate endpoint: The lasting pink tint is detectable at about \(10^{-6}\ \mathrm{mol\,dm^{-3}}\), giving only a small colour overshoot.
• Equivalence point: The oxidant and reductant are present in the mole ratio required by the balanced equation.
• Iodine in solution: Although iodine is sparingly soluble in water, it remains in a solution containing KI as \(\ce{KI3}\).

• Modern interpretation: Oxidation is viewed as a decrease in electron density around the atoms involved; reduction is an increase in electron density.
• Key idea: Oxidation number is a useful accounting tool, but redox can also be understood through changes in electron density.

Direct Electron Transfer

• Zinc and copper ions: When zinc is dipped in copper sulphate solution, zinc transfers electrons directly to copper ions; heat is evolved.
• Oxidation: Zinc loses electrons to form zinc ions; reduction: copper ions gain electrons to form copper metal.

Redox-Couple Notation

• Order: Write the oxidised form first, followed by a slash or vertical line for the interface, then the reduced form.
• Examples: Zinc and copper couples are written \(\ce{Zn^2+/Zn}\) and \(\ce{Cu^2+/Cu}\).

Cell Setup and Operation

• Set-up: A zinc rod sits in zinc sulphate solution and a copper rod in copper sulphate solution; the rods connect by a wire and the solutions by a salt bridge.
• Salt bridge: A U-tube containing an electrolyte such as potassium chloride or ammonium nitrate, often set in agar gel; it permits electrical contact without allowing the solutions to mix.
• Closed circuit: Electrons pass from zinc to copper through the external wire; ions migrate through the salt bridge to complete the circuit.
• Open switch: No reaction occurs in either beaker and no current flows through the wire.
• Potential difference: Current flows when a potential difference exists between the zinc and copper electrodes.

Reference and Meaning

• Reference: By convention, the standard potential of the hydrogen electrode, \(\ce{H+/H2}\), is \(0.00\ \mathrm{V}\).
• Significance: Electrode potential measures the relative tendency of the active species in a redox couple to remain in its oxidised or reduced form.
• Negative \(E^\circ\): The redox couple is a stronger reducing agent than the \(\ce{H+/H2}\) couple.
• Positive \(E^\circ\): The redox couple is a weaker reducing agent than the \(\ce{H+/H2}\) couple.

Trends in the Potential Table

• Oxidising strength: Increases upward in the standard reduction-potential table.
• Reducing strength: Increases downward in the table; the reducing forms of couples with more negative \(E^\circ\) are stronger.

Agents and Redox-Equation Methods

• Oxidising agent (oxidant): Accepts electrons and is itself reduced.
• Reducing agent (reductant): Donates electrons and is itself oxidised.
• Equation writing: Oxidation-number and ion–electron methods are useful for writing and balancing redox equations.