Comprehensive Study Guide for Redox Reactions

Learning Objectives for Redox Reactions

  • Oxidation Numbers: Use Roman numerals to indicate the oxidation number of an element in a compound.
  • Simultaneous Reactions: Define redox reactions as processes involving simultaneous oxidation and reduction.
  • Oxygen Transfer:
    • Oxidation: Defined as the gain of oxygen.
    • Reduction: Defined as the loss of oxygen.
  • Electron Transfer and Oxidation Numbers:
    • Oxidation: Defined as the loss of electrons or an increase in oxidation number.
    • Reduction: Defined as the gain of electrons or a decrease in oxidation number.
  • Identification of Redox: Identify redox reactions based on the gain and loss of oxygen, gain and loss of electrons, or changes in oxidation numbers.
  • Rules for Oxidation Numbers:
    • The oxidation number of elements in their uncombined state is 00.
    • The oxidation number of a monatomic ion is equal to the charge on the ion.
    • The sum of oxidation numbers in a neutral compound is 00.
    • The sum of oxidation numbers in a polyatomic ion is equal to the charge on the ion.
  • Chemical Agents:
    • Oxidising Agent: A substance that oxidises another substance and is itself reduced.
    • Reducing Agent: A substance that reduces another substance and is itself oxidised.
  • Laboratory Tests: Identify redox reactions by colour changes when using acidified aqueous potassium manganate(VII) or aqueous potassium iodide.

Oxidation and Reduction via Oxygen Transfer

  • Basic Definition: A redox reaction is where oxidation and reduction take place together at the same time in the same reaction.
  • Oxidation (Oxygen Gain):
    • Example: Magnesium burning in air.
    • Reaction: 2Mg(s)+O2(g)→2MgO(s)2Mg_{(s)} + O_{2(g)} \rightarrow 2MgO_{(s)}
    • Observation: Magnesium burns with a dazzling white flame to form white ash (magnesium oxide). The magnesium has gained oxygen and is therefore oxidised.
    • Example: Heating copper in air.
    • Reaction: 2Cu(s)+O2(g)→2CuO(s)2Cu_{(s)} + O_{2(g)} \rightarrow 2CuO_{(s)}
    • Observation: Surface of the copper becomes coated with black copper(II) oxide. This is an oxidation reaction even though it is not highly exothermic and produces no flame.
  • Reduction (Oxygen Loss):
    • Example: Passing hydrogen over heated copper(II) oxide.
    • Reaction: CuO(s)+H2(g)→Cu(s)+H2O(l)CuO_{(s)} + H_{2(g)} \rightarrow Cu_{(s)} + H_{2}O_{(l)}
    • Observation: The black copper(II) oxide turns pink as it loses oxygen to become copper metal. The copper(II) oxide is reduced.
    • Example: Reaction between zinc oxide and carbon.
    • Zinc oxide is reduced (loses oxygen) and carbon is oxidised (gains oxygen).
  • Specific Examples and Reasons:
    • Reaction: 2Mg+CO2→2MgO+C2Mg + CO_2 \rightarrow 2MgO + C (Magnesium is oxidised; Carbon dioxide is reduced because it lost oxygen).
    • Metal Extraction: When Co3O4Co_3O_4 is heated in hydrogen to form cobalt metal (Co3O4+4H2→3Co+4H2OCo_3O_4 + 4H_2 \rightarrow 3Co + 4H_2O), the Co3O4Co_3O_4 is reduced because it loses oxygen.
    • Nitric Acid Production: In the reaction 4NH3+5O2→4NO+6H2O4NH_3 + 5O_2 \rightarrow 4NO + 6H_2O, the substance oxidised is NH3NH_3 because it gains oxygen to become NONO.

Oxidation and Reduction via Electron Transfer (OIL RIG)

  • Definitions:
    • OIL: Oxidation Is Loss of electrons.
    • RIG: Reduction Is Gain of electrons.
  • Example: Magnesium Oxide Formation:
    • Equation: 2Mg(s)+O2(g)→2MgO(s)2Mg_{(s)} + O_{2(g)} \rightarrow 2MgO_{(s)}
    • Process: Each magnesium atom loses two electrons (Mg→Mg2++2e−Mg \rightarrow Mg^{2+} + 2e^-) and each oxygen atom gains two electrons per atom (O+2e−→O2−O + 2e^- \rightarrow O^{2-}).
    • Magnesium is oxidised (electron loss) and oxygen is reduced (electron gain).
  • Redox Without Oxygen:
    • Sodium and Chlorine: 2Na(s)+Cl2(g)→2NaCl(s)2Na_{(s)} + Cl_{2(g)} \rightarrow 2NaCl_{(s)}
    • Sodium half-equation (Oxidation): 2Na→2Na++2e−2Na \rightarrow 2Na^+ + 2e^-
    • Chlorine half-equation (Reduction): Cl2+2e−→2Cl−Cl_2 + 2e^- \rightarrow 2Cl^-
    • Sodium acts as the electron donor and is oxidised; chlorine acts as the electron acceptor and is reduced.
    • Displacement of Bromine: Cl2(g)+2KBr(aq)→2KCl(aq)+Br2(aq)Cl_{2(g)} + 2KBr_{(aq)} \rightarrow 2KCl_{(aq)} + Br_{2(aq)}
    • Chlorine half-equation (Reduction): Cl2+2e−→2Cl−Cl_2 + 2e^- \rightarrow 2Cl^-
    • Bromide ion half-equation (Oxidation): 2Br−→Br2+2e−2Br^- \rightarrow Br_2 + 2e^-
    • Ionic Equation: Cl2+2Br−→2Cl−+Br2Cl_2 + 2Br^- \rightarrow 2Cl^- + Br_2

Oxidation States and Numbers

  • Definition: Oxidation state indicates how many electrons each atom of an element has gained, lost, or shared in forming a compound. It shows the degree of oxidation or reduction.
  • Notation: Usually written with the sign first followed by the number (e.g., +7+7) or using Roman numerals (e.g., (VII)).
  • Standard Oxidation States in Compounds:
    • Group I metals (e.g., sodium): +I+I
    • Group II metals (e.g., calcium): +II+II
    • Aluminium: +III+III
    • Hydrogen: +1+1
    • Oxygen: −II-II (except in peroxides)
    • Group VII non-metals (without oxygen): −1-1
  • Transition Metal Variability:
    • Iron: +II+II and +III+III
    • Copper: +I+I and +II+II
    • Manganese: +II,+IV,+V,+VII+II, +IV, +V, +VII
    • Chromium: +III+III and +VI+VI
  • Trend Identification:
    • A rise in oxidation number signifies oxidation.
    • A fall in oxidation number signifies reduction.
    • Example: 2FeCl2+Cl2→2FeCl32FeCl_2 + Cl_2 \rightarrow 2FeCl_3
    • Fe changes from +II+II to +III+III (Oxidation).
    • Cl changes from 00 to −I-I (Reduction).

Oxidising and Reducing Agents

  • Oxidising Agent:
    • Definition: A substance that provides oxygen for another substance or takes electrons from another substance.
    • The agent itself is reduced.
    • Common Examples: Oxygen, Chlorine, Potassium manganate(VII), hydrogen peroxide.
  • Reducing Agent:
    • Definition: A substance that removes oxygen from another substance or gives electrons to another substance.
    • The agent itself is oxidised.
    • Common Examples: Hydrogen, Carbon, Carbon monoxide, reactive metals (e.g., sodium).
  • Summary of Roles:
    • Electron donor = Reducing agent (is oxidised).
    • Electron acceptor = Oxidising agent (is reduced).

Laboratory Tests and Colour Changes

  • Potassium Manganate(VII) (Acidified):
    • Function: Used as a test for reducing agents.
    • The manganese atom in KMnO4KMnO_4 is in oxidation state +VII+VII. It is strongly driven to move to the more stable +II+II state.
    • Equation: MnO_4^-_{(aq)} (purple) + 5e^- \rightarrow Mn^{2+}_{(aq)} (colourless)
    • Result: If a reducing agent is present, the purple solution turns colourless.
  • Potassium Iodide (KIKI):
    • Function: Used as a test for oxidising agents.
    • The iodide ion (I−I^-) is oxidised to iodine (I2I_2).
    • Equation: 2I(aq)−(colourless)→I2(aq)(yellow−brown)+2e−2I^-_{(aq)} (colourless) \rightarrow I_{2(aq)} (yellow-brown) + 2e^-
    • Result: If an oxidising agent is present, the colourless solution turns yellow-brown (or red-brown). If starch is added, it turns dark blue.

Everyday and Industrial Applications of Redox

  • Corrosion: Reactive metals attacked by air and water. The rusting of iron/steel in damp air produces iron(III) oxide (rust), which weakens structures like bridges and ships.
  • Rancid Food: Oxidation of fats and oils in butter/margarine leads to unpleasant tastes and smells. Antioxidants or airtight containers are used to prevent this.
  • Metal Extraction (Blast Furnace):
    • Carbon dioxide is reduced to carbon monoxide: CO2+C→2COCO_2 + C \rightarrow 2CO
    • Iron(III) oxide (Fe2O3Fe_2O_3) is reduced to iron by carbon monoxide: Fe2O3+3CO→2Fe+3CO2Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2
  • Household Cleaners: Often contain oxidising agents (like bleach) to kill bacteria.

Questions & Discussion

  • Question 1: Which reactions involve oxidation and reduction? (List provided: A. Hexane combustion, B. Magnesium burning, C. Calcium carbonate decomposition, D. Magnesium + Copper oxide, E. HCl + NaOH).
    • Answer: A, B, and D. (C is thermal decomposition; E is neutralisation).
  • Question 2: Which reactions usually involve burning?
    • Answer: A and B.
  • Question 3: What type of reaction happened to copper(II) oxide in equation D (Mg+CuO→MgO+CuMg + CuO \rightarrow MgO + Cu)?
    • Answer: Reduction.
  • Question 4: Is it possible to have oxidation without reduction in a chemical reaction?
    • Answer: No. Even in magnesium burning, magnesium is oxidised while oxygen is reduced (it is no longer a free element).
  • Question 5: Which definition of redox is more useful: oxygen transfer or electron transfer?
    • Answer: While oxygen transfer is simpler to understand, the electron-based definition is more useful because it includes more types of reactions.
  • Question 6: Calculate oxidation numbers for underlined elements:
    • Al‾3+\underline{Al}^{3+}: +3+3
    • Cl‾O3−\underline{Cl}O_3^-: +5+5 (since Cl+3(−2)=−1→Cl−6=−1Cl + 3(-2) = -1 \rightarrow Cl - 6 = -1).
    • O‾3\underline{O}_3: 00 (uncombined element).
    • P‾Cl3\underline{P}Cl_3: +3+3
    • Cr‾2O72−\underline{Cr}_2O_7^{2-}: +6+6 (since 2Cr+7(−2)=−2→2Cr−14=−2→2Cr=122Cr + 7(-2) = -2 \rightarrow 2Cr - 14 = -2 \rightarrow 2Cr = 12).
  • Question 7: Describe the colour change when bromine reacts with potassium iodide.
    • Answer: Colourless to yellow-brown.