redox reactions
Redox basics
- Redox reactions involve transfer of electrons between species; oxidation state changes track this transfer.
- Key definitions:
- Oxidation: increase in oxidation state of an element.
- Reduction: decrease in oxidation state of an element.
- Oxidising agent: the substance that is itself reduced in a redox reaction; it causes oxidation of the other substance. In other words, it accepts electrons.
- Reducing agent: the substance that is itself oxidised in a redox reaction; it donates electrons to the other substance.
- Oxidation states (oxidation numbers) are a bookkeeping method to keep track of electron transfer.
- Common way to identify agents:
- If a species undergoes an increase in oxidation state, it is oxidised and acts as a reducing agent (it reduces another species).
- If a species undergoes a decrease in oxidation state, it is reduced and acts as an oxidising agent (it oxidises another species).
- Simple rule-of-thumb examples from the transcript:
- Hydrogen oxidation: H₂ is oxidised to form water, when its oxidation state increases from 0 to +1 (in H₂O).
- Copper oxide reduction: Copper in CuO is reduced from Cu²⁺ to Cu⁰ (oxidation state decreases from +2 to 0). This occurs alongside the oxidation of another species (e.g., H₂ delivering electrons).
- In many redox examples, the oxidising agent is reduced and the reducing agent is oxidised.
Illustrative redox reactions from the transcript
- Methane oxidation example (common textbook redox):
- Reaction (overall):
ext{CH}4 + 2 ext{O}2
ightarrow ext{CO}2 + 2 ext{H}2 ext{O} - Oxidation state changes (conceptual):
- Carbon in CH₄ has an oxidation state of −4. In CO₂, carbon is +4. Change: increase by 8 electrons (carbon is oxidised).
- Oxygen in O₂ goes from 0 to −2 in CO₂ and H₂O (oxygen is reduced).
- This demonstrates the general oxidation-number bookkeeping used to identify oxidised and reduced species in a reaction.
- Reaction (overall):
- Copper oxide reduction example (CuO + H₂ → Cu + H₂O):
- Reactants and products:
ext{CuO + H}2 ightarrow ext{Cu + H}2 ext{O} - Oxidation state changes:
- Cu in CuO is +2; Cu in Cu is 0. Reduction of Cu²⁺ to Cu⁰.
- Hydrogen goes from 0 in H₂ to +1 in H₂O (oxidation).
- This shows CuO as the oxidising agent (it is reduced) and H₂ as the reducing agent (it is oxidised).
- Reactants and products:
- Hydration/oxidation context for H₂O formation:
- Hydrogen is oxidised to form water due to the increase in oxidation state from 0 to +1.
- Additional standard redox example mentioned in the transcript:
- Combustion-like oxidation of methane (as above) contrasted with reduction of methane to CO₂ and H₂O in the presence of an oxidant that accepts electrons.
Oxidising and reducing agents: practical identifications
- Oxidising agents (oxidants) – substances that get reduced:
- Often represented by species that gain electrons and cause another species to be oxidised.
- In acidic solution, permanganate ion is a classic oxidising agent:
- Half-reaction (in acidic solution):
ext{MnO}4^- + 8 ext{H}^+ + 5 e^- ightarrow ext{Mn}^{2+} + 4 ext{H}2 ext{O} - Observed color change: purple
→ colorless (overall MnO₄⁻ to Mn²⁺ change).
- Dichromate as oxidising agent:
- Common example is acidified potassium dichromate:
- Half-reaction (in acidic solution):
ext{Cr}2 ext{O}7^{2-} + 14 ext{H}^+ + 6 e^-
ightarrow 2 ext{Cr}^{3+} + 7 ext{H}_2 ext{O} - Typical color change: orange dichromate to green Cr³⁺ (in solution).
- Reducing agents – substances that get oxidised:
- Potassium iodide (I⁻) is a common reducing agent in redox chemistry:
- Oxidation half-reaction:
2 I^-
ightarrow I_2 + 2 e^- - I⁻ is oxidised to I₂; the oxidising agent in the paired reaction is reduced.
- In many reactions, iodide acts as the reducing agent (it donates electrons) and chlorine or other oxidants accept electrons.
- Other recognisable redox pair example from the transcript:
- Br₂ + SO₂ (with a hydrogen-containing medium) can react to form products such as HBr and SO₃ (as depicted in the transcript). The exact balanced form in the text is unclear; a representative redox form is:
ext{Br}2 + ext{SO}2 + ext{H}2 ext{O} ightarrow 2 ext{HBr} + ext{SO}3 - Note: The transcript’s exact coefficients/equation may contain transcription errors; use this as a qualitative example of a halogen/oxide redox couple.
- Br₂ + SO₂ (with a hydrogen-containing medium) can react to form products such as HBr and SO₃ (as depicted in the transcript). The exact balanced form in the text is unclear; a representative redox form is:
Practice questions and what to look for
- Identify oxidising agent and reducing agent in a given reaction:
- Determine which species increases in oxidation state (reducing agent) and which decreases (oxidising agent).
- Example identifications mentioned:
- In the reaction of Fe with Cl₂ to form iron chloride, chlorine acts as the oxidising agent (it is reduced) and iron acts as the reducing agent (it is oxidised). A balanced form of a related common redox equation is:
2 ext{Fe} + 3 ext{Cl}2 ightarrow 2 ext{FeCl}3
- In the reaction of Fe with Cl₂ to form iron chloride, chlorine acts as the oxidising agent (it is reduced) and iron acts as the reducing agent (it is oxidised). A balanced form of a related common redox equation is:
- The transcript also discusses the concept that chlorine can oxidise iron(II) iodide while iodide is oxidised to iodine, illustrating the mutual redox roles:
- Oxidation: iodide to iodine, reduction: chlorine to chloride.
- A typical related set of half-reactions:
- 2 I^-
ightarrow I_2 + 2 e^- - ext{Cl}_2 + 2 e^-
ightarrow 2 Cl^-
- Fill-in-the-blank redox definitions (as per transcript, corrected for clarity):
- (i) Oxidation is the loss of hydrogen or the gain of oxygen.
- (ii) Reduction is the gain of hydrogen or the loss of oxygen.
- Practice-application example from the transcript:
- Bromine and sulfur dioxide reaction question (text notes): identify whether bromine or sulfur dioxide is oxidised or reduced in the given reaction; the exact coefficients may be unclear in the source, but the exercise emphasizes determining oxidation state changes for each participant.
Quick recap of key concepts to memorize
- Oxidation state changes track electron transfer: rise in oxidation state = oxidation; fall = reduction.
- Oxidising agent is reduced; it oxidises the other substance.
- Reducing agent is oxidised; it reduces the other substance.
- Common redox couples with observable color changes:
- Permanganate in acid:
ext{MnO}_4^-
ightarrow ext{Mn}^{2+} (purple → colorless in solution) - Dichromate in acid:
ext{Cr}2 ext{O}7^{2-}
ightarrow 2 ext{Cr}^{3+} (orange → green in solution) - Iodide oxidation:
2 I^-
ightarrow I_2 + 2 e^- (colorless to brown/blue depending on solvent)
- Permanganate in acid:
- Representative balanced redox equations from foundational examples:
- ext{CH}4 + 2 ext{O}2
ightarrow ext{CO}2 + 2 ext{H}2 ext{O} - ext{CuO} + ext{H}2 ightarrow ext{Cu} + ext{H}2 ext{O}
- 2 ext{Fe} + 3 ext{Cl}2 ightarrow 2 ext{FeCl}3
- ext{CH}4 + 2 ext{O}2
Notes on transcript accuracy
- Several items in the transcript contain transcription errors or unclear formatting (e.g., the ammonia/oxygen combustion equation on Page 1, and the Br₂/SO₂ reaction notation on Page 3). The concepts and standard examples are retained, but some exact coefficients/equations are presented in their conventional, corrected forms where appropriate to ensure accuracy for studying.