Ch8: Redox Reactions

Oxidation of Copper & the Statue of Liberty

• Public misconception: Many believe the Statue of Liberty was always green.
• Original surface: bare copper (brown-reddish, penny-like).
• Green patina appeared gradually through oxidation.
• Net reaction (simplified): $4Cu<em>{(s)} + O</em>2 \longrightarrow 2Cu<em>2O</em>{(s)}$
• Balancing check: $4 Cu$ (left) $\rightarrow 4 Cu$ (right); $2 O$ (left) $\rightarrow 2 O$ (right).
• Oxidation-state analysis
  • $Cu : 0 \;\longrightarrow\; +1$ (increase $\Rightarrow$ oxidation)
  • $O : 0 \;\longrightarrow\; -2$ (decrease $\Rightarrow$ reduction)
• Consequence: brown $Cu$ turns into green $Cu_2O$ (and later basic copper carbonates not covered in this lecture) — the patina that we see.

Fundamentals of Redox Reactions

• Redox = Reduction + Oxidation occurring simultaneously.
• Key electron ideas
  • Reduction: gain of electrons.
  • Oxidation: loss of electrons.
• Electricity production in galvanic/voltaic/electrochemical/Daniell cells comes from electron transfer between substances.
• Every redox process can be written as two half-reactions (one oxidation, one reduction).

Mnemonic - "LEO the lion says GER"

• Loss of Electrons = Oxidation.
• Gain of Electrons = Reduction.
• Instructor phrased it humorously as “Leo the lion say scare.”

Galvanic/Voltaic Cell Example: Zn–Ag Battery

• Electrodes
  • Anode: Zinc (undergoes oxidation).
  • Cathode: Silver (undergoes reduction).
• Half-reactions
  • Oxidation (anode): $Zn \rightarrow Zn^{2+} + 2e^-$
  • Reduction (cathode): $Ag^+ + e^- \rightarrow Ag$
• Oxidation-state changes
  • $Zn: 0 \rightarrow +2$ (loses $2 e^-$.)
  • $Ag: +1 \rightarrow 0$ (gains $1 e^-$.)
• Overall: electron flow through external circuit = usable electrical energy.

Reducing vs. Oxidizing Agents

• Substance oxidized $\rightarrow$ acts as reducing agent (because it supplies $e^-$ to reduce another species).
• Substance reduced $\rightarrow$ acts as oxidizing agent (because it removes $e^-$ from another species).
• Rule of thumb provided:
  • “Reduced $\Rightarrow$ oxidizing agent.”
  • “Oxidized $\Rightarrow$ reducing agent.”

Identifying Half-Reactions Quickly

1. Check electron location:
   • $e^-$ on reactant side $\rightarrow$ reduction.
   • $e^-$ on product side $\rightarrow$ oxidation.
2. Check oxidation-number trend:
   • Number increases $\Rightarrow$ oxidation.
   • Number decreases $\Rightarrow$ reduction.

Practice Classification (Classroom Exercise A–E)

A. Electrons on reactant side & oxidation state decreases $\rightarrow$ Reduction.
B. Electrons on product side & oxidation state increases $\rightarrow$ Oxidation.
C. Initially mis-read; correct classification after examining $Mn^{+7} \rightarrow Mn^{+4}$ ($e^-$ on reactant side) $\rightarrow$ Reduction.
D. More challenging example 2H2O \rightarrow O2 + 4H^+ + 4e^-$

• e^-$on product side; O:$-2 \rightarrow 0$$ \Rightarrow$<strong>Oxidation</strong>.<br /> E. Another given equation with$e^-$on reactant side and oxidation state drop of$3 \Rightarrow Reduction.

Quick Rules & Tips Re-emphasized

• Oxidation number = hypothetical charge determined by electronegativity and bookkeeping.
• In water O$usually$-2$,$H$usually$+1$$.
• Electron bookkeeping in half-reactions must balance charge as well as atoms.
• Always verify with both methods (electron position & oxidation-number change) for confidence.

Contextual / Pedagogical Notes

• Instructor repeatedly asked students to raise hands, promoting active participation.
• Emphasized being “confident” in answers after applying the two diagnostic tests.
• Promised deeper coverage of oxidation-state assignment in subsequent chapter but supplied needed numbers for current exercise.
• Encouraged students to remember mnemonic and agent relationships for all future redox analyses.