Electrochemistry

Cheat Sheet  ·  Chemistry Gr 11
Emil Oliversen
REDOX BASICS
OIL RIG
Oxidation = Loss of e⁻
Reduction = Gain of e⁻
Oxidizing agent
accepts e⁻ → gets reduced
Reducing agent
donates e⁻ → gets oxidized
OXIDATION STATE RULES
Pure element
OS = 0
Monoatomic ion
OS = charge
O (usually)
−2 (H₂O₂: −1; OF₂: +2)
H (usually)
+1 (metal hydrides: −1)
Sum rule
sum = overall charge
Increase in OS → oxidized
Decrease in OS → reduced
WORKED OS EXAMPLES
Mn in KMnO₄
+1+Mn−8=0 → Mn=+7
Cr in Cr₂O₇²⁻
2Cr−14=−2 → Cr=+6
S in H₂SO₄
2+S−8=0 → S=+6
BALANCING REDOX (acidic)
Half-reaction method
  1. Split into half-reactions
  2. Balance atoms other than O and H
  3. Add H₂O to balance O
  4. Add H⁺ to balance H
  5. Add e⁻ to balance charge
  6. Multiply to equalize e⁻ in both halves
  7. Add half-reactions; cancel e⁻
Check: atoms balanced AND charges balanced
WORKED EXAMPLE
MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺ (acid)
Oxidation ×5
5Fe²⁺ → 5Fe³⁺ + 5e⁻
Reduction ×1
MnO₄⁻+8H⁺+5e⁻→Mn²⁺+4H₂O
Final
MnO₄⁻+5Fe²⁺+8H⁺→Mn²⁺+5Fe³⁺+4H₂O
SIMPLE EXAMPLE
Oxidation
Zn → Zn²⁺ + 2e⁻
Reduction
Cu²⁺ + 2e⁻ → Cu
Net
Zn + Cu²⁺ → Zn²⁺ + Cu
GALVANIC CELLS
Anode
oxidation · loses e⁻ · (−)
Cathode
reduction · gains e⁻ · (+)
e⁻ flow
anode → wire → cathode
Ion flow
through salt bridge
Cell notation:
A | A⁺ || B²⁺ | B
(anode left, cathode right, || = salt bridge)
ELECTRODE POTENTIALS
E°cell
E°cathode − E°anode
Positive E°cell
spontaneous (galvanic) ✓
Negative E°cell
non-spontaneous (electrolytic)
SHE (standard hydrogen electrode) = 0.00 V reference
WORKED EXAMPLE
Zn | Zn²⁺ || Cu²⁺ | Cu
E°(Cu²⁺/Cu) = +0.34 V (cathode)
E°(Zn²⁺/Zn) = −0.76 V (anode)
E°cell = 0.34 − (−0.76) = +1.10 V → spontaneous
ELECTROLYSIS
Driven by
external electrical energy
Cathode (−)
reduction (same as galvanic)
Anode (+)
oxidation (connected to + of battery)
Faraday F
96 500 C/mol e⁻
Q = I × t
mol e⁻ = Q / F
ACTIVITY SERIES (most → least reactive)
MetalE° (V)Role
K−2.93Best reducing agent
Na−2.71
Mg−2.37
Al−1.66
Zn−0.76
Fe−0.44
Ni−0.25
Sn−0.14
Pb−0.13
H0.00SHE reference
Cu+0.34
Ag+0.80
Au+1.50Best oxidizing agent
More active = better reducing agent  ·  Less active = better oxidizing agent
COMMON MISTAKES
  • Anode = oxidation ALWAYS (regardless of cell type)
  • In galvanic: anode is negative (−)
  • In electrolytic: anode is positive (+)
  • OS of O in H₂O₂ = −1, not −2
  • E°cell = E°cathode − E°anode (not sum)
  • Reducing agent gets oxidized (loses e⁻)
  • Oxidizing agent gets reduced (gains e⁻)
AN OX · RED CAT
ANOde=OXidation · REDuction at CAThode