Electrode Potentials and Electrochemical Series

Introduction

Why does zinc dissolve in copper sulfate solution but copper doesn’t dissolve in zinc sulfate? Why can’t we store ferrous sulfate solution in a copper container? The answer lies in electrode potentials - a ranking system that predicts which metals are stronger reducing agents and which ions are better oxidizing agents!

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Interactive: Electrochemical Series

Explore the electrochemical series and predict reactions:

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Standard Electrode Potential (E°)

Definition

The standard electrode potential (E°) is the potential difference developed between the electrode and the electrolyte when:

• Concentration of ions = 1 M
• Temperature = 298 K (25°C)
• Pressure (for gases) = 1 bar
• The other electrode is Standard Hydrogen Electrode (SHE) at E° = 0.00 V

Convention

All electrode potentials are written as reduction potentials:

$$\boxed{\text{Oxidized form} + ne^- \rightarrow \text{Reduced form}} \quad E°$$

Examples:

• $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$    E° = +0.34 V
• $\text{Zn}^{2+} + 2e^- \rightarrow \text{Zn}$    E° = -0.76 V
• $\text{Ag}^+ + e^- \rightarrow \text{Ag}$    E° = +0.80 V

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Measuring Electrode Potential

Setup with Standard Hydrogen Electrode

To measure E°(Cu²⁺/Cu):

Cell:

$$\text{Pt} | \text{H}_2(1 \text{ bar}) | \text{H}^+(1M) || \text{Cu}^{2+}(1M) | \text{Cu}$$

Measured EMF: +0.34 V

Since H₂ is oxidized (anode) and Cu²⁺ is reduced (cathode):

$$E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}}$$ $$0.34 = E°(\text{Cu}^{2+}/\text{Cu}) - 0$$ $$\boxed{E°(\text{Cu}^{2+}/\text{Cu}) = +0.34 \text{ V}}$$

For Negative Electrode Potentials

To measure E°(Zn²⁺/Zn):

Cell:

$$\text{Zn} | \text{Zn}^{2+}(1M) || \text{H}^+(1M) | \text{H}_2(1 \text{ bar}) | \text{Pt}$$

Measured EMF: +0.76 V

Here Zn is oxidized (anode) and H⁺ is reduced (cathode):

$$0.76 = 0 - E°(\text{Zn}^{2+}/\text{Zn})$$ $$\boxed{E°(\text{Zn}^{2+}/\text{Zn}) = -0.76 \text{ V}}$$

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The Electrochemical Series

The electrochemical series arranges elements in order of their standard reduction potentials (E°).

Complete Electrochemical Series

Half-Reaction	E° (V)	Strength
F₂ + 2e⁻ → 2F⁻	+2.87	↑
Au³⁺ + 3e⁻ → Au	+1.50
Cl₂ + 2e⁻ → 2Cl⁻	+1.36	Strong
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O	+1.51	Oxidizing
Br₂ + 2e⁻ → 2Br⁻	+1.09	Agents
Ag⁺ + e⁻ → Ag	+0.80
Fe³⁺ + e⁻ → Fe²⁺	+0.77
I₂ + 2e⁻ → 2I⁻	+0.54
Cu²⁺ + 2e⁻ → Cu	+0.34	↓
2H⁺ + 2e⁻ → H₂	0.00	Reference
Pb²⁺ + 2e⁻ → Pb	-0.13	↑
Sn²⁺ + 2e⁻ → Sn	-0.14
Ni²⁺ + 2e⁻ → Ni	-0.25
Fe²⁺ + 2e⁻ → Fe	-0.44	Strong
Zn²⁺ + 2e⁻ → Zn	-0.76	Reducing
Al³⁺ + 3e⁻ → Al	-1.66	Agents
Mg²⁺ + 2e⁻ → Mg	-2.37
Na⁺ + e⁻ → Na	-2.71
Ca²⁺ + 2e⁻ → Ca	-2.87
K⁺ + e⁻ → K	-2.93
Li⁺ + e⁻ → Li	-3.05	↓

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Interpreting the Electrochemical Series

1. As Oxidizing Agents (Getting Reduced)

Higher E° = Stronger oxidizing agent

• Top of series (F₂, Au³⁺, Cl₂) are excellent oxidizing agents
• They readily gain electrons (get reduced)
• F₂ is the strongest oxidizing agent known!

Order of oxidizing strength:

$$\text{F}_2 > \text{MnO}_4^- > \text{Cl}_2 > \text{Br}_2 > \text{Ag}^+ > \text{Cu}^{2+} > \text{H}^+$$

2. As Reducing Agents (Getting Oxidized)

Lower E° = Stronger reducing agent

• Bottom of series (Li, K, Na, Mg) are excellent reducing agents
• They readily lose electrons (get oxidized)
• Li is the strongest reducing agent in the series!

Order of reducing strength:

$$\text{Li} > \text{K} > \text{Na} > \text{Mg} > \text{Al} > \text{Zn} > \text{Fe} > \text{H}_2$$

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Predicting Reaction Spontaneity

Rule for Spontaneity

A redox reaction is spontaneous if:

$$\boxed{E°_{\text{cell}} = E°_{\text{cathode}} - E°_{\text{anode}} > 0}$$

Equivalently:

$$\boxed{E°_{\text{reduction}} > E°_{\text{oxidation}}}$$

The Higher-Lower Rule

Spontaneous reaction occurs when:

• Species with higher E° gets reduced (cathode)
• Species with lower E° gets oxidized (anode)

Example 1: Zn + Cu²⁺

Question: Will Zn reduce Cu²⁺?

Half-reactions:

• $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$    E° = +0.34 V
• $\text{Zn}^{2+} + 2e^- \rightarrow \text{Zn}$    E° = -0.76 V

Analysis:

• Cu²⁺ has higher E° → gets reduced (cathode)
• Zn has lower E° → gets oxidized (anode)

E°cell = 0.34 - (-0.76) = +1.10 V > 0

Answer: Yes, spontaneous! ✓

Example 2: Cu + Zn²⁺

Question: Will Cu reduce Zn²⁺?

Analysis:

• Zn²⁺/Zn: E° = -0.76 V (lower)
• Cu²⁺/Cu: E° = +0.34 V (higher)

For Cu to reduce Zn²⁺:

• Zn²⁺ must be reduced (cathode)
• Cu must be oxidized (anode)

E°cell = -0.76 - 0.34 = -1.10 V < 0

Answer: No, non-spontaneous! ✗

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Applications of Electrochemical Series

1. Predicting Displacement Reactions

Rule: A metal can displace another metal from its salt solution if it has lower reduction potential.

Example: Can Zn displace Cu from CuSO₄?

• E°(Zn²⁺/Zn) = -0.76 V (lower)
• E°(Cu²⁺/Cu) = +0.34 V (higher)

Yes! Zn is a stronger reducing agent.

$$\text{Zn} + \text{CuSO}_4 \rightarrow \text{ZnSO}_4 + \text{Cu}$$

2. Reactivity of Metals with Acids

Rule: Metals with negative E° (below hydrogen) can displace H₂ from acids.

Can react with dilute acids (E° < 0):

• Zn, Fe, Al, Mg, Na, K (all below H in series)
• Example: $\text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2$

Cannot react with dilute acids (E° > 0):

• Cu, Ag, Au (all above H in series)
• Cu + HCl → No reaction

3. Galvanic Cell Design

Rule: Use two half-cells with large difference in E° for maximum EMF.

Best combination:

• Anode: Li/Li⁺ (E° = -3.05 V)
• Cathode: F₂/F⁻ (E° = +2.87 V)
• E°cell = 2.87 - (-3.05) = 5.92 V!

(Impractical due to reactivity, but shows the principle)

4. Corrosion Prevention

Rule: Metals with lower E° corrode preferentially (sacrificial protection).

For protecting iron ships (E°(Fe²⁺/Fe) = -0.44 V):

• Attach Zn blocks (E° = -0.76 V) → Zn corrodes instead
• Attach Mg blocks (E° = -2.37 V) → Even better protection!

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Comparing Oxidizing and Reducing Powers

Strongest Oxidizing Agents

Species	E° (V)	Common Use
F₂/F⁻	+2.87	Uranium enrichment
MnO₄⁻/Mn²⁺ (acidic)	+1.51	Titrations in labs
Cl₂/Cl⁻	+1.36	Water purification
Cr₂O₇²⁻/Cr³⁺ (acidic)	+1.33	Oxidizing agent

Strongest Reducing Agents

Species	E° (V)	Common Use
Li/Li⁺	-3.05	Batteries
K/K⁺	-2.93	Reduction reactions
Na/Na⁺	-2.71	Metallurgy
Mg/Mg²⁺	-2.37	Grignard reagents
Al/Al³⁺	-1.66	Thermite reactions

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Standard Electrode Potentials - Important Values

For JEE: Must Memorize

Half-Reaction	E° (V)
Halogens
F₂ + 2e⁻ → 2F⁻	+2.87
Cl₂ + 2e⁻ → 2Cl⁻	+1.36
Br₂ + 2e⁻ → 2Br⁻	+1.09
I₂ + 2e⁻ → 2I⁻	+0.54
Metals
Ag⁺ + e⁻ → Ag	+0.80
Cu²⁺ + 2e⁻ → Cu	+0.34
2H⁺ + 2e⁻ → H₂	0.00
Pb²⁺ + 2e⁻ → Pb	-0.13
Ni²⁺ + 2e⁻ → Ni	-0.25
Fe²⁺ + 2e⁻ → Fe	-0.44
Zn²⁺ + 2e⁻ → Zn	-0.76
Al³⁺ + 3e⁻ → Al	-1.66
Mg²⁺ + 2e⁻ → Mg	-2.37
Na⁺ + e⁻ → Na	-2.71
Li⁺ + e⁻ → Li	-3.05
Important Ions
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O	+1.51
Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O	+1.33
Fe³⁺ + e⁻ → Fe²⁺	+0.77

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Practice Problems

Level 1: JEE Main

Q1. Which is the strongest reducing agent?

• (a) Li (b) Na (c) K (d) Mg

Q2. Which metal cannot displace hydrogen from dilute HCl?

• (a) Zn (b) Fe (c) Cu (d) Mg

Q3. Arrange in order of increasing oxidizing power: Cl₂, Br₂, I₂, F₂

Q4. Calculate E°cell for:

$$\text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu}$$

Given: E°(Cu²⁺/Cu) = +0.34 V, E°(Fe²⁺/Fe) = -0.44 V

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Level 2: JEE Main/Advanced

Q5. Predict whether the following reactions are spontaneous:

• (a) $2\text{Ag} + \text{Zn}^{2+} \rightarrow 2\text{Ag}^+ + \text{Zn}$
• (b) $\text{Mg} + 2\text{Ag}^+ \rightarrow \text{Mg}^{2+} + 2\text{Ag}$

Q6. Which of the following can reduce Fe³⁺ to Fe²⁺?

• (a) Cu (b) Zn (c) I⁻ (d) Br⁻ Given: E°(Fe³⁺/Fe²⁺) = +0.77 V, E°(I₂/I⁻) = +0.54 V

Q7. In the electrochemical series, if E°(M²⁺/M) = -0.50 V:

• (a) Can M displace hydrogen from acids?
• (b) Can M reduce Cu²⁺ to Cu?
• (c) Can M reduce Ag⁺ to Ag?

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Level 3: JEE Advanced

Q8. A solution contains Fe²⁺, Cu²⁺, and Ag⁺ ions. What happens when zinc metal is added? Given: E°(Ag⁺/Ag) = +0.80 V, E°(Cu²⁺/Cu) = +0.34 V, E°(Fe³⁺/Fe²⁺) = +0.77 V, E°(Fe²⁺/Fe) = -0.44 V, E°(Zn²⁺/Zn) = -0.76 V

Q9. Three metals X, Y, Z have E° values -0.76 V, +0.34 V, and +0.80 V respectively. Which metal will:

• (a) Corrode most easily?
• (b) Be best for sacrificial protection of iron?
• (c) Be used as cathode in a cell with maximum EMF?

Q10. For the reaction:

$$a\text{MnO}_4^- + b\text{Fe}^{2+} + c\text{H}^+ \rightarrow d\text{Mn}^{2+} + e\text{Fe}^{3+} + f\text{H}_2\text{O}$$

Find a:b ratio and calculate E°cell. Given: E°(MnO₄⁻/Mn²⁺) = +1.51 V, E°(Fe³⁺/Fe²⁺) = +0.77 V

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Solutions to Practice Problems

A1. (a) Li - Most negative E° = -3.05 V

A2. (c) Cu - E° = +0.34 V (above hydrogen)

A3. I₂ < Br₂ < Cl₂ < F₂ (increasing E°)

A4. E°cell = 0.34 - (-0.44) = 0.78 V

A5.

• (a) E°cell = -0.76 - 0.80 = -1.56 V → Non-spontaneous
• (b) E°cell = 0.80 - (-2.37) = +3.17 V → Spontaneous

A6. (c) I⁻ - E°(I₂/I⁻) = +0.54 V < E°(Fe³⁺/Fe²⁺) = +0.77 V So I⁻ can reduce Fe³⁺

A7.

• (a) Yes (E° < 0)
• (b) Yes (E° = -0.50 < +0.34)
• (c) Yes (E° = -0.50 < +0.80)

A8. Zn will reduce all three: Ag⁺ → Ag, Cu²⁺ → Cu, and Fe²⁺ → Fe (all have higher E° than Zn)

A9.

• (a) X (-0.76 V, lowest/most negative)
• (b) X (needs lower E° than Fe)
• (c) Z (+0.80 V, highest)

A10. a:b = 1:5 (from electron balance); E°cell = 1.51 - 0.77 = 0.74 V

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Common Mistakes to Avoid

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Key Takeaways for JEE

Must Remember Concepts

1. Electrochemical series ranks elements by E° (reduction potential)
2. Higher E° → Better oxidizing agent (top of series)
3. Lower E° → Better reducing agent (bottom of series)
4. Spontaneity: E°cell > 0 (reduction > oxidation potential)
5. Metals below H₂ (E° < 0) react with dilute acids
6. E° is intensive - doesn’t multiply with stoichiometry

Quick Decision Tree

graph TD
    A[Compare two E° values] --> B{Higher E° species}
    B --> C[Gets REDUCED]
    B --> D[Acts as OXIDIZING agent]
    B --> E[Goes to CATHODE]

    A --> F{Lower E° species}
    F --> G[Gets OXIDIZED]
    F --> H[Acts as REDUCING agent]
    F --> I[Goes to ANODE]

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Real-Life Applications

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Related Topics

Within Electrochemistry

• Oxidation-Reduction — Foundation of redox reactions
• Electrochemical Cells — Building cells using E° values
• Nernst Equation — E° at non-standard conditions
• Batteries — Practical applications of electrode potentials

Cross-Chapter Connections

• Thermodynamics — ΔG° = -nFE° relationship
• Chemical Equilibrium — Connection via equilibrium constant
• Periodic Properties — Trends in E° values

Physics Connections

• Current Electricity — Potential difference concepts
• Electrostatics — Electric potential fundamentals

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