Standard electrode potential—often written as E°—is essential in IB Chemistry Topic 9 (Redox Processes). It helps you determine which species are stronger oxidizing or reducing agents, predict the direction of redox reactions, and calculate the overall cell potential of voltaic (galvanic) cells. Understanding E° values is crucial for analyzing redox behavior and electrochemical cells.
What Is Standard Electrode Potential?
Standard electrode potential (E°) is the voltage of a half-cell measured relative to the standard hydrogen electrode (SHE) under standard conditions.
It reflects how easily a species gains electrons (is reduced).
Standard conditions:
- 298 K (25°C)
- 100 kPa pressure (for gases)
- 1.0 mol dm⁻³ ion concentrations
- Electrodes in pure form
All E° values listed in the IB Data Booklet are measured under these conditions.
Why Use the Standard Hydrogen Electrode (SHE)?
To compare reduction potentials, we need a universal reference.
The standard hydrogen electrode is defined as:
- H₂(g) at 100 kPa
- 1.0 mol dm⁻³ H⁺(aq)
- Platinum electrode
Its potential is set to 0.00 V by definition.
All other electrode potentials are measured against this reference.
What an E° Value Tells You
A species with a more positive E°:
- Gains electrons more easily
- Is a stronger oxidizing agent
- Is more likely to be reduced
A species with a more negative E°:
- Gains electrons less easily
- Is a stronger reducing agent
- Is more likely to be oxidized
This allows prediction of spontaneous reaction direction.
Reading the IB Data Booklet Table
IB always presents reduction half-equations:
Example:
Cu²⁺ + 2e⁻ → Cu(s) E° = +0.34 V
This means copper(II) ions readily accept electrons.
Another example:
Mg²⁺ + 2e⁻ → Mg(s) E° = −2.37 V
Magnesium metal is a strong reducing agent because its E° is very negative.
Predicting Spontaneous Redox Reactions
A redox reaction is spontaneous when:
E°cell = E°(reduction) − E°(oxidation) > 0
Steps:
- Identify the half-equation with the more positive E°.
• This one undergoes reduction. - The other half-equation undergoes oxidation.
- Combine them and calculate E°cell.
- If E°cell is positive → reaction is spontaneous.
This rule is heavily tested in Paper 2.
Example of Predicting Spontaneity
Given:
Ag⁺ + e⁻ → Ag E° = +0.80 V
Zn²⁺ + 2e⁻ → Zn E° = −0.76 V
Silver has the more positive E°, so it is reduced.
Oxidation (reverse zinc equation):
Zn → Zn²⁺ + 2e⁻
E°cell = 0.80 − (−0.76) = +1.56 V → spontaneous reaction.
Standard Electrode Potentials in Electrochemical Cells
In a voltaic cell:
- Electrons flow from lower E° to higher E°
- The more positive E° is the cathode (reduction)
- The more negative E° is the anode (oxidation)
This explains why magnesium corrodes easily and copper does not.
Why E° Values Matter in Chemistry
Standard electrode potentials allow you to:
- Rank oxidizing and reducing agents
- Predict feasibility of reactions
- Calculate cell potentials
- Analyze corrosion and metal reactivity
- Understand electrolysis and plating
- Solve IB exam problems involving redox tables
They form a bridge between chemical reactivity and electrical energy.
Common IB Misunderstandings
“More positive E° means stronger reducing agent.”
Incorrect. More positive means stronger oxidizing agent.
“E° becomes more positive when you reverse a reaction.”
No. E° does not change sign when multiplied or divided. It only flips when the reaction direction changes.
“You always subtract larger from smaller.”
Incorrect. Use the formula:
E°cell = E°(cathode) – E°(anode)
FAQs
Why are all E° values reduction potentials?
For consistency. IB presents all half-reactions in reduction form to avoid confusion.
Does changing concentration affect E°?
Yes, but then it becomes a non-standard potential. Standard conditions must be used for E° values.
Why can’t E° predict reaction rate?
It predicts thermodynamic feasibility, not kinetics.
Conclusion
Standard electrode potential measures the tendency of a species to be reduced under standard conditions. The more positive the E°, the stronger the oxidizing agent. E° values allow you to predict reaction direction, determine spontaneous processes, and calculate cell potentials. Mastering E° is essential for understanding electrochemical cells in IB Chemistry.
