Faraday’s constant is a core idea in IB Chemistry Topic 9 (Redox Processes) and Topic 19 (HL). It links charge, electrons, and moles together, making it essential for calculating mass deposited during electrolysis, gas volumes produced, and the quantity of electric charge needed for chemical change. Once you understand how Faraday’s constant works, electrolysis problems become straightforward.
What Is Faraday’s Constant?
Faraday’s constant (F) is the amount of electric charge carried by one mole of electrons.
Its value is:
F = 96485 C mol⁻¹
(rounded to 96500 C mol⁻¹ in many IB calculations)
This number connects electrical charge and chemical change.
Why Faraday’s Constant Has This Value
Faraday’s constant is based on two fundamental quantities:
1. Charge of a single electron (e):
1.602 × 10⁻¹⁹ C
2. Avogadro’s number (NA):
6.022 × 10²³ mol⁻¹
Multiplying them gives the charge carried by one mole of electrons:
F = e × NA
F ≈ (1.602 × 10⁻¹⁹) × (6.022 × 10²³)
F ≈ 96485 C mol⁻¹
This makes Faraday’s constant a physical bridge between atomic-scale charges and real-world electric current.
Why Faraday’s Constant Matters in Chemistry
Electrolysis involves moving electrons from an external power supply into ions in solution or molten form. To calculate how much product forms, you must know:
- How many electrons are transferred
- How much charge passes through the system
Faraday’s constant directly connects charge and moles.
Faraday’s Laws of Electrolysis
Faraday’s constant appears in both laws.
Faraday’s First Law
The mass of substance produced is proportional to the total charge passed.
Faraday’s Second Law
Different substances deposit amounts proportional to their molar masses and number of electrons transferred.
These laws enable all electrolysis calculations.
Key Formulae Using Faraday’s Constant
1. Total Charge
Q = I × t
Where:
Q = charge (C)
I = current (A)
t = time (s)
2. Moles of electrons
moles of e⁻ = Q / F
3. Moles of product
If a reaction requires n electrons:
moles of product = (Q / F) / n
4. Mass of product formed
mass = moles × molar mass
These formulas appear frequently in IB Paper 2 and Paper 3.
Example Calculation (IB Style)
Question:
A current of 2.0 A is passed through molten MgCl₂ for 30 minutes. How many grams of Mg form?
Step 1: Calculate charge
t = 30 min = 1800 s
Q = I × t = 2.0 × 1800 = 3600 C
Step 2: Moles of electrons
moles e⁻ = Q / F = 3600 / 96485 ≈ 0.0373 mol e⁻
Step 3: Determine electrons per Mg atom
Mg²⁺ + 2e⁻ → Mg
So n = 2
moles Mg = 0.0373 / 2 = 0.0187 mol
Step 4: Mass formed
mass = 0.0187 × 24.3 ≈ 0.454 g Mg
This is a typical IB electrolysis calculation.
Faraday’s Constant in Electrochemical Cells
Although Faraday’s constant is mostly used in electrolysis calculations, it is also important in galvanic cells:
- It connects electric charge to ΔG (in HL thermodynamics)
- It helps relate cell potential to maximum work
The formula appears in HL Topic 19:
ΔG = −nFE°
This connects:
- Gibbs free energy
- Number of electrons
- Faraday’s constant
- Cell potential
Common IB Misconceptions
“Faraday’s constant is the same as Avogadro’s number.”
False. Faraday’s constant is the charge carried by a mole of electrons.
“Faraday’s constant changes with the reaction.”
It is a physical constant—always the same.
“Faraday’s constant measures current.”
It measures charge per mole, not current.
FAQs
Why does electrolysis require Faraday’s constant?
Because electrolysis depends on the transfer of electrons, and Faraday’s constant links electrons to measurable charge.
Does Faraday’s constant ever change?
No. It is a universal constant based on fundamental particles.
Why is it sometimes rounded to 96500?
For simpler calculations; the IB often accepts either value.
Conclusion
Faraday’s constant (96485 C mol⁻¹) is the amount of charge carried by one mole of electrons. It underpins all electrolysis calculations, connects chemical reactions to electrical charge, and appears throughout redox chemistry. Mastering Faraday’s constant makes quantitative electrochemistry much more intuitive in IB Chemistry.
