Electron transfer reactions (redox and electrochemistry)

IB Chemistry SL· Topic 9: Redox processes (SL content)· 25 min read

1. Oxidation and Reduction Core Definitions★★☆☆☆⏱ 5 min

All redox reactions involve transfer of one or more electrons between reacting species. Early definitions described oxidation as reaction with oxygen, but the modern electron transfer definition is universally applied for IB SL questions.

📘 Definition

Redox reaction

A paired reaction where one species loses electrons (oxidized) and a second species gains electrons (reduced), with no net creation or destruction of electrons.

📐 Worked Example

Identify which species is oxidized and which is reduced in the reaction: $Zn (s) + Cu^{2 +} (a q) \to Zn^{2 +} (a q) + Cu (s)$

1
Count electrons on each side for zinc: Neutral Zn metal becomes Zn²⁺, losing 2 electrons
2
Count electrons on each side for copper: Cu²⁺ becomes neutral Cu metal, gaining 2 electrons
3
Final classification: Zn is oxidized, Cu²⁺ is reduced

✓ Quick check

Which process corresponds to reduction?
- Gain of protons
- Gain of electrons
- Loss of electrons
- Loss of neutrons
Reveal answer
Gain of electrons —
Reduction always refers to a net gain of negatively charged electrons, decreasing the species' oxidation number.

2. Oxidation Number Assignment Rules★★★☆☆⏱ 6 min

Oxidation numbers let you track electron transfer without drawing full Lewis structures. IB SL exam questions almost always include at least one 1-mark question asking you to assign the oxidation number of a specified element in a polyatomic ion or compound.

Species type	Default oxidation number
Uncombined element (e.g. O₂, Fe)	0
Group 1 metal in compounds	+1
Group 2 metal in compounds	+2
Fluorine in all compounds	-1
Oxygen (most compounds)	-2
Hydrogen (most compounds)	+1

📐 Worked Example

Calculate the oxidation number of manganese in the permanganate ion $MnO_{4}^{-}$

1
Sum of oxidation numbers for all atoms in the ion equals the overall charge of -1
2
Assign default oxidation number of -2 to each of the 4 oxygen atoms, total O contribution = $4 \times - 2 = - 8$
3
$\text{Mn} + (-8) = -1 \implies \text{Mn} = +7$

3. Balancing Redox Half-Equations★★★★☆⏱ 7 min

🚫 No Calculator

Half-equation balancing is a 3-4 mark common IB SL question, and marks are deducted for missing steps or unbalanced charge. All half-equations you will be asked to balance at SL use acidic conditions, so you will not need to add OH⁻ ions.

📐 Worked Example

Balance the reduction half-equation for $Cr_{2} O_{7}^{2 -}$ turning into $Cr^{3 +}$ in acidic solution

1
Step 1: Balance chromium atoms: $Cr_{2} O_{7}^{2 -} \to 2 Cr^{3 +}$
2
Step 2: Balance oxygen by adding 7 H₂O to the right side: $Cr_{2} O_{7}^{2 -} \to 2 Cr^{3 +} + 7 H_{2} O$
3
Step 3: Balance hydrogen by adding 14 H⁺ to the left side: $Cr_{2} O_{7}^{2 -} + 14 H^{+} \to 2 Cr^{3 +} + 7 H_{2} O$
4
Step 4: Balance total charge: Left side total charge = -2 +14 = +12, right side total charge = 2*(+3) = +6. Add 6 electrons to left to balance: $Cr_{2} O_{7}^{2 -} + 14 H^{+} + 6 e^{-} \to 2 Cr^{3 +} + 7 H_{2} O$

4. Voltaic and Electrolytic Cell Fundamentals★★★☆☆⏱ 7 min

Electrochemical cells are divided into two core categories you must distinguish for IB SL: spontaneous voltaic cells that generate electricity, and non-spontaneous electrolytic cells that use external electricity to drive redox reactions.

Methods compared

Key differences between the two cell types:

Voltaic Cell

Spontaneous redox reaction, no external power source, anode is negative, cathode is positive, converts chemical energy to electrical energy

+ Pros: Generates usable electricity

− Cons: Only runs for thermodynamically favoured reactions

Electrolytic Cell

Non-spontaneous reaction, requires external DC power source, anode is positive, cathode is negative, converts electrical energy to chemical energy

+ Pros: Can drive unfavourable reactions like metal extraction

− Cons: Requires continuous energy input

📐 Worked Example

Identify the direction of electron flow in a standard zinc-copper voltaic cell

1
Zinc is more reactive than copper, so zinc metal is oxidized at the anode
2
Electrons are released from the zinc anode, which has a negative charge
3
Electrons flow through the external wire from the zinc anode directly to the copper cathode, where Cu²⁺ ions are reduced

5. Common Pitfalls

Wrong move:

Assigning oxygen an oxidation number of -1 in all oxygen-containing compounds

Why:

This rule only applies to peroxides; oxygen has an oxidation number of +2 in oxygen difluoride and -2 in all other common compounds

Correct move:

Check for peroxide O-O bonds or fluorine bonding first before applying default oxidation number values

Wrong move:

Forgetting to balance oxygen atoms with H₂O when balancing acidic half-equations

Why:

Unbalanced oxygen leads to incorrect total charge calculations and lost partial marks

Correct move:

Follow the strict order: balance non-oxygen/hydrogen atoms first, then O with H₂O, then H with H⁺, then charge with electrons

Wrong move:

Stating electrons flow from cathode to anode in a voltaic cell

Why:

Electrons are released when species are oxidized at the anode, so they always leave the anode first

Correct move:

Remember electrons flow from the negative terminal (anode) to the positive terminal (cathode) in all voltaic cells

Wrong move:

Assigning hydrogen an oxidation number of +1 in metal hydrides

Why:

Metal hydrides contain the H⁻ anion, so hydrogen has an oxidation number of -1

Correct move:

If hydrogen is bonded to a group 1 or group 2 metal, its oxidation number is -1, not +1

Wrong move:

Assuming oxidation occurs at the cathode for electrolytic cells

Why:

The 'AN OX RED CAT' rule applies to all cell types: oxidation always happens at the anode, reduction at the cathode regardless of cell charge polarity

Correct move:

Use the AN OX RED CAT mnemonic for all electrochemical cell questions to avoid mixing up electrode processes

6. Quick Reference Cheatsheet

Species / Rule	Default Oxidation Number	Common Exception
Uncombined element	0	None
Group 1 metal	+1	None
Group 2 metal	+2	None
Fluorine	-1	None
Oxygen	-2	Peroxides (-1), OF₂ (+2)
Hydrogen	+1	Metal hydrides (-1)

When this came up on past exams

AI-estimated based on syllabus patterns — cross-check with official past papers for accuracy. Use only as revision-focus signals.

2024 · Paper 2
Balancing acidic redox half-equations
2023 · Paper 1
Oxidation number identification in ions
2022 · Paper 2
Voltaic cell component labelling

What's Next

Mastering electron transfer redox reactions is the foundational skill you will use for all subsequent electrochemistry and organic oxidation state questions in your IB SL exam. This content is frequently tested alongside reaction spontaneity calculations and redox titration practical questions, which carry up to 8% of your total SL exam marks. You will now build on this base to learn how to calculate standard cell potentials, predict spontaneous reaction directions, and interpret results from common redox titration practicals such as iron(II) and permanganate titrations. Complete the end-of-spoke practice questions to lock in your half-equation balancing skills before progressing.