# Electron transfer reactions (redox and electrochemistry)

> IB Chemistry SL · IB Diploma Programme Chemistry Standard Level
> Source: https://www.owlsprep.com/study/ib-chemistry-sl-u6-electron-transfer-reactions/

This module covers oxidation/reduction definitions, oxidation number rules, balancing redox half-equations, and core properties of voltaic and electrolytic cells for IB SL assessment.

**Prerequisites:** [Basic atomic structure and ion formation](https://www.owlsprep.com/study/ib-chemistry-sl-u2-atomic-structure/); [Chemical equations and conservation of mass](https://www.owlsprep.com/study/ib-chemistry-sl-u1-stoichiometric-relationships/)

## Learning objectives

- Define oxidation and reduction in terms of electron loss/gain and oxidation number change
- Assign oxidation numbers to elements in compounds and monatomic/polyatomic ions correctly
- Balance redox half-equations in acidic and neutral conditions per IB SL marking criteria
- Distinguish between core properties of voltaic and electrolytic cells

## Oxidation and Reduction Core Definitions

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.

**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.

> **Core Redox Mnemonic**
>
> OIL RIG: Oxidation Is Loss of electrons, Reduction Is Gain of electrons

**Worked example:** Identify which species is oxidized and which is reduced in the reaction: $\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)$

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

**Check your understanding**

1. Which process corresponds to reduction?

   - Gain of protons
   - Gain of electrons
   - Loss of electrons
   - Loss of neutrons

   *Why:* Reduction always refers to a net gain of negatively charged electrons, decreasing the species' oxidation number.

## Oxidation Number Assignment Rules

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 $\text{MnO}_4^-$

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. $$\text{Mn} + (-8) = -1 \implies \text{Mn} = +7$$

> **tip**
>
> Always write + or - for non-zero oxidation numbers to avoid losing marks in IB SL exams, do not write just '7' for +7.

## Balancing Redox Half-Equations

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.

**Exam command terms**

IB exam command terms for this topic have strict marking expectations:

- **Deduce the half-equation** — Show full balancing steps for atoms and charge, no skipped steps

- **State the oxidation number** — Only the final value is required, no working needed

**Worked example:** Balance the reduction half-equation for $\text{Cr}_2\text{O}_7^{2-}$ turning into $\text{Cr}^{3+}$ in acidic solution

1. Step 1: Balance chromium atoms: $\text{Cr}_2\text{O}_7^{2-} \rightarrow 2\text{Cr}^{3+}$
2. Step 2: Balance oxygen by adding 7 H₂O to the right side: $\text{Cr}_2\text{O}_7^{2-} \rightarrow 2\text{Cr}^{3+} + 7\text{H}_2\text{O}$
3. Step 3: Balance hydrogen by adding 14 H⁺ to the left side: $\text{Cr}_2\text{O}_7^{2-} + 14\text{H}^+ \rightarrow 2\text{Cr}^{3+} + 7\text{H}_2\text{O}$
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: $\text{Cr}_2\text{O}_7^{2-} + 14\text{H}^+ + 6e^- \rightarrow 2\text{Cr}^{3+} + 7\text{H}_2\text{O}$

*Calculator:* forbidden

## Voltaic and Electrolytic Cell Fundamentals

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.

**Comparing methods**

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

## Common pitfalls

- **Wrong:** Assigning oxygen an oxidation number of -1 in all oxygen-containing compounds
  - Why it fails: This rule only applies to peroxides; oxygen has an oxidation number of +2 in oxygen difluoride and -2 in all other common compounds
  - Correct: Check for peroxide O-O bonds or fluorine bonding first before applying default oxidation number values
- **Wrong:** Forgetting to balance oxygen atoms with H₂O when balancing acidic half-equations
  - Why it fails: Unbalanced oxygen leads to incorrect total charge calculations and lost partial marks
  - Correct: Follow the strict order: balance non-oxygen/hydrogen atoms first, then O with H₂O, then H with H⁺, then charge with electrons
- **Wrong:** Stating electrons flow from cathode to anode in a voltaic cell
  - Why it fails: Electrons are released when species are oxidized at the anode, so they always leave the anode first
  - Correct: Remember electrons flow from the negative terminal (anode) to the positive terminal (cathode) in all voltaic cells
- **Wrong:** Assigning hydrogen an oxidation number of +1 in metal hydrides
  - Why it fails: Metal hydrides contain the H⁻ anion, so hydrogen has an oxidation number of -1
  - Correct: If hydrogen is bonded to a group 1 or group 2 metal, its oxidation number is -1, not +1
- **Wrong:** Assuming oxidation occurs at the cathode for electrolytic cells
  - Why it fails: 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: Use the AN OX RED CAT mnemonic for all electrochemical cell questions to avoid mixing up electrode processes

## 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) |

## 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.

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