# Proton transfer reactions (acids and bases)

> Chemistry · IB
> Source: https://www.owlsprep.com/study/ib-chemistry-sl-u6-proton-transfer-reactions/

We will cover Bronsted-Lowry definitions, conjugate pairs, balanced proton transfer equations, and common reaction examples for IB SL Chemistry.

**Prerequisites:** [Basic chemical equation balancing](https://www.owlsprep.com/study/ib-chemistry-sl-u1-chemical-equations/); [Ion formation and ionic charge rules](https://www.owlsprep.com/study/ib-chemistry-sl-u2-ionic-bonding/)

## Learning objectives

- Define Bronsted-Lowry acids and bases as proton donors and acceptors
- Identify conjugate acid-base pairs in proton transfer reactions
- Write balanced half-equations for proton transfer processes
- Classify common lab compounds as Bronsted-Lowry acids or bases

## Bronsted-Lowry Acid and Base Definitions

The Bronsted-Lowry model of acid-base chemistry centers entirely on the transfer of a single proton (H⁺ ion) between two species, rather than the older Arrhenius definition that only applies to aqueous solutions. This model works for all solvent systems, including non-aqueous reactions.

**Bronsted-Lowry Acid** — Proton donor, releases one H⁺ ion during reaction

*Notation:* HA

*Example:* Hydrochloric acid HCl, ethanoic acid CH₃COOH

**Bronsted-Lowry Base** — Proton acceptor, uses a lone electron pair to bind the incoming H⁺

*Notation:* B

*Example:* Ammonia NH₃, hydroxide ion OH⁻

**Worked example:** Classify the following species as Bronsted-Lowry acid, base, or both: H₂O, HNO₃, NH₂⁻

1. Step 1: Check if the species can donate an H⁺
2. $$HNO_3 \rightarrow H^+ + NO_3^-$$
3. HNO₃ can only donate a proton, so it is a Bronsted-Lowry acid.
4. Step 2: Check if the species can accept an H⁺
5. $$NH_2^- + H^+ \rightarrow NH_3$$
6. NH₂⁻ can only accept a proton, so it is a Bronsted-Lowry base.
7. Step 3: H₂O can both donate and accept protons, so it is amphiprotic, acting as acid or base depending on the reaction partner.

> **info**
>
> Amphiprotic species can act as acid or base, the most common example you will see in IB SL questions is water.

**Exam command terms**

IB exam command terms for this topic have specific expectations:

- **Identify** — Only name the acid/base, no justification required

- **Classify** — State if the species is a proton donor or acceptor to support your answer

## Conjugate Acid-Base Pairs

When an acid donates a proton, the remaining species is its conjugate base, which can accept a proton to reform the original acid. When a base accepts a proton, the resulting species is its conjugate acid, which can donate a proton to reform the original base. The two species in a pair differ by exactly one H⁺ unit.

$$HA + B \rightleftharpoons A^- + HB^+$$

**Worked example:** Identify all conjugate acid-base pairs in the reaction between ethanoic acid and water: CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺

1. Step 1: Find the species that loses an H⁺ on the left side
2. CH₃COOH loses an H⁺ to become CH₃COO⁻, so this is the acid → conjugate base pair.
3. Step 2: Find the species that gains an H⁺ on the left side
4. H₂O gains an H⁺ to become H₃O⁺, so this is the base → conjugate acid pair.
5. Step 3: Confirm each pair differs by exactly one H⁺, no other atom or charge changes.

**Check your understanding**

Test your understanding:

1. Which of the following is a valid conjugate pair?

   - H₂O and OH⁻
   - HCl and H₂O
   - NH₃ and O²⁻
   - CH₃COOH and H₂SO₄

   *Why:* H₂O donates a proton to form OH⁻, so they differ by exactly one H⁺.

## Writing Balanced Proton Transfer Equations

All proton transfer reactions must balance for atoms and total charge on both sides of the equation. The total charge on the reactant side must equal the total charge on the product side, as no electrons are lost or gained in the proton transfer process.

**Derivation:** Derive the full balanced equation for the reaction between hydrochloric acid and ammonia

*Starting from:* Half equation 1: HCl donates a proton

1. $$HCl \rightarrow H^+ + Cl^-$$
2. Half equation 2: NH₃ accepts the proton
3. $$NH_3 + H^+ \rightarrow NH_4^+$$
4. Add the two half equations together, cancel the free H⁺ term on both sides

*Conclusion:* Full balanced equation: $HCl + NH_3 \rightarrow NH_4^+ + Cl^-$

> **tip**
>
> If you are unsure of the charge of the conjugate species, calculate total charge on each side to confirm it is equal.

## Common Proton Transfer Reaction Examples

IB SL Chemistry exam questions almost exclusively use a small set of common proton transfer reactions, including neutralisation between strong acids and strong bases, reaction of weak acids with carbonates, and aqueous dissociation of weak bases.

| Reaction Type | Example Equation | Acid | Base |
| --- | --- | --- | --- |
| Strong acid + strong base | $HCl + NaOH \rightarrow NaCl + H_2O$ | HCl | OH⁻ |
| Weak acid + carbonate | $2CH_3COOH + Na_2CO_3 \rightarrow 2CH_3COONa + H_2O + CO_2$ | CH₃COOH | CO₃²⁻ |
| Ammonia in water | $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$ | H₂O | NH₃ |

## Common pitfalls

- **Wrong:** Identifying H₂O as always an acid or always a base
  - Why it fails: Water is amphiprotic, so its role depends entirely on the other reaction partner
  - Correct: Check if water loses an H⁺ (acts as acid) or gains an H⁺ (acts as base) in the specific reaction given
- **Wrong:** Counting conjugate pairs that differ by more than one H⁺
  - Why it fails: Conjugate pairs must differ by exactly one proton, no more no less
  - Correct: Count the number of H atoms and the charge difference between the two species to confirm they only differ by one H⁺
- **Wrong:** Forgetting to balance total charge across the proton transfer equation
  - Why it fails: Unbalanced charge is an automatic mark deduction in IB P2 questions
  - Correct: Sum all charges on reactant side and product side before finalising your equation to confirm they are equal
- **Wrong:** Using the Arrhenius definition to classify bases in non-aqueous reactions
  - Why it fails: Arrhenius bases only produce OH⁻ in water, which does not apply to reactions in other solvents
  - Correct: Always use the Bronsted-Lowry proton donor/acceptor definition for all IB SL acid-base classification questions
- **Wrong:** Classifying the conjugate base of a strong acid as a strong base
  - Why it fails: Conjugate bases of strong acids are extremely weak, almost neutral species
  - Correct: Remember: strong acid → very weak conjugate base, weak acid → stronger conjugate base

## Cheatsheet

| Term | Definition | Key Rule |
| --- | --- | --- |
| Bronsted-Lowry Acid | Proton donor | Loses 1 H⁺ |
| Bronsted-Lowry Base | Proton acceptor | Gains 1 H⁺ |
| Conjugate Pair | Two species differing by 1 H⁺ | Charge differs by +1 between base and its conjugate acid |
| Amphiprotic | Can donate or accept H⁺ | Most common example = H₂O |

## What's next

Now that you have mastered the core proton transfer model of Bronsted-Lowry acids and bases, you are ready to explore the properties of strong and weak acids, including pH calculations, dissociation constants, and neutralisation reaction stoichiometry. These concepts are heavily weighted in Paper 1 multiple choice and Paper 2 structured response questions, and build directly on the conjugate pair identification skills you practiced in this guide. Mastering proton transfer first will make all subsequent acid-base topics far easier to learn, and will help you avoid common mistakes that cost marks on exam day.

---

From [OwlsPrep](https://www.owlsprep.com) — free study guides for A-Level, IB, AP and IGCSE, written against the official syllabus. Canonical page: https://www.owlsprep.com/study/ib-chemistry-sl-u6-proton-transfer-reactions/