# Strong and weak acids and bases, and the pH scale

> IB Chemistry SL · IB
> Source: https://www.owlsprep.com/study/ib-chemistry-sl-u6-strong-and-weak-acids-and/

This guide covers definitions of strong and weak acids/bases, their key observable properties, pH scale calculations, and common question patterns you will encounter in IB Chemistry SL exams.

**Prerequisites:** [Basics of acid-base definitions (Bronsted-Lowry)](https://www.owlsprep.com/study/ib-chemistry-sl-u6-bronsted-lowry-acids-bases/); [Logarithm rules for chemistry calculations](https://www.owlsprep.com/study/ib-chemistry-sl-math-logarithms-for-chemistry/)

## Learning objectives

- Distinguish between strong and weak acids/bases in terms of degree of dissociation
- Relate acid/base strength to conductivity, reaction rate, and pH for equal concentrations
- Calculate pH of strong monoprotic acids and strong monobases from concentration
- Identify common examples of strong and weak acids and bases for IB exam contexts

## Defining Strong vs Weak Acids and Bases

**Strong and Weak Electrolytes** — Compounds that fully dissociate into ions in water are strong electrolytes, while those that only partially ionise are weak electrolytes. All strong acids and bases are strong electrolytes, all weak acids and bases are weak electrolytes.

$$HA_{(aq)} \rightarrow H^+_{(aq)} + A^-_{(aq)} \quad \text{(Strong acid, 100% dissociation)}$$

$$HA_{(aq)} \rightleftharpoons H^+_{(aq)} + A^-_{(aq)} \quad \text{(Weak acid, partial dissociation, equilibrium)}$$

> **Strong Acids Mnemonic**
>
> Strong acids list mnemonic: strong acids: HCl, HBr, HI, HNO₃, H₂SO₄ = Sulfuric, Hydroiodic, Hydrobromic, Nitric, Hydrochloric

**Worked example:** Classify each of the following as strong or weak acid: HCl, CH₃COOH, HNO₃, H₂CO₃

1. Step 1: Recall the 6 common strong acids specified for IB SL exams
2. Step 2: HCl (hydrochloric) and HNO₃ (nitric) are on the strong acid list, fully dissociated in water
3. Step 3: CH₃COOH (ethanoic acid, vinegar) and H₂CO₃ (carbonic acid, soda) are not on the list, so they are weak, partially dissociated at equilibrium

**Check your understanding**

Test your understanding of dissociation:

1. Which of the following 0.1 M solutions has the highest number of free ions?

   - Weak ethanoic acid
   - Strong hydrochloric acid
   - Weak ammonia
   - Ethanol (non-acid)

   *Why:* Strong HCl fully dissociates into H+ and Cl-, giving twice as many ions as the original undissociated acid molecules.

## Observable Properties to Distinguish Acid/Base Strength

For equal molar concentrations of acids or bases, you can tell strong and weak variants apart using three simple, measurable observations: electrical conductivity, rate of reaction with active metals or carbonates, and pH value.

| Property | Equal [ ] Strong Acid | Equal [ ] Weak Acid |
| --- | --- | --- |
| Electrical conductivity | High (high ion concentration) | Low (low ion concentration) |
| Reaction rate with Mg | Fast, vigorous bubbling | Slow, gentle bubbling |
| pH value | Low (higher H+ concentration) | Higher (lower H+ concentration) |

**Exam command terms**

IB exam questions often use specific command terms for this topic, these are the accepted marking scheme cues:

- **Describe a test to distinguish** — You must name a measurable observation difference, not just state definitions

- **Equal concentration** — This is a critical cue that strength, not concentration, is the variable being tested

**Worked example:** You have two unlabelled 0.1 mol dm⁻³ solutions of HCl and CH₃COOH. Describe one simple test to tell them apart, no pH meter allowed.

1. Step 1: Take 10 cm³ of each solution in separate identical beakers
2. Step 2: Add a 1 cm strip of magnesium ribbon to each beaker at the same time
3. Step 3: The solution that produces bubbles of hydrogen gas much faster is the strong HCl, the slower reaction is weak CH₃COOH

## The pH Scale and Calculations for Strong Species

**pH** — A base-10 logarithmic scale that quantifies the hydrogen ion concentration of an aqueous solution, ranging from 0 (very acidic) to 14 (very basic) at 25°C.

*Notation:* $pH = -\log_{10}[H^+_{(aq)}]$

$$[H^+] = 10^{-pH}$$

**Derivation:** Derive pH of 0.1 mol dm⁻³ HCl

*Starting from:* HCl is strong, fully dissociates: $HCl \rightarrow H^+ + Cl^-$

1. 1 mole of HCl produces 1 mole of H+ ions, so $[H^+] = 0.1$ mol dm⁻³
2. Substitute into pH formula: $pH = -\log_{10}(0.1)$
3. $-\log_{10}(10^{-1}) = 1$

*Conclusion:* pH of 0.1 M HCl is exactly 1

**Worked example:** Calculate the pH of 0.05 mol dm⁻³ NaOH, a strong monobasic base at 25°C, given $K_w = 1.0 \times 10^{-14}$

1. Step 1: NaOH fully dissociates, so $[OH^-] = 0.05$ mol dm⁻³
2. Step 2: Use $[H^+] = K_w / [OH^-] = 1.0 \times 10^{-14} / 0.05 = 2 \times 10^{-13}$
3. Step 3: $pH = -\log_{10}(2 \times 10^{-13}) = 12.7$

> **SL Exam Scope Note**
>
> IB SL only requires pH calculations for strong monoprotic acids and strong monobases. You do NOT need to calculate pH of weak acids at SL level.

## Reference of Common Strong and Weak Species

| Category | Common Strong Examples | Common Weak Examples |
| --- | --- | --- |
| Acids | HCl, H₂SO₄, HNO₃ | CH₃COOH, H₂CO₃, H₃PO₄ |
| Bases | NaOH, KOH, LiOH | NH₃, CH₃NH₂, Na₂CO₃ |

**Check your understanding**

Confirm you can identify species correctly:

1. Which of the following will have a pH > 7 at 0.1 M concentration?

   - Nitric acid
   - Sodium hydroxide
   - Hydrochloric acid
   - Ethanoic acid

   *Why:* NaOH is a strong base, so it will produce a pH of 13 at 0.1 M concentration.

## Common pitfalls

- **Wrong:** Confusing acid concentration with acid strength
  - Why it fails: Concentration is how many moles you have per litre, strength is how much of those moles dissociate, two independent variables
  - Correct: Always check if the question specifies equal concentration before comparing strength directly to pH
- **Wrong:** Stating weak acids have no H+ ions
  - Why it fails: Weak acids partially dissociate, they just have lower H+ concentration than equal concentration strong acids
  - Correct: Always write that weak acids form an equilibrium with partial ionisation, not zero ionisation
- **Wrong:** Trying to calculate pH of a weak acid at SL exam
  - Why it fails: Weak acid pH requires Ka values which are not given or tested at SL level
  - Correct: For SL, if you see a weak acid, you only need to estimate its pH is higher than equal concentration strong acid
- **Wrong:** Forgetting pH is a log scale
  - Why it fails: A difference of 1 pH unit means 10x difference in H+ concentration, not 2x
  - Correct: When asked how much more acidic pH 2 is than pH 4, the answer is 100x, not 2x
- **Wrong:** Including water as a strong acid
  - Why it fails: Water is a neutral substance with very low dissociation, not classified as strong or weak acid in IB syllabus
  - Correct: Only use the 5 strong acids (HCl, HBr, HI, HNO₃, H₂SO₄) named in the syllabus and group 1 hydroxides as strong bases for your answers

## Cheatsheet

| Species Type | Dissociation Behaviour | pH at 0.1 M 25°C |
| --- | --- | --- |
| Strong monoprotic acid | 100% dissociation | 1 |
| Weak monoprotic acid | Partial equilibrium dissociation | 2-3 |
| Strong monobase | 100% dissociation | 13 |
| Weak base | Partial equilibrium dissociation | 11-12 |

## What's next

Now that you have mastered the core differences between strong and weak acids, bases, and the pH scale, you are ready to move on to Ka calculations for weak acids (HL only) and buffer solutions, as well as exploring titration curves which build directly on these pH calculation skills. This topic is a foundational prerequisite for 15-20% of the marks in IB Chemistry SL Paper 2 Section B acid-base questions, so make sure you practice distinguishing strong and weak species in multiple choice questions to avoid easy lost marks on exam day.

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