AQA A-Level Chemistry (7405)

## Learning Outcomes

- Identify the main features of titration curves for strong and weak acids and bases
- Explain why equivalence points can be at pH 7, above pH 7, or below pH 7
- Distinguish clearly between the equivalence point and the end-point
- Select a suitable indicator by matching its transition range to the steep part of a titration curve
- Interpret weak acid–strong base curves, including the buffer region and half-neutralisation point
- Apply mole calculations to titration data, including monoprotic and diprotic acids
- Recognise common practical and exam errors in titration curve questions

## AQA A-Level Chemistry (7405) Syllabus

For the AQA A-Level Chemistry (7405), you are expected to understand titration curves and indicator choice, with a focus on the following syllabus points:

- Know that a titration curve is a graph of pH against volume of titrant added
- Identify the typical shapes of curves for strong acid–strong base, weak acid–strong base, strong acid–weak base, and weak acid–weak base titrations
- Explain why the equivalence point is about pH 7 for a strong acid–strong base titration, above pH 7 for a weak acid–strong base titration, and below pH 7 for a strong acid–weak base titration
- Understand the difference between the equivalence point and the end-point
- Use the steep vertical section of a titration curve to choose a suitable indicator
- Know the transition ranges of methyl orange and phenolphthalein and when each is suitable
- Recognise the buffer region and half-neutralisation point on a weak acid–strong base titration curve
- Use titration data to calculate concentrations, including cases involving diprotic acids

## Test Your Knowledge

_Attempt these questions before reading this article. If you find some difficult or cannot remember the answers, look more closely at that area during your revision._

1. Why is the equivalence point not always at pH 7?
2. Which indicator is usually suitable for a weak acid–strong base titration, and why?
3. What is the difference between an end-point and an equivalence point?
4. What happens at the half-neutralisation point in a weak acid–strong base titration?
5. Why is a weak acid–weak base titration usually followed with a pH meter rather than a simple indicator?

## Introduction

Titration curves are one of the best synoptic topics in AQA Chemistry because they join together acid-base theory, equilibria, pH, indicators, practical skills, and calculations. A question may ask you to sketch a curve, explain its shape, choose an indicator, identify the equivalence point, or use the graph in a calculation.

This topic is often tested in a way that checks whether you can explain chemistry, not just recall shapes. You need to know why one curve has a buffer region, why another has an equivalence point below 7, and why an indicator works for one titration but fails for another.

> **Key Term: titration curve**
> A graph of pH against volume of titrant added during a titration.

### Equivalence point, end-point, and the basic logic of a titration curve

In an acid-base titration, one solution is added from a burette into another solution of known volume. As the reaction progresses, the pH changes. Plotting pH against volume added gives the titration curve.

The most important point on the curve is the **equivalence point**. This is the exact stoichiometric point where the acid and base have reacted in the required mole ratio. For a monoprotic acid reacting with a hydroxide base, this means moles of H+ supplied equal moles of OH− supplied.

> **Key Term: equivalence point**
> The stoichiometric point in a titration where the acid and base have reacted in the exact required molar ratio.

Students often mix this up with the **end-point**. They are not identical. The equivalence point is the true chemical point of neutralisation. The end-point is what you observe experimentally, usually from an indicator colour change. A good indicator is chosen so that the end-point is as close as possible to the equivalence point.

> **Key Term: end-point**
> The observed point in a titration at which the indicator changes colour.

Why do curves have a steep section? Near the equivalence point, a very small extra volume of titrant changes which substance is in excess. Before equivalence, one reagent controls the pH. Just after equivalence, the other does. That switch can cause a very large pH change over a tiny volume interval.

For strong acid–strong base titrations, this pH jump is very large because both species are fully dissociated. For titrations involving weak acids or weak bases, the jump is smaller because partial dissociation and hydrolysis effects matter more.

A common exam trap is writing that "the end-point is where the solution becomes neutral." That is only safely true for some titrations, and even then the wording is weak. Better wording is: the end-point is where the indicator changes colour, and it should be close to the equivalence point.

> **Test Tip:** If a question asks you to explain indicator choice, do not just name the indicator. State that its transition range lies within the steep vertical section of the titration curve, so the colour change occurs close to the equivalence point.

### The main titration curves and why they have different shapes

You need to recognise four standard cases.

#### Strong acid–strong base

This curve starts at very low pH because the acid is fully dissociated. As strong base is added, the pH rises gradually at first, then very sharply near the equivalence point, then levels off at high pH once base is in excess.

The equivalence point is about **pH 7** at 25 °C. This is because the salt formed does not hydrolyse significantly. The solution contains ions such as Na+ and Cl−, which do not react with water enough to shift pH away from 7.

The vertical section is very steep and covers a wide pH range, so several indicators can work.

#### Weak acid–strong base

This curve starts at a higher pH than a strong acid of the same concentration because the weak acid is only partially dissociated. As the strong base is added, a mixture of weak acid and its conjugate base forms. That creates a **buffer region**, so the pH rises more slowly at first.

> **Key Term: buffer region**
> The part of a weak acid–strong base titration curve where both a weak acid and its conjugate base are present, so pH changes only gradually.

Near equivalence, the curve still becomes steep, but the vertical section is less dramatic than for strong acid–strong base. The equivalence point is **above pH 7**. This is a very common exam question.

Why above 7? At equivalence, the weak acid has been converted into its salt, containing the conjugate base A−. That ion reacts with water:

A− + H2O ⇌ HA + OH−

This hydrolysis produces OH−, so the solution is alkaline.

Students often say "because a strong base was used." That is not enough. At equivalence, the base added has been exactly used up. The solution is alkaline because of the hydrolysis of the conjugate base of the weak acid, not because excess strong base is present.

This curve also contains the **half-neutralisation point**. This is where half the weak acid has been converted into its conjugate base. At this point, [HA] = [A−], so from the Henderson–Hasselbalch relationship:

pH = pKa + log([A−]/[HA])

Since [A−]/[HA] = 1, log(1) = 0, so:

pH = pKa

> **Key Term: half-neutralisation point**
> The point in a weak acid–strong base titration where half the weak acid has been neutralised, so [HA] = [A−] and pH = pKa.

This is useful because you can read **pKa** directly from the graph. First find the equivalence volume. Then take half of that volume. The pH at that half-volume is the pKa.

> **Key Term: pKa**
> A value defined by pKa = -log10Ka; for a weak acid–strong base titration, pH = pKa at the half-neutralisation point.

#### Strong acid–weak base

This curve starts at low pH because the acid is strong. As weak base is added, the pH rises, but the vertical section is smaller than in a strong acid–strong base titration.

The equivalence point is **below pH 7**. Again, the explanation depends on hydrolysis. At equivalence, the solution contains the conjugate acid of the weak base. For example, if ammonia is titrated with hydrochloric acid, the salt contains NH4+:

NH4+ + H2O ⇌ NH3 + H3O+

This produces H3O+, so the solution is acidic.

A weak answer says "because the acid is strong." A strong answer explains that the conjugate acid of the weak base hydrolyses water and makes the solution acidic.

#### Weak acid–weak base

This is the least suitable for visual indicators. The curve has no sharp vertical section because the pH changes too gradually around equivalence. That means no simple indicator gives a clean, reliable end-point. A pH meter or probe is much better.

Also, the equivalence pH is not fixed at 7. It depends on the relative strengths of the weak acid and weak base.

> **Key Term: hydrolysis**
> The reaction of an ion with water that changes the pH of the solution.

A good comparison summary is:

- strong acid–strong base: equivalence about pH 7
- weak acid–strong base: equivalence above pH 7
- strong acid–weak base: equivalence below pH 7
- weak acid–weak base: no sharp vertical section
  > **Test Tip:** When explaining why equivalence pH is above or below 7, always refer to the hydrolysis of the salt ion formed. This is the step many students miss.

### Choosing indicators from a titration curve

An indicator is suitable only if its transition range lies entirely within the steep part of the titration curve. If it changes colour outside that region, the colour change happens too early or too late, so the measured titre is inaccurate.

> **Key Term: indicator transition range**
> The narrow pH range over which an indicator changes colour.

Two indicators appear again and again:

> **Key Term: methyl orange**
> An acid-base indicator with a transition range of about pH 3.1 to 4.4; red in acid and yellow in alkali.
>
> **Key Term: phenolphthalein**
> An acid-base indicator with a transition range of about pH 8.3 to 10.0; colourless in acid and pink in alkali.

Use them like this:

- **Strong acid–strong base:** methyl orange or phenolphthalein usually both work, because the vertical section is very large.
- **Weak acid–strong base:** phenolphthalein is suitable; methyl orange changes colour too early.
- **Strong acid–weak base:** methyl orange is suitable; phenolphthalein changes colour too late.
- **Weak acid–weak base:** neither is usually suitable; use a pH meter.

Why is this true? Because for weak acid–strong base, the steep section sits in the alkaline region, so phenolphthalein matches it. For strong acid–weak base, the steep section sits in the acidic region, so methyl orange matches it.

Worked example:

A titration curve has an equivalence point at pH 8.8, and the steep section runs roughly from pH 7.5 to pH 10.5.

Step 1: Identify which indicator ranges fit that vertical region.

- Methyl orange: pH 3.1 to 4.4
- Phenolphthalein: pH 8.3 to 10.0

Step 2: Compare the ranges.

- Methyl orange lies well below the steep region, so it would change colour too early.
- Phenolphthalein lies inside the steep region, so it would change colour near the equivalence point.

Answer: **phenolphthalein** is suitable.

A common mistake is to choose the indicator based only on the pH at the equivalence point. That is incomplete. You must match the whole transition range to the steep vertical section, not just one pH value.

### Calculations, practical method, and common errors

Titration questions often combine curves with mole calculations. The core formula is:

moles = concentration × volume in dm3

If volume is in cm3, divide by 1000 first.

Worked example:

A 25.0 cm3 sample of NaOH is neutralised by 20.0 cm3 of 0.150 mol dm−3 HCl. Find the concentration of the NaOH.

Step 1: Write the equation.

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

Step 2: Find moles of HCl.

Volume of HCl = 20.0 cm3 = 0.0200 dm3

Moles of HCl = 0.150 × 0.0200 = 0.00300 mol

Step 3: Use the mole ratio.

The ratio HCl : NaOH is 1 : 1, so moles of NaOH = 0.00300 mol

Step 4: Find concentration of NaOH.

Volume of NaOH = 25.0 cm3 = 0.0250 dm3

Concentration = 0.00300 / 0.0250 = 0.120 mol dm−3

Answer: **0.120 mol dm−3**

Now a diprotic example:

32.0 cm3 of KOH is neutralised by 18.0 cm3 of 0.400 mol dm−3 H2SO4. Find the concentration of KOH.

Step 1: Equation.

H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2H2O(l)

Step 2: Moles of H2SO4.

18.0 cm3 = 0.0180 dm3

Moles = 0.400 × 0.0180 = 0.00720 mol

Step 3: Use the mole ratio.

1 mol H2SO4 reacts with 2 mol KOH

Moles of KOH = 2 × 0.00720 = 0.0144 mol

Step 4: Concentration of KOH.

32.0 cm3 = 0.0320 dm3

Concentration = 0.0144 / 0.0320 = 0.450 mol dm−3

Answer: **0.450 mol dm−3**

Students often make these errors:

- forgetting to convert cm3 to dm3
- missing the stoichiometric ratio
- assuming all acids are monoprotic
- using end-point and equivalence point as if they are identical
- confusing acid strength with concentration

That last point matters in curve interpretation. A weak acid can still be concentrated, and a strong acid can be dilute. Strength means degree of dissociation, not amount present.

In required practical work, reliability and resolution matter. A rough titre is used first to estimate the end-point. Then accurate titres are collected. Concordant titres are close together, usually within about 0.10 cm3. When using a pH probe to draw a full curve, use larger volume additions at first, then much smaller additions near equivalence, because that is where the pH changes fastest.

Good practical points include:

- rinse the burette with the titrant
- rinse the pipette with the solution it will deliver
- read the burette at eye level
- add titrant dropwise near the end-point
- repeat until concordant results are obtained
- calibrate the pH probe if one is used
- stir throughout so the pH is uniform

## Key Point Checklist

_This article has covered the following key knowledge points:_

- A titration curve is a graph of pH against volume of titrant added
- The equivalence point is the stoichiometric point where acid and base have reacted in the correct mole ratio
- The end-point is the observed indicator colour change and should be close to the equivalence point
- Strong acid–strong base titrations have an equivalence point at about pH 7
- Weak acid–strong base titrations have an equivalence point above pH 7 because the conjugate base hydrolyses water to produce OH−
- Strong acid–weak base titrations have an equivalence point below pH 7 because the conjugate acid hydrolyses water to produce H3O+
- Weak acid–strong base titrations show a buffer region before equivalence
- At the half-neutralisation point in a weak acid–strong base titration, pH = pKa
- Methyl orange is usually suitable for strong acid–weak base titrations
- Phenolphthalein is usually suitable for weak acid–strong base titrations
- Either methyl orange or phenolphthalein is usually suitable for strong acid–strong base titrations
- Weak acid–weak base titrations are usually followed with a pH meter because there is no sharp vertical section

## Key Terms and Concepts

- titration curve
- equivalence point
- end-point
- buffer region
- half-neutralisation point
- pKa
- hydrolysis
- indicator transition range
- methyl orange
- phenolphthalein


Acids and bases - Titration curves and indicators