Neutralisation (OCR A-Level Chemistry A): Revision Notes

Neutralisation

Introduction to pH titration curves

Acid-base titrations are analytical techniques used to determine the concentration of an acid or base by monitoring neutralisation reactions. During a titration, the pH changes as one solution is gradually added to another. These pH changes can be monitored and analyzed to provide valuable information about the solutions involved.

Using a pH meter

A pH meter uses an electrode placed in the solution being tested. The electrode connects to a digital meter that displays the pH reading with high precision.

Standard procedure for acid-base titration:

1. Add a measured volume of acid to a conical flask using a pipette
2. Place the pH meter electrode in the flask, ensuring it is fully immersed
3. Add the aqueous base from a burette in $1~\text{cm}^3$ portions
4. After each addition, swirl the flask contents and record both the pH and total volume of base added
5. Continue adding base in $1~\text{cm}^3$ portions until the pH begins to change more rapidly
6. Once rapid pH changes are observed, add the base dropwise (smaller volumes) to capture the vertical section accurately
7. Return to $1~\text{cm}^3$ additions once the pH is changing slowly again
8. Continue until excess base has been added and the pH stabilizes at a high value

The results are then plotted as a graph of pH (y-axis) against total volume of base added (x-axis), producing a characteristic S-shaped titration curve.

Key features of pH titration curves

When analyzing a pH titration curve for a strong acid-strong base titration, several distinct regions can be identified:

Initial excess acid region: When base is first added, the acid is present in large excess. The pH increases very slowly as small amounts of acid are neutralized. The pH remains low (around 1-3) during this stage.

Vertical section: As more base is added, the acid concentration decreases significantly. Eventually, the pH increases very rapidly with the addition of only a few drops of base. This produces the characteristic steep, vertical section of the curve. During this section, the pH can change by 8-10 units with the addition of less than $0.1~\text{cm}^3$ of base.

Excess base region: After passing through the equivalence point, the base is now in excess. Further additions of base cause the pH to increase slowly again as the basic solution becomes more concentrated. The pH stabilizes at a high value (around 12-13).

Adding acid to a base

pH titration curves can also be constructed by adding acid to base, rather than adding base to acid. The shape of the curve is essentially inverted - it starts at high pH and decreases to low pH.

Acid-base indicators

The end point

Indicators provide a visual method for detecting when a neutralisation reaction is complete. An acid-base indicator is a weak acid (represented as HA) where the molecular form and its conjugate base (A⁻) display distinctly different colors.

For methyl orange, a commonly used indicator:

• The weak acid form, HA(aq), appears red
• The conjugate base form, A⁻(aq), appears yellow

The indicator exists in equilibrium between these two forms:

$\text{HA}(aq) \rightleftharpoons \text{A}^-(aq) + \text{H}^+(aq)$

Explaining indicator color changes

Indicators function through equilibrium shifts driven by changes in hydrogen ion concentration. Understanding Le Chatelier's principle helps explain these color changes.

When strong base is added to strong acid containing methyl orange:

Initially, the high concentration of H⁺ ions drives the equilibrium position far to the left, favoring the acid form HA. The solution appears red.

When OH⁻(aq) ions are added:

1. OH⁻ ions react with H⁺ ions in the solution: $\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$
2. This removes H⁺ ions, causing the weak acid HA to dissociate further
3. The equilibrium shifts to the right
4. The color changes progressively from red through orange (at the end point) to yellow as the equilibrium position shifts

When acid is added to base containing methyl orange:

The process reverses:

1. H⁺(aq) ions react with the conjugate base form A⁻(aq)
2. This removes A⁻ from the equilibrium
3. The equilibrium shifts to the left
4. The color changes from yellow through orange (at the end point) to red

Sensitivity of the end point

Different indicators undergo color changes over different pH ranges, determined by their acid dissociation constant ($K_a$) values.

At the end point of an indicator:

• $[\text{HA}(aq)] = [\text{A}^-(aq)]$
• Therefore: $K_a = [\text{H}^+(aq)]$
• This means: $pK_a = pH$

The pH ranges and colors for several common indicators are shown below:

Indicator	pH range
Thymol blue	1.2-2.8
Bromophenol blue	3.6-4.6
Methyl red	4.4-6.2
Metacresol purple	7.4-9.0
Indigo carmine	11.4-13.0

Choosing the indicator

Selecting an appropriate indicator is crucial for accurate titration results. The indicator's color change range must coincide with the vertical section of the pH titration curve.

Ideal scenario: The end point and equivalence point should coincide. In practice, they may differ slightly (by one or two drops, equivalent to $0.05-0.10~\text{cm}^3$), but this difference is negligible in analytical terms.

Strong acid-strong base titration (e.g., HCl and NaOH):

• Large vertical section spanning pH 3-11
• Both methyl orange (pH 3-4) and phenolphthalein (pH 8-10) are suitable
• Either indicator will change color sharply at the equivalence point

Weak acid-strong base titration (e.g., CH₃COOH and NaOH):

• Smaller vertical section on the basic side (approximately pH 7-11)
• Phenolphthalein is suitable (pH 8-10 falls within the vertical section)
• Methyl orange is NOT suitable (pH 3-4 is too acidic and not in the vertical section)

Strong acid-weak base titration (e.g., HCl and NH₃):

• Smaller vertical section on the acidic side (approximately pH 3-7)
• Methyl orange is suitable (pH 3-4 falls within the vertical section)
• Phenolphthalein is NOT suitable (pH 8-10 is too basic and not in the vertical section)