Required practical 9: pH changes for weak acid with strong base and strong acid with weak base

Introduction to Required Practical 9

Welcome to one of the most visual and data-driven practicals in your A-level Chemistry course! In Required Practical 9, we step beyond simple "colour-change" titrations to look at the actual "story" of a neutralisation reaction. By using a pH probe, we can track exactly how the acidity or alkalinity of a solution changes as we add a titrant.

This practical focuses on two specific scenarios: reacting a weak acid with a strong base and a strong acid with a weak base. Understanding these curves is essential for mastering Paper 1 and Paper 3, as they reveal the secrets of buffer solutions and indicator choices.

The Core Equipment: The pH Meter

While universal indicator is great for a rough estimate, it isn't precise enough for scientific curves. We use a pH meter (or pH probe) which provides a digital reading of the \( [H^+] \) concentration.

Calibration: Why and How?

Don't worry if your pH meter seems a bit "moody" at first—they are sensitive instruments! Before starting, you must calibrate the probe because pH readings can "drift" over time.
1. Rinse the probe with distilled water.
2. Place it into a buffer solution of a known pH (usually \( pH = 4.0 \), \( 7.0 \), or \( 10.0 \)).
3. Adjust the meter until it matches the known value.
4. Repeat with at least one other buffer to ensure accuracy across the whole scale.

Quick Tip: Always rinse the probe between different solutions to avoid cross-contamination!

Step-by-Step Procedure

Whether you are reacting \( CH_3COOH \) (weak acid) with \( NaOH \) (strong base) or \( HCl \) (strong acid) with \( NH_3 \) (weak base), the steps are very similar:

1. Set up the burette: Fill it with your standard solution (the base or acid of known concentration).
2. Prepare the reaction vessel: Use a pipette to measure a fixed volume (e.g., \( 25.0 \text{ cm}^3 \)) of your other reactant into a beaker.
3. Initial reading: Place the pH probe into the beaker and record the initial \( pH \).
4. Add the titrant: Add the solution from the burette in small portions (e.g., \( 2.0 \text{ cm}^3 \)) and stir thoroughly using a magnetic stirrer or glass rod.
5. Record: Record the \( pH \) after each addition.
6. Approach the end-point: As the \( pH \) starts to change more rapidly, reduce the additions to \( 0.2 \text{ cm}^3 \) or even dropwise. This is crucial for capturing the "vertical section" of the curve!
7. Post-equivalence: Continue adding the titrant until it is in excess and the \( pH \) levels off again.

Analyzing the pH Curves

When you plot your results (Volume added on the x-axis, \( pH \) on the y-axis), you will see distinct shapes based on the strength of your reactants.

1. Weak Acid (\( WA \)) with Strong Base (\( SB \))

Example: Ethanoic acid and Sodium Hydroxide
- Starting pH: Around \( pH = 3 \).
- The Buffer Region: The \( pH \) rises slowly at first because a buffer is formed (a mixture of the weak acid and its salt).
- Equivalence Point: This occurs at a pH > 7. This is a common mistake—students often think neutralisation always means \( pH = 7 \), but for a weak acid/strong base, the resulting salt is slightly basic.
- Ending pH: Around \( pH = 13 \)- \( 14 \).

2. Strong Acid (\( SA \)) with Weak Base (\( WB \))

Example: Hydrochloric acid and Ammonia
- Starting pH: Around \( pH = 1 \).
- The Shape: The \( pH \) rises steadily.
- Equivalence Point: This occurs at a pH < 7. The salt formed in this reaction is slightly acidic.
- Ending pH: Around \( pH = 10 \)- \( 11 \) (weak bases don't reach the high \( pH \) levels of \( NaOH \)).

Key Concept: The Equivalence Point and Indicators

The equivalence point is the exact volume where the amount of base added is stoichiometrically equal to the amount of acid. On your graph, this is the mid-point of the vertical section.

Choosing the Right Indicator

An indicator is only useful if its colour change occurs entirely within the vertical section of the \( pH \) curve.
- Phenolphthalein: Changes colour between \( pH = 8.3 \) and \( 10.0 \). Perfect for WA/SB titrations.
- Methyl Orange: Changes colour between \( pH = 3.1 \) and \( 4.4 \). Perfect for SA/WB titrations.
- The "Golden Rule": Never use an indicator for a Weak Acid/Weak Base titration because there is no sharp vertical section!

Using the Curve to find \( pK_a \)

One of the coolest things about a Weak Acid/Strong Base curve is the Half-Neutralisation Point. This is the point where exactly half of the acid has been turned into its salt.
At this specific point:
\( [HA] = [A^-] \)
If we look at the \( K_a \) expression: \( K_a = \frac{[H^+][A^-]}{[HA]} \)
The \( [A^-] \) and \( [HA] \) cancel out, leaving:
\( K_a = [H^+] \) which means \( pK_a = pH \).
So, just find the volume at the equivalence point, divide it by two, and read the \( pH \) at that volume on your graph. That is your \( pK_a \)!

Safety and Hazards

Even though we call them "weak," acids like ethanoic acid can still be irritants.
- Corrosive/Irritant: Concentrated acids and bases can damage skin and eyes. Always wear safety goggles and lab coats.
- Spills: If you spill any chemical, wash with plenty of water immediately.
- Safe Handling: Use a funnel to fill the burette, but always remove it before starting your titration to prevent extra drops falling in.

Summary Checklist

- Did you calibrate the pH probe? (Necessary for accuracy).
- Did you add titrant dropwise near the end-point? (Necessary to see the vertical section).
- Is the equivalence point > 7 or < 7? (Depends on whether the base or acid is "stronger").
- Does your indicator match the vertical section? (Essential for a valid titration).

Note: For more on how to calculate pH from concentrations, see the chapter on 3.1.12 Acids and Bases. To see how this compares to other practical techniques, check out Required Practical 1.