Chemistry (0620) | Acid–base titrations

🧪 Acid–Base Titrations: Measuring Exactly What You Have!

Hello! This chapter is all about a super important practical technique in chemistry called titration. It sounds complicated, but it's really just a very careful, precise way of doing a neutralisation reaction to find out exactly how much acid or alkali you have in a solution.

Mastering titrations is vital because it links your practical lab skills (using apparatus correctly) with your calculation skills (stoichiometry and moles). Ready to become a chemical detective? Let’s dive in!

Key Takeaway from the Introduction

Titration is a method of quantitative analysis used to find the unknown concentration of a solution.

1. Understanding Acid–Base Titration

What is Titration? (Volumetric Analysis)

An acid-base titration is a method of analysis that allows us to accurately determine the volume of one solution (the acid) required to react completely with a measured volume of another solution (the alkali).

The core chemical reaction involved is neutralisation:

Acid + Alkali (Base) \(\rightarrow\) Salt + Water

Remember the ionic equation for neutralisation (Syllabus 7.1 Core 8):
H⁺(aq) + OH^-(aq) \(\rightarrow\) H₂O(l)

Key Terminology in Titration

• Titrant: The solution added from the burette (usually the one whose concentration is accurately known).
• Analyte (or Titrate): The solution in the conical flask (the one whose concentration is unknown).
• Titre: The volume of the titrant added from the burette.
• End-Point: The point during the titration where the indicator changes colour. We assume this is the same as the equivalence point for the calculation.
• Equivalence Point: The theoretical point where the moles of acid have exactly reacted with the moles of alkali, according to the stoichiometric ratio of the balanced equation.

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2. The Essential Titration Apparatus (12.2 Core 1)

To perform a titration accurately, you need specialised glassware designed for precise volume measurement.

(a) The Burette

The burette is a long, graduated glass tube with a tap (stopcock) at the bottom.

• Purpose: To deliver variable, precise volumes of the titrant solution.
• Precision: Burettes are usually $50 \text{ cm}^3$ and allow measurements to be taken to two decimal places (e.g., $25.50 \text{ cm}^3$), with the smallest division being $0.1 \text{ cm}^3$.
• How to Read: Always read the volume from the bottom of the meniscus (the curve of the liquid surface) and read vertically at eye level to avoid parallax error.

(b) The Volumetric Pipette

A volumetric pipette is designed to measure and deliver a single, fixed, highly accurate volume (e.g., $10.0 \text{ cm}^3$ or $25.0 \text{ cm}^3$).

• Purpose: To measure an exact, fixed volume of the analyte solution into the conical flask.
• Safety Tip: You must always use a pipette filler (never your mouth!) to suck the liquid up to the mark.

(c) The Conical Flask and White Tile

• Conical Flask: Used to hold the analyte and indicator. Its shape allows you to swirl the solution vigorously without spilling, ensuring the acid and alkali mix rapidly.
• White Tile: Placed under the flask to make the indicator colour change much clearer and easier to spot.

(d) The Suitable Indicator (12.2 Core 1 & 2)

The indicator is crucial because it tells you exactly when the reaction is complete (the end-point).

• Role: To change colour abruptly over a small range of pH values, signalling neutralisation.
• Examples:
  • Methyl Orange: Red in acid, Yellow in alkali.
  • Phenolphthalein (often used): Colourless in acid, Pink/Magenta in alkali.

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3. Performing the Practical Titration

Titration requires extreme care to ensure the volumes recorded are as accurate as possible. Follow these steps:

Step 1: Preparation

1. Rinse the burette with a small amount of the titrant (the solution going in it) to ensure no residual water or previous chemical contaminates the solution. Fill the burette above the zero mark and then run some through the jet (the tip) to ensure there are no air bubbles. Record the initial reading.
2. Use the volumetric pipette and filler to accurately measure a fixed volume (e.g., $25.0 \text{ cm}^3$) of the analyte solution into the conical flask.
3. Add 2 or 3 drops of a suitable indicator to the conical flask. (Adding too much indicator can affect the end-point accuracy).

Step 2: The Rough Titration (Finding the Approximate Titre)

The first titration is done quickly to get an idea of the volume needed.

• Run the titrant from the burette into the flask, swirling constantly, until the indicator changes colour permanently.
• This reading is called the rough titre. It is usually not used in calculations but helps you work quickly and accurately in the next step.

Step 3: The Accurate Titrations (Finding the End-Point)

Now we aim for high precision.

1. Refill the burette and record the new initial reading.
2. Add the titrant quickly until you are about $2 \text{ cm}^3$ short of the rough titre volume.
3. Slow down dramatically! Add the titrant **drop by drop**, swirling constantly.
4. When a single drop causes the indicator to change colour permanently (the end-point), immediately close the tap.
5. Record the final burette reading.

Step 4: Recording and Calculating the Mean Titre

Repeat the accurate titration until you obtain concordant results (volumes that agree closely, typically within $0.10 \text{ cm}^3$ of each other).

• Discard the rough titre.
• Calculate the average (mean) of the concordant titres. This is your average volume of titrant used in the reaction.

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4. Using Titration Data to Calculate Concentration (3.3 Supplement 6)

This section is key for Extended candidates. The titration data allows you to calculate the moles involved, and thus the unknown concentration.

Prerequisites: Concentration and Volume

Remember the relationship involving moles (\(n\)), concentration (\(C\)), and volume (\(V\)):

$$n \text{ (mol)} = C \text{ (mol/dm}^3) \times V \text{ (dm}^3)$$

Crucial Unit Check: Volume must always be in cubic decimetres ($dm^3$) for concentration calculations. Since your pipette and burette volumes are typically in $cm^3$:

$$V \text{ (dm}^3) = \frac{V \text{ (cm}^3)}{1000}$$

Step-by-Step Calculation Guide

Let's assume you are titrating Sulphuric acid (\(H_2SO_4\)) with Sodium Hydroxide (\(NaOH\)).

Step 1: Write the Balanced Chemical Equation

This step provides the essential mole ratio.

$$H_2SO_4 + 2NaOH \rightarrow Na_2SO_4 + 2H_2O$$

Mole Ratio: 1 mole of acid reacts with 2 moles of alkali.

Step 2: Calculate the Moles of the KNOWN Solution

Use the measured volume and the known concentration of the titrant (e.g., $NaOH$).

$$n_{NaOH} = C_{NaOH} \times V_{NaOH} \text{ (in dm}^3)$$

Step 3: Use the Mole Ratio to Find Moles of the UNKNOWN Solution

Using the mole ratio from Step 1 (1:2 ratio):

$$n_{H_2SO_4} = n_{NaOH} \times \frac{1}{2}$$

(If the ratio was 1:1, the moles would be equal.)

Step 4: Calculate the Concentration of the UNKNOWN Solution

You now have the moles of the unknown acid (\(n_{H_2SO_4}\)) and you know the exact volume you measured into the flask with the pipette (\(V_{H_2SO_4}\)).

$$C_{H_2SO_4} = \frac{n_{H_2SO_4}}{V_{H_2SO_4} \text{ (in dm}^3)}$$

Did You Know?

Titrations are not just used in labs! In manufacturing and medicine, titrations are routinely used to ensure quality control, such as checking the exact acidity of soft drinks or determining the precise amount of active ingredient in a medication tablet. The accuracy matters in the real world!