Acid-Base Titrations - Expanded Notes

I. Core Concepts

Acid-base titrations are quantitative analytical techniques used to determine the concentration of an unknown acid or base solution by reacting it with a solution of known concentration. This reaction, a neutralization reaction, involves the transfer of protons (H⁺) from an acid to a base. The point at which the acid and base have completely reacted is called the equivalence point. In practice, we observe the endpoint, which is indicated by a color change of an added indicator. While ideally the equivalence point and endpoint are the same, slight differences can occur due to the indicator's properties.

Key Terminology:

• Analyte (Titrand): The solution with the unknown concentration (the substance being analyzed).
• Titrant (Reagent): The solution of known concentration added to the analyte.
• Equivalence Point: The theoretical point in the titration where the moles of acid and base are stoichiometrically equal.
• Endpoint: The point in the titration where the indicator changes color, signaling the approximate equivalence point.
• Indicator: A substance that changes color at a specific pH range, allowing visual detection of the endpoint. The choice of indicator is crucial; it should change color near the expected pH at the equivalence point.

II. Titration Procedure: A Detailed Look

1. Preparation of the Analyte:

   • Accurately measure a known volume of the analyte solution using a pipette or burette and transfer it to a clean Erlenmeyer flask.
   • Add a few drops of an appropriate indicator. The amount of indicator should be minimal to avoid affecting the results.

2. Filling the Burette:

   • Rinse the burette thoroughly with the titrant solution to ensure there is no contamination.
   • Fill the burette with the titrant solution above the zero mark. Carefully drain the titrant to precisely zero the burette, eliminating any air bubbles in the tip. Record the initial burette reading.

3. Titration:

   • Slowly add the titrant from the burette to the analyte in the flask, swirling constantly to ensure thorough mixing.
   • As the endpoint is approached, the addition of titrant should be slowed to a drop-wise addition. This ensures the endpoint is not overshot.
   • Observe the color change of the indicator carefully. The color change should be gradual initially, then rapid as the endpoint is approached.

4. Endpoint Determination:

   • The endpoint is reached when the color change of the indicator persists for at least 30 seconds. This indicates that the reaction is complete.
   • Record the final burette reading.

5. Calculations:

   • Calculate the volume of titrant used by subtracting the initial burette reading from the final burette reading.
   • Use the stoichiometry of the balanced chemical equation to determine the moles of analyte.
   • Calculate the concentration of the analyte using the known volume and the calculated moles.

III. Indicator Selection: Crucial Considerations

Indicator selection is paramount for accurate titration results. A good indicator exhibits a sharp, readily observable color change near the equivalence point. The indicator's pKa (the negative logarithm of its acid dissociation constant) should be close to the pH at the equivalence point. This ensures a rapid and distinct color change, minimizing error. Different indicators are appropriate for different titrations (strong acid-strong base, weak acid-strong base, etc.). The choice depends on the pH at the equivalence point which depends on the strength of the acid and base involved.

IV. Types of Titrations

There are several types of acid-base titrations, categorized based on the strength of the acid and base involved:

• Strong Acid-Strong Base Titrations: These titrations involve a strong acid (like HCl) and a strong base (like NaOH). The equivalence point occurs at pH 7.
• Weak Acid-Strong Base Titrations: The equivalence point will be above pH 7 (basic).
• Strong Acid-Weak Base Titrations: The equivalence point will be below pH 7 (acidic).
• Weak Acid-Weak Base Titrations: These are less common because the equivalence point is difficult to determine precisely.

V. Calculations: A Worked Example (Strong Acid-Strong Base)

Let's revisit the HCl/NaOH example from the original notes:

Given:

• 50.00 mL of HCl solution (analyte)
• 0.1000 M NaOH solution (titrant)
• 10.00 mL of NaOH required to reach the endpoint

Balanced Equation: HCl(aq) + NaOH(aq) → H₂O(l) + NaCl(aq)

Calculations:

1. Moles of NaOH: (0.1000 mol/L) * (0.01000 L) = 0.001000 mol NaOH

2. Moles of HCl: From the balanced equation, the mole ratio of HCl to NaOH is 1:1. Therefore, 0.001000 mol HCl reacted.

3. Molarity of HCl: (0.001000 mol HCl) / (0.05000 L) = 0.02000 M HCl*

VI. Sources of Error

Several factors can introduce error in acid-base titrations:

• Indicator error: The endpoint may not exactly coincide with the equivalence point.
• Parallax error: Incorrect reading of the burette due to improper eye level.
• Improper mixing: Incomplete mixing of the analyte and titrant can lead to inaccurate results.
• Contamination: Impurities in the solutions or glassware can affect the results.

By carefully following the procedure and understanding potential sources of error, accurate and reliable results can be obtained in acid-base titrations. Remember that meticulous technique and careful observation are crucial for success in this type of quantitative analysis.