Unit 1: Preparation of Buffers - Subjective Questions

Definition: A buffer solution is a solution that resists significant changes in pH when small quantities of a strong acid or strong base are added.

Composition:

• An acidic buffer contains a weak acid and its salt with a strong base, such as acetic acid and sodium acetate.
• A basic buffer contains a weak base and its salt with a strong acid, such as ammonium hydroxide and ammonium chloride.

Action of an acidic buffer:

• Added hydrogen ions are consumed by the conjugate base: .
• Added hydroxide ions react with the weak acid: .

Thus, the ratio of the weak acid to its conjugate base changes only slightly, so the pH remains nearly constant.

Molarity () is the number of moles of solute dissolved per litre of solution.

Since:

Substituting this relationship:

Therefore:

This formula is used to calculate the quantity of a pure solute needed for buffer preparation.

Normality () is the number of gram-equivalents of solute present per litre of solution.

The relationship between normality and molarity is:

where is the -factor or equivalence factor.

Differences:

• Molarity is based on moles, whereas normality is based on equivalents.
• Molarity is generally independent of the type of reaction, whereas normality depends on the reaction.
• For a monoprotic acid such as HCl, .
• For sulfuric acid in a complete neutralization reaction, because one mole provides two equivalents of hydrogen ions.

Procedure:

1. Calculate the required amounts of the weak acid and its conjugate salt using the Henderson–Hasselbalch equation.
2. Accurately weigh the required chemicals using a calibrated balance.
3. Dissolve each component separately in a small quantity of distilled water.
4. Transfer the solutions quantitatively into a volumetric flask.
5. Rinse the beakers and add the washings to the flask.
6. Add distilled water up to approximately the calibration mark.
7. Mix thoroughly and allow the solution to reach room temperature.
8. Make the final volume exactly up to the mark and mix by inverting the flask several times.
9. Label the solution with its composition, concentration, pH, date, and expiry information.

The final pH should be checked using a properly calibrated pH meter.

For a weak acid buffer containing a weak acid and its conjugate base , the dissociation equation is:

The acid dissociation constant is:

Rearranging and taking negative logarithms gives:

For a buffer prepared from a salt and a weak acid, concentrations may often be replaced by the amounts or molar concentrations of the conjugate base and weak acid.

The equation shows that the pH depends mainly on the of the weak acid and the ratio of conjugate base to weak acid. Maximum buffering occurs when , because then .

A basic buffer contains a weak base and its conjugate acid , usually supplied by a salt such as ammonium chloride. The equilibrium is:

The equation for a basic buffer is:

Since:

an alternative form is:

where refers to the conjugate acid of the weak base.

• is the concentration of the weak base.
• is the concentration of its conjugate acid.
• The equation is used to calculate the pH or the proportions of components required for buffer preparation.

Use the formula:

Given:

• Molar mass
• Volume

Therefore:

Thus, 13.61 g of sodium acetate trihydrate should be accurately weighed, dissolved in distilled water, and diluted to a final volume of .

The required mass is calculated using:

where:

Laboratory procedure:

1. Determine the molar mass and -factor of the compound.
2. Calculate its equivalent weight.
3. Calculate the mass required for the desired normality and volume.
4. Accurately weigh the solid.
5. Dissolve it in a small volume of distilled water.
6. Transfer the solution to a volumetric flask.
7. Wash the container and add the washings to the flask.
8. Dilute to the calibration mark with distilled water.
9. Stopper and mix thoroughly.
10. Label the prepared solution with its normality and preparation details.

Normality is commonly used in acid–base and redox titrations, whereas molarity is widely used for general solution preparation and quantitative calculations.

A volumetric flask is designed to contain one highly accurate volume at a specified temperature.

Importance in buffer preparation:

• It permits accurate adjustment of the final solution volume.
• Its narrow neck allows precise reading of the meniscus.
• It minimizes volume errors during preparation.
• It is suitable for preparing standard and analytical solutions.

A measuring cylinder is less suitable because:

• Its graduations are less precise.
• The wider shape makes accurate meniscus reading more difficult.
• The final volume may have a larger uncertainty.
• Small volume errors can alter the concentrations and pH of the buffer.

Therefore, a volumetric flask should be used whenever an accurately known concentration is required.

During dilution, the number of moles of solute transferred from the stock solution remains constant. Therefore:

where:

• is the concentration of the stock solution.
• is the volume of stock solution required.
• is the desired final concentration.
• is the final volume.

The required stock volume is calculated as:

Application: The calculated volume of concentrated stock is measured accurately with a pipette, transferred to a volumetric flask, and diluted to the final mark with distilled water. The flask is then mixed thoroughly. The same relationship can be applied to molar or normal concentrations, provided the concentration units are consistent.

Buffer capacity is the amount of strong acid or strong base required to change the pH of one litre of a buffer solution by one pH unit.

Buffer capacity is affected by:

• Total concentration: A concentrated buffer has a greater capacity than a dilute buffer with the same pH.
• Ratio of components: Capacity is greatest when the concentrations of the weak acid and conjugate base are approximately equal.
• pH relative to : A buffer is most effective within approximately one pH unit above or below its .
• Volume: The total amount of acid or base required depends on the volume of the buffer.
• Temperature and ionic strength: These can influence dissociation equilibria and the measured pH.

A buffer may have an appropriate pH but poor capacity if its total concentration is too low.

A stock solution is a concentrated solution prepared accurately and stored for subsequent dilution or use. A working buffer solution is a solution prepared at the concentration and pH required for an experiment.

Differences:

• Stock solutions have higher concentrations than working solutions.
• Working solutions are usually prepared by direct weighing or dilution of a stock solution.
• Stock solutions reduce repeated weighing and preparation errors.
• Working solutions are used directly in laboratory procedures.
• Both solutions must be labeled with concentration, pH, date of preparation, and storage conditions.

The dilution equation is used to prepare a working buffer from a stock solution.

Method:

1. Calibrate the pH meter using suitable standard buffer solutions, commonly at two or three pH values.
2. Rinse the electrode with distilled water and gently blot it dry.
3. Transfer a portion of the prepared buffer to a clean beaker.
4. Immerse the electrode without touching the sides or bottom.
5. Stir gently and wait for a stable reading.
6. If adjustment is required, add small amounts of dilute acid or base while stirring.
7. Allow the reading to stabilize after each addition.
8. Transfer the adjusted solution to a volumetric flask and restore the final volume if necessary.
9. Record the final pH and temperature.

Large additions of acid or base should be avoided because they change the buffer composition and concentration.

The neutralization reaction is:

The stoichiometric ratio is .

Using the relation for equal equivalents:

Given:

• Acetic acid concentration
• Acetic acid volume
• Sodium hydroxide concentration

Therefore, 50 mL of sodium hydroxide is required for complete neutralization.

Accuracy at each stage is essential because the buffer pH depends on the concentrations and ratio of its components.

• Accurate weighing: Ensures that the correct number of moles or equivalents is used.
• Complete dissolution: Prevents undissolved material from causing an incorrect concentration.
• Quantitative transfer: Rinsing the original container ensures that all solute reaches the volumetric flask.
• Correct volume adjustment: The final volume must be adjusted exactly to the calibration mark; overfilling produces a solution that is too dilute.
• Thorough mixing: Ensures uniform concentration throughout the solution.
• Temperature control: Volume and pH can vary with temperature, so measurements should be made under controlled conditions.

Errors in any of these steps can change the buffer pH, buffer capacity, and reproducibility of experimental results.

For an acidic buffer, the Henderson–Hasselbalch equation is:

Rearranging:

Let the total buffer concentration be and the final volume be . If , then:

The required moles are:

Therefore, the masses are:

These calculated masses are weighed, dissolved, combined, and diluted to the required final volume.

Both types resist pH changes through reversible acid–base reactions.

Equivalent weight is the mass of a substance that supplies, reacts with, or replaces one mole of reactive units, such as one mole of hydrogen ions in an acid–base reaction.

The formula is:

For sulfuric acid:

• Molar mass of approximately.
• In complete neutralization, one mole of sulfuric acid supplies two moles of .
• Therefore, the -factor is 2.

Thus, the equivalent weight of sulfuric acid for complete neutralization is 49 g per equivalent. Its normality is related to its molarity by .

Common errors and precautions:

• Incorrect calculations: Recheck molar mass, -factor, volume, and concentration units.
• Impure or hydrated chemicals: Use the correct chemical form and account for water of crystallization.
• Inaccurate weighing: Use a calibrated analytical balance and avoid loss during transfer.
• Incomplete dissolution: Ensure that all solid dissolves before final dilution.
• Overfilling the volumetric flask: Add water slowly near the mark and use a dropper for the final adjustment.
• Parallax error: Read the bottom of the meniscus at eye level.
• Poor mixing: Invert the stoppered volumetric flask several times.
• Uncalibrated pH meter: Calibrate the instrument before measuring pH.
• Contamination: Use clean glassware, distilled water, and uncontaminated reagents.
• Temperature variation: Prepare and measure solutions at a controlled temperature.

These precautions improve the accuracy, stability, and reproducibility of the buffer.