Lab Report: Determination of Dissociation Constants (Ka) of Weak Acids

Abstract:

The aim of this experiment was to determine the dissociation constants (Ka) of weak acids, namely ethanoic acid, chloroethanoic acid, and dichloroethanoic acid. This was achieved by titrating 0.10M sodium hydroxide (NaOH) solution with 20.0cm³ of each acid and monitoring the pH of the resulting solutions. The Ka values were calculated using the titration data and compared with literature values.

Introduction:

Weak acids are substances that partially dissociate in solution, producing both hydrogen ions (H⁺) and their conjugate bases (X^-). The equilibrium constant for this dissociation is called the dissociation constant (Ka), which is a measure of the strength of the acid. In this experiment, we aimed to determine the Ka values for three different weak acids: ethanoic acid (CH₃COOH), chloroethanoic acid, and dichloroethanoic acid.

The pH meter was calibrated using a buffer solution with a known pH value to ensure accurate pH measurements during the experiment.

Materials and Methods:

Materials:

• 0.10M Sodium Hydroxide (NaOH) Solution
• Ethanoic Acid (CH₃COOH)
• Chloroethanoic Acid
• Dichloroethanoic Acid
• Phenolphthalein Indicator
• Distilled Water

Methods:

1. The pH meter was calibrated using a buffer solution with a known pH.
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2. 20.0cm³ of 0.10M ethanoic acid was pipetted into a conical flask.
3. 0.10M sodium hydroxide solution was titrated into the flask using phenolphthalein as an indicator until the solution turned pink.
4. Another 20.0cm³ of the same ethanoic acid solution was added to the flask and thoroughly mixed.
5. The pH of the resulting solution was determined.

Experimental Procedure:

The titration was carried out, and the volume of NaOH used in each trial was recorded.

The results are presented in the table below:

The pH values for chloroethanoic acid, dichloroethanoic acid, and ethanoic acid are given. The calculation of Ka is based on the assumption that the concentration of the acid (HX) tends to the total acid concentration, and the concentration of the conjugate base (X^-) is mostly contributed by the salt resulting from neutralization.

Using these assumptions, we have:

[H⁺] = Ka

If the concentrations of the acid and salt in the mixture are equal:

[H⁺] = Ka = 1

[H⁺] = Ka

-log Ka = -log [H⁺]

pKa = pH

The aim of neutralization is to neutralize all the acid and make the concentration of the salt the same as the concentration of the 20.0cm³ acid. Therefore, the titre need not be exactly 20.0cm³.

If the titre were, say, 22.0cm³, a further 20.0cm³ of acid is still added. The mole of the salt from neutralization is the same as the mole of the 20.0cm³ acid. As an equal mole of salt and acid is needed, a further 20.0cm³ of acid is still added.

The Ka values calculated for ethanoic acid, chloroethanoic acid, and dichloroethanoic acid are as follows:

• Ethanoic Acid (CH₃COOH): Ka = 2.514 × 10^-5 mol dm^-3
• Chloroethanoic Acid: Ka = 0.001 mol dm^-3
• Dichloroethanoic Acid: Ka = 6.310 × 10^-3 mol dm^-3

It is important to note that the Ka value for chloroethanoic acid obtained in this experiment shows the greatest divergence from the accepted values. This divergence can be attributed to the assumptions made in the calculation, as the slightly ionized acid contributes to the discrepancy.

One possible explanation for the relative values of Ka for the three acids is that the more chlorine atoms an acid has, the greater the dissociation percentage for that acid. This results in a higher concentration of the conjugate base (X^-) and a lower concentration of the acid (HX), leading to a higher Ka value.

This experimental method can also be used to determine the dissociation constant (Kb) of weak bases. For example, for ammonia (NH₃), the following equation can be used:

Kb = [OH^-] / [NH₄⁺]

The procedure can be repeated with NH₃ and HCl, and by using the known pH, the Kb value can be calculated.

Discussion:

1. Distilled water can be added to a buffer solution, but it should not be added to the conical flask during the experiment, as it can alter the pH of the solution.
2. The pH meter should be calibrated by washing it with deionized water and then with the acid to ensure accurate pH measurements.

Conclusion:

In this experiment, we successfully determined the dissociation constants (Ka) of three weak acids: ethanoic acid, chloroethanoic acid, and dichloroethanoic acid. The Ka values obtained were as follows:

• Ethanoic Acid (CH₃COOH): Ka = 2.514 × 10^-5 mol dm^-3
• Chloroethanoic Acid: Ka = 0.001 mol dm^-3
• Dichloroethanoic Acid: Ka = 6.310 × 10^-3 mol dm^-3

These values showed variations from the accepted literature values, with the Ka of chloroethanoic acid exhibiting the greatest divergence. This discrepancy can be attributed to the assumptions made during the calculations and the influence of the partial ionization of the acids.

Recommendations:

Further experiments should be conducted to refine the determination of Ka values for weak acids, taking into account the partial ionization and other factors that may affect the accuracy of the results. Additionally, the method demonstrated in this experiment can be extended to determine the dissociation constants (Kb) of weak bases, offering a broader understanding of acid-base chemistry.