Acid-base titrations are commonly used when finding the amount of a given basic or acidic solution through an acid base reaction. The analyte is usually the solution with the unknown concentration while the solution used as a titrant has a known molarity (Levie). An indicator must be used in this type of reaction to indicate the end point. The indicator to be used should however have a strong color that changes immediately near its pKa and its pKa should also be equivalent to the pH of the titration’s endpoint to minimize errors (Levie).
The analyte is usually prepared through dissolving a substance to be studied into a solution inside a flask. A suitable indicator is added and a reagent run slowly from a burette to the solution. The end point is indicated by a sharp color change of the solution. Since we are dealing with two solutions, on of known molarity while the other one has molarity unknown, we determine the unknown through calculation using the formula MaVa=MbVb
at the end of the experiment (Levie). The reaction between acid and base is a neutralization reaction and give salt and water as the final products.
Acid + Base→ Salt + Water
Purpose of the experiment
This experiment aims at:
1. Learning how to prepare solutions using proper glassware.
2. Determine the endpoint of an acid-base titration using a chemical indicator.
Procedure
Solution preparation
50.0 mL of a 0.01 M hydrochloric acid solution was prepared from a 4.0 M stock solution in a 50 mL volumetric flask.
100.0 mL of a 0.10 M sodium hydroxide solution was prepared from solid in a 100 mL volumetric flask.
Equipment set up
Titration apparatus were set up as demonstrated by the instructor
Using deionized water, filling the burette with water was practiced until it was comfortable filling and using the burette.
The remaining water was drained form the burette.
The walls of the burette were washed by adding approximately 10 mL of the 0.10 M sodium hydroxide solution.
The burette was filled with an easily readable volume of 0.1 M sodium hydroxide.
The starting volume was recording to the nearest 0.01 mL.
Standardization of NaOH
10 mL of the 0.10 M hydrochloric acid solution and 2 drops of phenolphthalein were added in a 125 mL Erlenmeyer flask. A magnetic stir bar was slide gently down the side of the flask into the solution.
The burette was positioned over the flask using a stir pad on a low speed.
The base was added slowly to the acid until the color turned to light pink.
The volume of the base delivered to the flask was noted.
The processes was repeated with more precision, to achieve a very faint pink color that remained for at least 1 minute.
The procedure was repeated three times, and the volume of base delivered to the nearest 0.01 mL.
The concentration of the sodium hydroxide was calculated using the concentration of the acid as the standard.
All waste was disposed through draining on the sink.
The procedure above was repeated three more times using 10 mL of clear soda.
The concentration of the acid was determined using the determined concentration of sodium hydroxide as the standard.
Results
Solution of hydrochloric acid prepared
1.25 mL
Titration using sodium hydroxide
Trail
1
2
3
Final burette reading of NaOH (mL)
45.4
37.3
46.4
Initial burette reading of NaOH (mL)
36.0
27.4
37.3
Volume of NaOH used (mL)
9.4
9.9
9.1
Color at end point
Bright pink
Bright pink
Pale pink
Average= 9.8 mL
Titration using clear soda
Trial
1
2
3
Final burette reading ()
50.5
28.7
31.0
Initial burette reading ()
46.4
25.2
28.7
Volume of clear soda used ()
4.1
3.5
2.3
Color at end point
Pale pink
Pale pink
Pale pink
Average used= 3.3 mL
Calculation
Preparation of HCl solution
V1= (M2V2)/M1
((0.1M) × (50Ml))/4.0M
=1.25 mL
Concentration of acid
MaVa= Mbvb
Ma× 10= 0.10 × 3.3
Ma= (0.10 × 3.3)/10= 0.033 M
Solution of NaOH
100mL× 1L/1000Ml ×0.1 mol NaOH/1L × 40.0 g NaOH/ 1 Mol NaOH= 0.4g NaOH
Discussion
Reaction of an acid and base result to a neutralization realization. Adding sodium hydroxide or clear soda into the flask containing hydrochloric acid results to formation of salt and water. The end point of the reaction is indicated by the change of color of the solution since it contains phenolphthalein as indicator. The reactants in this reaction are strong base and strong acid and therefore phenolphthalein acts as the suitable indicator since it charges color sharply at the endpoint. The results obtained are almost similar since the bases used in the experiments are both strong. Sources of error in the reaction could arise due to use of contaminated reagents.
Conclusion
The aim of the experiment was achieved since it was possible to determine the molarity of the unknown solution. However, the success the reaction and accuracy of the results is dependent on the appropriateness of the indicator used (Levie).
Works Cited
Levie, Robert De. Aqueous Acid-Base Equilibria and Titrations. New York: Oxford University Press, 1999.
