This experiment aimed to determine the molarity of acetic acid in vinegar. The process involved titrating the acetic acid with a strong base, sodium hydroxide, to identify the equivalence point of this chemical reaction. Phenolphthalein was used as an indicator because it changes color on the basic side of the pH scale. To ensure accuracy, two conductors conducted three trials to reach the best endpoint possible.
After punching in numbers and calculating results through a stoichiometry equation, we obtained a resultant molarity of 0.960M, but the actual molarity was only 0.845M. We discovered that it took 9.60 mL of NaOH to neutralize 10.00 mL of acetic acid, resulting in a percent error of approximately 13.6%. There were several areas where human error could have impacted the experiment’s findings; however, if burettes are used correctly, these issues should become less noticeable.
Introduction and Background Information:
Acetic acid is the second most abundant component found within vinegar.
This experiment explores the presence of acids and bases in common everyday forms. The individuals conducting the experiment will determine the concentration of acetic acid by taking a sample of vinegar and titrating it with a strong base, sodium hydroxide. The chemical equation for this reaction is shown below:
HC2H3O2 (aq) + NaOH (aq) → NaC2H3O2 (aq) + H2O (l)
The equivalence point of the titration is reached when all H+ ions from acetic acid are neutralized by an equal number of OH- ions from sodium hydroxide.
The equivalence point of weak-acid and strong-base titrations is generally basic. Therefore, an indicator that changes color in the basic region of the pH scale must be chosen. For this experiment, phenolphthalein was used as an indicator because it is colorless in acid and red/magenta in base.
The purpose of this experiment is to determine the concentration of acetic acid in a vinegar sample by titrating the acetic acid (HC2H2O2) with sodium hydroxide (NaOH), which is a strong base.
Materials List
- Pipets
- Vinegar
- Indicator (phenolphthalein)
Experimental Set-up:
- Sodium Hydroxide (1.00M)
- 50 mL burette for NaOH
- One 150 mL Erlenmeyer flask
Procedures:
- Obtain 10.00 mL of vinegar into the Erlenmeyer flask.
- Record the brand of vinegar used.
- Add 2-3 drops of indicator solution (phenolphthalein). Swirl the flask to ensure that the solution is mixed well.
- Record the molarity of base (NaOH).
- Drain enough sodium hydroxide into the waste bucket so that the tip of the burette is filled without air bubbles. Make sure to go slowly or else, there may be a chance that amount of base may go below meniscus.
Record the initial volume reading of NaOH to the nearest 0.01mL. Titrate the first sample with good mixing until reaching the indicator endpoint. The solution should be a pale pink that is neither too light nor too dark. Record the final burette reading to the nearest 0.01mL. When refilling the burette, use only as much hydroxide solution as needed. Finally, clean up all glassware and station used.
Conclusion and Discussion of Results
In this experiment, we attempted to determine the molarity of acetic acid in vinegar. My partner and I used 10.00 mL of vinegar to conduct a titration with sodium hydroxide, a strong base. The acetic acid in the vinegar was titrated with sodium hydroxide to determine the equivalence point (endpoint) of this chemical reaction. We used phenolphthalein as an indicator because it changes color when the solution shifts towards the basic side of the pH scale. When all of the acetic acid had reacted with sodium hydroxide, we reached our endpoint.
After being neutralized with sodium hydroxide, the solution turned pale pink due to the indicator. To ensure the experiment’s accuracy, two conductors attempted three trials to reach the best endpoint possible. In the first trial, around 9.75 mL of NaOH was used, resulting in a deep pink or magenta solution. In the second trial, ironically using the same amount of NaOH as in the first trial led to a dark solution once again.
Lastly, in our final trial, we successfully titrated acetic acid with 9.60 mL of NaOH. After calculating the results using a stoichiometry equation, the resultant molarity was 0.960M, but the actual molarity was 0.845M. The standard deviation of our entire class data was calculated to three significant figures, ending in 0.0430M.
We found that it took 9.60 mL of 1.00 M NaOH to titrate 10.00 mL of vinegar based on our results which had a percent error of around 13.6%. There were many places where human error could have affected the results of this experiment.
There is a possibility that NaOH may have spilled out of the burette, leading to an inaccurate measurement of the amount used. Similarly, an incorrect measurement of vinegar could have occurred if someone accidentally blew out the last drop from the pipette. However, if burettes and pipettes are used correctly, these issues should become less significant and result in a lower percent error.
If the tip of the burette was not filled with sodium hydroxide before recording its initial volume reading, it would affect the resultant molarity of acetic acid.
The volume of NaOH would be greater than the recorded amount, resulting in an increase in the molarity of acetic acid.
If a few drops splashed out of the Erlenmeyer flask during titration, the resultant molarity of acetic acid would be higher than expected.
If the last drop of vinegar solution were blown out of the pipet into the Erlenmeyer flask, it would result in a lower than expected molarity of acetic acid.
If the wet burette was not rinsed with sodium hydroxide before filling, the resultant molarity of acetic acid would be affected. This is because any residual impurities in the burette could react with the acetic acid solution and alter its concentration. Water alone does not affect the molarity of acetic acid since it does not alter its acidity or basicity due to the 1:1 ratio of H+ and OH- ions present.
If a volume of water was added to the Erlenmeyer flask, it would dilute the concentration of acetic acid. This means that there would be less solute per unit volume, resulting in a decrease in molarity. The answer for this question is similar to Number 4.
If methyl orange was used as an indicator, it would not directly affect the resultant molarity of acetic acid. However, if too much indicator is added, it could potentially react with some of the acetic acid molecules and cause a decrease in concentration. Therefore, it is important to use an appropriate amount of indicator for accurate results.
When comparing Phenolphthalein and Methyl on the pH scale, Phenolphthalein is more basic. If a different indicator such as Methyl was used, it would require a different amount of base to change the color of the solution, which would offset the molarity of acetic acid with Phenolphthalein.
The color of the endpoint fades upon standing because eventually all H+ and OH- ions will be neutralized in the solution, changing the pH to where the indicator is. NaOH will eventually ionize and change color afterward.
