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Lina Marchi

Megan Hill

Formal Laboratory Report

C125 Experimental Chemistry I

Section 20509


November 9, 2018

Analysis of Acetic Acid in Vinegar Using Titration


The experiment Vinegar Analysis is designed to investigate the concentration of acetic acid in the vinegar used in the experimental titrations. The use of titration is critical to be able to determine the equivalence point of an acid-base reaction. The major problem to solve in this experiment is to determine the equivalence point in the reaction of acetic acid and sodium hydroxide.

Acids are defined as “substances that produce H ions in an aqueous solution”.  Conversely, bases are defined as “substances that produce OH ions in an aqueous solution”.  Acid-base reactions involving titration are used to solve for the equivalence point of the reaction. The equivalence point of this specific reaction is the point at which the base (sodium hydroxide) neutralizes the acid (vinegar/acetic acid) . This is nearly impossible to obtain; thus, the end point is actually found during titration using a chemical pH indicator. The use of the chemical indicator phenolphthalein in this lab is indicative of neutralization of these two substances as the end point in titration is “characterized by color change of the solution”.  The equation shown below demonstrates the process of titration:

CH3 COOH(aq) + NaOH(aq)¬¬ → NaCH3COO(aq) + H2O(l) (1)

The purpose of this lab is to determine the end point of titration in sodium hydroxide and acetic acid as well as determine the concentration of acetic acid in the vinegar solution used. This reaction produces a combination of a salt and water. In the titration, small amounts of the base are dropped into the acidic solution. The combination of the acid and base will produce a pink color with each drop added due to the chemical indicator used. The end point is reached when the pink color persists for about thirty seconds.  However, if more base is added to the acid than what is needed to reach the end point then the solution will turn completely purple. This means that the titration was not as accurate as if it had been pale pink.


100 mL of the standard .1 M NaOH was measured into a 250 mL beaker. The buret was rinsed with 2 times with 5 mL portions of standard .1 M NaOH .10 mL of vinegar was measured into a 10 mL graduated cylinder and poured into a 125 mL Erlenmeyer flask where 2 drops of phenolphthalein indicator were also added to the flask. NaOH was added in 1 mL portions to the flask to titrate. The flask was swirled with each addition and a pink color was observed until it lasted for 30 seconds. The volume of NaOH added was too much to find the proper end point of the titration. The process was repeated for 3 additional NaOH and vinegar titrations to find the nearest end point of the neutralization.


The data in the following tables provide the results of the experiment described in the previous section. Table 1 shows the data of the standardized NaOH solution used in three titrations. The initial volume of NaOH was measured using the buret before titration and the final volume of NaOH was measured using the buret reading after the titration. The difference in final volume of NaOH and the initial volume of NaOH produced the total volume of NaOH used in each titration (2). The moles of NaOH used in titration was then calculated using the equation for molarity (3).

Final Volume NaOH-Initial Volume NaOH=Total Volume NaOH used in Titration (2)

Molarity=Moles of Solute/Moles of Solution (3)

Table 1: Amount of NaOH Solution Used in Titration

Trial # Initial Volume of NaOH solution (mL) Final Volume of NaOH solution (mL) Volume of NaOH used in titration (mL) Moles of NaOH used in titration (mol)

Trial 1 1.20 28.4 27.2 .008

Trial 2 0.40 27.8 27.4 .008

Trial 3 .30 27.7 27.4 .008

Data Table 2 expresses the data of vinegar and acetic acid that was used in each titration. The moles of acetic acid was determined through the use of stoichiometry. The mass of NaOH for each titration was converted to moles of acetic acid to find the amount of acetic acid in vinegar. The mass percent of acetic acid in vinegar was calculated by determining the mass of acetic acid in grams from the previous mole calculation and then using the mass percent equation (4).

Mass % of Solute= Mass of Solute (g)/Volume of Solution (L)   (4)

Table 2: Determination of Mass% of Acetic Acid in Vinegar

Trial # Moles of Acetic Acid in Vinegar (mol) Mass of Acetic Acid in Vinegar (g) Mass % of Acetic Acid in Vinegar (%)

Trial 1 .008 .519 5.85

Trial 2 .008 .519 5.85

Trial 3 .008 .519 5.85

Results and discussion

The average volume of NaOH solution used in titration was 27.3 ± .1 mL. The average percent of acetic acid in vinegar was 5.85± .1%. The same volume of vinegar was used in each trial, but the variable changed was the volume of NaOH used to titrate in each trial. The equilibrium point of titration is nearly impossible to find, but the general end point of titration where the acid is neutralized by the base could be relatively obtained. Each of the three trials were all visibly over-titrated; however, the averages were determined using the data with this knowledge in mind. The second and third trials were more consistent with one another than the first trial. The initial trial showed a volume of 27.2 mL to be titrated with the vinegar whereas the second and third trials used 27.4 mL of NaOH to titrate.

Although each trial was over-titrated, this was a source of consistent error in the experiment. The moles of acetic acid and concentration proved to be the same for all the trials despite the over-titration errors. The experimental process can be improved using an even more precise buret to obtain even smaller volumes of sodium hydroxide for the titration. The smaller portions of solution used, the more accurate the experiment will become. The instant at which the end point is found would use less NaOH solution which could account for the over-titration errors found in this experiment.

The largest source of error in finding the end point and the percentage of acetic acid in vinegar most likely resulted from over titrating each sample. As the NaOH solution was added, it became increasingly difficult to obtain small drops to the vinegar. The solution was initially clear and with the addition of NaOH a bright pink color was produced. The more drops of NaOH added to the vinegar solution, the darker the solution became. Adding a larger portion of NaOH near the end of titration caused each trial to become over-titrated.

Another source of error could have been the starting volume measurement of NaOH. The meniscus on a graduated pipet can be difficult to obtain a clear reading from. The initial volume of NaOH solution in the graduated pipet should have been more consistent than the actual readings recorded. Obtaining the final volume left in the graduated pipet was also difficult to find as the meniscus is complicated to gauge in the reading.

The final probable source of error could have been the drops of NaOH adhering to the sides of the flask. Some drops from the stopcock were sliding down the sides of the flask and were stuck to the sides. This most likely contributed to the high volume of NaOH used in the titration. The over-titration also could have resulted from this problem as well. The volume used of NaOH could have been slightly more as drops that stuck to the sides of the flask slipped down into the vinegar solution.


The Vinegar Analysis lab is used to determine the concentration of acetic acid used in the titration of sodium hydroxide and vinegar. The average volume of NaOH solution was 27.3 ± .1 mL and the average percent of acetic acid in vinegar was 5.85± .1%. The results stem from the over-titrated trials and the standard deviation is reasonably low for the data collected. The results can be deemed precise based on the measurements obtained; however, the data is not completely accurate due to the errors in titration. The accuracy could be improved by running more careful titrations to improve the average value of NaOH volume used to titrate. The results are still considered significant as the values can be used to improve future experiments in titration. Improving the accuracy of the experiment is most important and can be done through more experimentation.


1. Anliker, Keith, Hongqiu Zhao. Laboratory Manual for Experimental Chemistry I, 2016. 17-22.

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