Determining the Mass % of Acetic Acid in Vinegar PDF

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Mass Percent of Vinegar C125 Lab Report...

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Determining the Mass Percent of Acetic Acid in Vinegar using Acid-Base Titration Introduction Experimentation in the laboratory requires the knowledge of the specific concentrations of the chemicals utilized. Major consumer products like food and household supplies contain acid and bases. In order to determine the exact concentration of an unknown chemical, one can use the process of titration. This process is utilized in determining the amounts of specific substances in commercial products. This allows companies to understand the makeup of their products for safety purposes as well. The purpose of this specific experiment was to determine the mass percent of acetic acid in vinegar by perform an acid-base titration. Titration is the process of qualitative analysis of an acid by slowly adding a base solution of known concentration into known volume of an acid with an unknown concentration containing an indicator. This process can also be utilized by using a known amount of acid into an unknown amount of base. Neutralization occurs when an acid and base react with each other leaving no excess hydroxide or hydrogen ions present in the solution. Another important part of titration is the use of an indicator. An indicator is used to show when neutralization occurs and the indicator changes the color of the solution to make the end point visible. Many substances used in everyday life contain acids and bases. Acids are substances that are identified by their production of H+, or hydrogen ions in an aqueous solution. On the other hand, bases are substances that are identified by their production of OH-, or hydroxide ions in an aqueous solution. When acids and bases are combined, the hydrogen and hydroxide ions bind to make water and a salt, and the solution is neutralized when there is an equal amount of acid and base in the solution. CH3COOH(aq) + NaOH(aq)  NaCH3COO(aq) + H2O(l)

(1)

Equation 1 represents the chemical reaction equation between the acid and base utilized in this experiment. The acid-base titration was completed by reacting NaOH, or sodium hydroxide, with the acetic acid found in vinegar. NaOH had a known concentration while the concentration of acetic acid was unknown. Determining the concentration of acetic acid was completed by filling a buret with NaOH and slowly adding drops of NaOH to a known volume of acetic acid containing the pH indicator phenolphthalein. The equivalence point of the experiment is exactly when neutralization occurs, and this ideally would be the conclusion of the experiment. However, it is nearly impossible to reach the equivalent point, so the pH indicator

was used to determine the end point of a reaction. A color change in the solution represents an excess of hydroxide ions and the end point of the reaction has been reached. Once the reaction reaches the end point, the amount of NaOH needed to neutralize the acetic acid in vinegar was known and the mass percent of acetic acid was in vinegar was able to be calculated, specific calculations can be utilized to determine the concentration of the acid. Experimental A 50mL buret was initially prepared for titration by rinsing the buret with NaOH 3 times. After the buret was prepared, it was filled until it was between the 0-3mL mark on the buret with a .3599M NaOH solution. After the biuret was prepared with the NaOH, a 250mL Erlenmeyer was weighed with an analytical scale to determine its weight. Then the flask was filled with 10mL of vinegar and weighed again on an analytical scale and the new weight was recorded. After approximating the weight of the vinegar, 3 drops of the pH indicator phenolphthalein were added to the flask. Once the flask was prepared, the crude titration began. The flask was placed directly under the 50mL biuret and the NaOH was allowed to slowly drip one drop at a time into the vinegar solution and this process continued until the vinegar solution turned a light pink color while the flask was being mixed vigorously for more than 30 seconds. This indicated the end point had been reached. After the crude titration had been completed, three different trials of this experiment were performed and the results were recorded. Before each new trial, the amount of NaOH in the buret was recorded and then after the trial the amount was recorded also. Also, the flask was weighed on the analytical scale before and after adding vinegar, as was done in the crude trial. After each trial the mass percent of acetic acid in vinegar was determined. Data and Calculations The results from the vinegar analysis procedure can be seen in the three Data Tables. In Table 1, the moles of NaOH are found from using the molarity and liters of solution. The moles of NaOH are used to find the moles of acetic acid due to it being a 1 to 1 reaction. The molarity of NaOH was calculated prior to the lab, however to calculate molarity one must divide the moles of NaOH by the volume of liters of the solution. These calculations were completed using equations 2 and 3. Moles of NaOH=Moles of Acetic Acid=Molarity x Liters Acetic Acid Molarity of NaOH=Moles of Solute/Liters of Solution Table 1: Standardization of NaOH Solution

(2) (3)

Trial 1 Trial 2 Trial 3

Initial

Final

Volume of

Molarity of

Moles of

Moles of

Volume

Volume of

NaOH

NaOH

NaOH used in

Acetic Acid in

of NaOH

NaOH

Delivered

Solution

Titration (mol)

Titration (mol)

(mL) 1.3mL 23.0mL 1.1mL

(mL) 23.0mL 45.1mL 21.9mL

(mL) 21.7mL 22.1mL 20.8mL

(mol/L) .3599M .3599M .3599M

.0078mol .0078mol .0078mol

.0078mol .0078mol .0078mol

Table 2 represents the data of the percent of acetic acid in vinegar. The mass of vinegar was measured with an analytical scale by putting 10mL of vinegar into a flask and measuring the weight in grams. That mass of vinegar was then used to find the mass % of acetic acid in the vinegar by using the moles of acetic acid in vinegar converted to grams. The moles of acetic acid in vinegar was found by equating the moles of NaOH in the solution since it was a 1 to 1 ratio. The mass % was then found using equation 4. Mass of solute × 100 %=mass % concentration ( Massof solution )

(4)

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

Mass of

Mass of

Volume of

Moles of

Mass of

Mass % of

Flask

Flask +

Vinegar

NaOH

NaOH=Moles of

Acetic Acid

Acetic Acid

(g)

Vinegar

(g)

Delivered

Acetic Acid

in Vinegar

in Vinegar

Trial 1

119.96g

(g) 129.46

9.50g

(mL) 21.7mL

(mol) .0078mol

(g) .4689g

(%) 4.70%

Trial 2

127.51

g 137.03

9.52g

22.1mL

.0078mol

.4774g

4.78%

Trial 3

g 95.20g

g 104.68

9.48g

20.8mL

.0078mol

.6051g

6.00%

Table 3 includes the average mass % of acetic acid in vinegar and the mean and standard deviation found for the reaction. The average mass % was calculated by adding together the individual mass % of acetic acid in vinegar from each separate trial and then dividing by the 3 trials as seen in equation 5. The standard deviation was found by utilizing equation 6 with the results from the 3 trials. The percent error was found by taking the average mass % of acetic acid in vinegar found from the experiment and subtracting the manufacturer listed mass % of acetic acid in vinegar from that and then dividing by the manufactured value. Equation 7 shows the formula for this calculation.

Average Mass % of Acetic Acid∈Vinegar=

Mass % of Acetic Acid ∈Vinegar Total Number of Trials

X ¿ (¿ i− x´ ¿)2 n−1

(5)

(6)

n

∑¿ i=1

Standard Deviation=¿ % Error=

( Approximate − Exact ) Exact

(7)

Table 3: Statistical Analysis of Results Average % Mass of Acetic Acid in Vinegar Standard Deviation of Mass % of Acetic Acid

5.16% 5.2  0.7

in Vinegar Manufacture listed Mass % of Acetic Acid in

5%

Vinegar % Error of the Mass % of Acetic Acid in

3.2%

Vinegar Results and Discussion The average mass and standard deviation of acetic acid in vinegar was 5.2  0.7 respectively. To determine the average mass % of acetic acid, the mass % of acetic acid in each trial were added together and then divided by the number of trials. The mass % of acetic acid in vinegar from trial 3 was higher than the other two trials included in the experiment with it being 6.00%. Trial 3 also had a higher mass of acetic acid in the vinegar in comparison to the other two trials. Trials 1 and 2 appeared to be consistent and trial 3 had a higher mass %. The trial with the higher mass % also had the lowest volume of NaOH used in the titration, which caused the weight of the vinegar in acetic acid to be higher in trial 3 than in the first two trials. The % error calculated from the experiment was 3.2%. The most probable reason for errors within the experiment was most likely over titration. The drops of NaOH from the biuret are difficult to control and, the pink color from the pH indicator became dark pink instead of the pale pink indicating the endpoint of the reaction. This means that more NaOH was used then actually needed, increasing the moles of NaOH and subsequently the moles of acetic acid increased.

Another possible error would be losing the drops of NaOH in the buret or on the walls of the flask. This would increase the volume of NaOH used in the experiment and would also increase the mass % concentration of acetic acid in vinegar. The increase in the volume of NaOH used would increase the number of moles calculated to have been used in the titration. The increase in moles of NaOH would then affect the overall mass of acetic acid in vinegar. This is because the moles of NaOH is a 1 to 1 ratio with the moles of acetic acid, and the moles of acetic acid were used to find the mass of acetic acid in the vinegar. This would also increase the mass % concentration of acetic acid in vinegar. In order to reduce potential errors in future experiments, one could pay more specific attention to the buret specifically and avoid splashing the NaOH on the walls of the flask or in the buret. One could also make note to avoid touching the flasks with their hands as the oils and deposits from the fingers can affect the weight of the flask and alter the results of the mass of vinegar. Experimental errors could have also been made when adding the NaOH to the vinegar and recording the amount of NaOH used. Due to how quickly the solution changed color, it was very easy to over titrate and include much more NaOH than necessary to reach neutralization. If one of the trials was over-titrated, the results could imply that there were more moles of NaOH in the solution and subsequently more moles of acetic acid in the solution which would increase the mass % of acetic acid in vinegar. Overall, there are several different ways that one could easily alter the results of an acid-base titration and it is essential to practice good lab techniques and pay special attention to any errors that could have been made so they can be recorded and accounted for. Conclusion The average mass % of acetic acid in vinegar from three different trials was determined to be 5.16%. The standard deviation determined for this experiment was determined to be 0.7. This standard deviation was fairly low due to the low 3.2% percent error determined for the mass % of acetic acid in vinegar, which speaks to the accuracy of the performed experiment in the three trials of the experiment. The manufacturer’s listed mass % of acetic acid in vinegar is 5% and the mass % of acetic acid in vinegar determined in the experiment was 5.16%. The definition of percent error addresses the closeness in percentage of the manufacturer listed mass % of acetic acid in vinegar to the mass % determined in lab implies that the experiment showed both fairly precise and accurate results. This experiment is highly essential in finding the amount of a

particular substance with a precise and accurate method. This process can be utilized for many different experiments to find important information about a solution. The ability to know the concentrations of specific substances is essential in providing the information and safety measures required to produce products that are used by countless individuals across the world on a regular basis....