Formal Lab Report of Vinegar Lab PDF

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There is one formal lab report required for the Chem-C125 lab and it is this one....

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Formal Laboratory Report C125 Experimental Chemistry I Section 20500 IUPUI November 09, 2018

Finding Unknown Concentration using Titration INTRODUCTION Most people have faced some kind of stomach problem in their lives. One of the most common stomach problem is having too much stomach acid. This problem can be easily dealt with by taking antacids (base) such as Tums, which can reduce the amount of stomach acid. According to belmarrahealth.com, stomach acid helps process food and kills harmful bacteria, however, too much of stomach acid can lead to severe disorders if it is not treated. Acids and bases can be found in common every day foods, in-taking large quantities of either can cause many problems. So, how does one know how much of each substance is needed to neutralize each other? How much concentration can be neutralized with the given amount of acid or base? They use a process called acid-base titration. Acids are molecules that produce H+ ions in aqueous solution, bases on the other hand produce OH- ions. In an acid-base reaction, the H+ ions and the OH- ions mix together and neutralize the reaction, the reaction usually produces water and a salt. For example= NaOH (base) + C2H4O2 (acid)  NaC2H3O2 (salt) + H20 (water) (1) The purpose of an acid-base titration is to find the concentration of an unknown acid or base using a known value. In the given example above (1), theoretically, enough amount of NaOH will neutralize the C2H4O2; this is called the equivalent point of the titration. However, due to the microscopic measurements that people can’t do yet, it is impossible to get an actual equivalent point; this is where pH indicators come in. pH indicators can be used to find the end point of the titration, the end point can be noticed by a color change of the solution. In the example above (1), if the color of the solution changes, then it means that a small amount of NaOH is excess in the solution. The concentration of the unknown can be found using molarity. Molarity is the number of

moles of solute per 1 liter of solution (mol/L). A more preferred way in today’s world is percent concentration. The percent concentration can be found by volume percent concentration or mass percent concentration. Vinegar, which was used in this experiment, was expressed in mass concentration, mass of solute per mass of solution multiplied by 100%. The purpose of this lab was to determine the concentration of acetic acid in vinegar. In this experiment, known concentration of NaOH was used to titrate unknown concentration of acetic acid in vinegar. The chemical formula of the reaction was the following = NaOH (base) + C2H4O2 (acid)  NaC2H3O2 (salt) + H20 (water) The pH indicator used in this lab was phenolphthalein which turned pink at the end point. The mass percent of acetic acid in vinegar was found by taking the mass of acetic acid and diving it by the mass of the whole sample (g).

EXPERIMENTAL A buret was filled with 0.3005M +/- 0.0006mol/L NaOH up to 50mL. A crude titration was first performed using a 250-mL Erlenmeyer flask with 10mL vinegar and 3 drops of phenolphthalein indicator. The volume of NaOH used for this crude titration was used as a guide for the actual titrations. After the crude titration, a 250-mL Erlenmeyer flask was weighed. Following that, approximately 10.00mL of vinegar was added in the flask using a 10mL graduated cylinder. After that, approximately 10mL of deionized water was added to the flask along with 3 drops of phenolphthalein indicator (Note: Adding too much indicator could have decreased the accuracy). NaOH from the buret was titrated into the Erlenmeyer flask until a light-pink solution was formed and stayed for 30 seconds. The procedures after the crude titration were performed two additional times, a total of three times using approximately the same values and 3 drops of phenolphthalein indicator. The buret containing NaOH was filled when needed.

DATA AND CALCULATIONS

The following results were obtained from the procedures above. The Molarity of NaOH was written on the container which had NaOH, no calculations were done to find the Molarity. The mass of acetic acid in vinegar was found by taking the molarity of NaOH and converting it to mass of acetic acid. First, the Molarity of NaOH was converted to moles of NaOH using molarity= moles of NaOH/L of NaOH. After that, using 1:1 ratio, moles of NaOH were converted to moles of acetic acid (C2H4O2). Following, the moles of acetic acid were converted to grams by simply multiplying by the molar mass of acetic acid. The mass% of acetic acid in vinegar was found by diving the grams of acetic acid in vinegar by the mass of vinegar and multiplying by 100%. Molarity of

Trial 1 0.3005+/-

Trial 2 0.3005+/-

Trial 3 0.3005+/-

NaOH(mol/L) Mass of flask (g) Mass of Vinegar Volume of NaOH

0.0006mol/L 112.94g 9.89g 27.89mL

0.0006mol/L 117.46g 9.79g 23.05mL

0.0006mol/L 124.31g 9.79g 27.44mL

used (mL) Moles of NaOH

1.49mol

1.23mol

1.46mol

used (mol) Mass of acetic acid

0.503g

0.416g

0.495g

in vinegar (g) Mass % of acetic

5.08%

4.25%

5.06%

acid in vinegar (%)

Average mass % of acetic acid in vinegar= 4.80%. Standard deviation of mass % of acetic acid in vinegar= 0.224%. Manufacturer listed mass % of acetic acid in vinegar= 5%. % error of the mass % of acetic acid in vinegar= 4%.

RESULTS AND DISCUSSION As given, the molarity of NaOH used for this titration was 0.3005M +/0.0006mol/L. The Average mass % of acetic acid in vinegar was 4.80% with a standard

deviation of 0.224%. There was a 4% percent error in the experiment. All three trial were considered for finding the average mass % of acetic acid. However, trial 2 reached the end point using less volume of NaOH. There was approximately a difference of 4mL in trial 2 when compared to trials 1 and 3. A possible explanation for this could be error in measurement; the user may have misread the buret and noted the wrong volume of NaOH. Another explanation could be that the Erlenmeyer flask was not properly cleaned with deionized water. This may have caused some particles to stay in the flask from previous experiments or trials, which could have caused the solution to reach the end point earlier than expected. An error that was present in all three trials was over-titration. Even though extreme precautions were considered when the vinegar was almost titrated, a single drop of NaOH changed the solution from colorless to dark pink. Adding in extra amounts of NaOH could result in using up more NaOH then what was actually need for the end point. Connecting this to mass acetic acid in vinegar (g), the mass of acetic acid calculated would be higher than what was actually used at the end point. An error of 4% wasn’t necessarily bad, however, it could have been improved. A couple of procedural errors could have been eliminated, resulting in a better yield. First, each Erlenmeyer flask should have been washed with deionized water thoroughly to clean off any left-over substances from previous experiments; this would have resulted in a precise result. An improvement would have been to perform an additional crude titration to use a better volume as a guide. If the volume of NaOH used for both crude titrations were precise, those values could have been better to use as a guide for the actual titrations. This would have resulted in a decreased amount of over-titrations. Lastly, A few drops of NaOH were stuck on the sides of the Erlenmeyer flask due to the constant swirling of the flask. A drop was also left hanging outside the buret after the titration was completed. Simple procedural errors like this could be eliminated to provide a better yield and results in the future. Though there were a few errors in this experiment, a percent error of 4% was acceptable considering the over-titration and standard deviation. The main source of error could have been from trial 2 as discussed earlier, it could have been eliminated if

additional trials were performed; however, since only three trials were performed, it was best to consider the data from trial 2 as well.

CONCLUSION The molarity of the NaOH solution used for this titration was 0.3005M +/0.0006mol/L. The average mass % of acetic acid in vinegar was 4.80 +/- .2%. The percent error, when compared to the manufacturer’s listed mass of 5% acetic acid in vinegar, was calculated to be 4%. The standard deviation was a bit high due to errors in trial 2, but the average still resulted in a number that was close to the theoretical percentage (5%). For this experiment to be more accurate and precise in the future, additional trials could be performed to reduce errors or even discard some of the trials. This experiment was significant because finding the concentration of an unknown substance could be helpful in finding the concentration of stomach acid or even discarding waste acid for neutralization.

REFERENCES Pr e La bPo we rPo i n tVi n e g a rAna l y s i s , Ca n v a s , 2 0 16 . Ex p e r i me n t a l Ch e mi s t r yI , La bMa n u a l , Zh a o , 20 1 6 . Libretexts. “Titration.” Chemistry LibreTexts, National Science Foundation, 25 July 2016, chem.libretexts.org/Demos,_Techniques,_and_Experiments/General_Lab_Techniques/Tit ration....