Lab write up Experiment 7 november 18

# Lab write up Experiment 7 november 18 - Since we calculated...

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In this experiment, students used acid-base titration to determine the percent acetic acid in vinegar. This was used to determine the neutralization capacity of antacid tablets, and to learn the technique of standardizing a solution. In this experiment, we determined the % acetic acid in vinegar by taking the calculated concentration of CH3COOH in vinegar and dividing it by the 17.4M glacial CH3COOH. Then this number is multiplied by 100 to reach the percent acetic acid in vinegar. The average # acetic acid in vinegar between the two trials was 3.015%. Balanced net ionic equation: In this experiment, in order to determine the mass of CaCO3 in the antacid tablet, first we needed to calculate HCl millimoles; since we know that the mole ratio of CaCO3 to HCl is 1:2. In order to calculate the HCl reacted with base in antacid, take the total volume of NaOH titrated- millimoles of NaOH titrated.
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Unformatted text preview: Since we calculated the HCl millimoles, we just divided it by two to find the millimoles of CaCO3. Then we multiplied the millimoles by the mass of CaCO3. The two calculated masses of CaCO3 were .523919 and .563006; the average of these two is . 5434625. Compared to the 500 mg that is supposed to be in each tablet, there is more CaCO3 in each tablet than the supposed amount. Balanced equation: Some sources of error could have resulted from overheating the flask when dissolving the antacid tablet in HCl. Another source of error could have resulted from the NaOH reacting with CO2 in the air producing H2CO3. H2CO3 dissociates to produce more H+ concentration in solution. The first error could have created an incorrect concentration of CaCO3 because some of the HCl could have evaporated out of the flask. The second error could have created the same issue....
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