Erica Jenson Shelley Koenig Chemistry Acid Base Titration Lab Report
! Prelab Questions 1) Write the equation for the neutralization reaction between ammonia and HCl. NH4OH(aq) + HCl(aq) !! H20(L) + NH4Cl(aq) 2) Write the equation for the neutralization reaction between acetic acid and sodium hydroxide. CH3COOH(aq) + NaOH(aq) !! H20(L) + CH3COONa(aq) 3) Why should a minimal amount of water be used to cover the red cabbage leaves in part 1? A minimal amount of water should be used so that the strength of the indicator is greater. 4) What is the function of test tube A in part 1? Test tube A functions as the control 5) When performing titrations, why should the Erlenmeyer flask be constantly swirled? The Erlenmeyer flask should be constantly swirled to mix the solution and to increase the rate of reaction. It also allows for a more correct experiment, since when the entire solution becomes titrated, it will be known since it is properly mixed entirely. 6) Why were the titrations in part 2 repeated?
The titrations in part 2 were repeated so that experimental error would become more obvious. It would be hoped that both results would be the same, but if they weren’t—a recognizable error would be found.
! Introduction In this lab, the process of titration was used to determine the concentration of the acetic acid solution of vinegar. The molarity and percent of the vinegar solution could then be calculated. This lab is completed by having vinegar mix with sodium hydroxide. The equivalence point between the two solution was seen by a change in the color from the indicator. The two solutions in this lab are vinegar (CH3COOH) and sodium hydroxide (NaOH) The vinegar is an acid and the sodium hydroxide is a base. The combination of an acid and a base is a neutralization reaction. In a neutralization reaction, the reactants are a base and an acid and the products created are a salt and water. The equation for this reaction is CH3COOH(aq) + NaOH(aq) !! H20(L) + CH3COONa(aq). Titration is the process in which a known concentration of acid or base can be added to another solution with a possible unknown concentration to create a neutral solution. A weak acid or base is used as an indicator of a pH change. In this lab, phenolphthalein is used as the indicator. The pH changes greatly when the equivalence point is reached. Therefore, the solution with phenolphthalein in it will change color as the solution is neutralized. The vinegar was the acid so from the testing of the affect of phenolthalein in part A , it was recognized that until the solution became slightly basic,
the color would remain clear. 1 drop before the color kept the pink color was when the reaction was neutralized and had equivalence. At that point, the solution was neither a base or an acid but instead neutral with a pH of 7. The goal from these results was to find the concentration of the vinegar. From knowing the concentration of the one solution, along with the data collected, the concentration of the other solution (the vinegar) can be found with stoichiometry. The titration process relates to chemistry class as it uses knowledge of concentration, pH, acids, bases, equilibrium, and neutralization. The world outside the chemistry class uses the titration technic in many different applications. One application is Blood Sugar Testing. A small machine called the blood glucose meter measures the amount of glucose in a diabetic’s blood. A small sample of blood is put on a test strip, mixed with reactants and given a slight electrical current. The affect the blood has tells the concentration of the reactants and is then used to measure the amount of glucose in the blood.
! Data Table
! Part 1: Affect of Phenolphalein Indicator on Reactants Part 1
Reagents Added
Observations
Tube A
None
None
Tube B
NaOH
Cloud of bright pink
Tube C
HCl
Clear solution
Tube D
Vinegar
White cloud at top
Tube E
Ammonia
Pink cloud at top
! Part 2: Vinegar and NaOH Titration Part 2
Trial 1
Volume of Vinegar
.0005 Liters
Molarity of NaOH
.5 M
Volume of HCl (Initial reading)
6mL
Volume of HCl (Final reading)
59mL
! ! Calculations 1. For the titration of vinegar, calculate the volume of NaOH needed to reach the equivalence point. 59mL – 6mL = 53mL = 0.053L 2. Calculate the number of moles of NaOH needed to neutralize the vinegar. 0.053L NaOH
!
x
.5mol NaOH 1L
=
0.0265mol NaOH
3. How many moles of acetic acid are contained in the sample of vinegar? .0265mol NaOH
!
=
.0265mol Acetic Acid
4. Calculate the mass of the acetic acid in the vinegar. .0265mol Acetic Acid x
60g Acetic Acid = 1.59g Acetic Acid 1mol Acetic Acid
!
! 5. Calculate the mass of the vinegar sample. .005L Vinegar
x
!
1g H2O .001L
=
5g H2O
5g H20 + 1.59g Acetic Acid = 6.59 g Vinegar
!
!
6. Calculate the percent acetic acid in vinegar by dividing the mass of the solute by the mass of the solution and multiplying by 100.
! 1.59g Acetic Acid 5g H2O
x
100
=
31.8% acetic acid in vinegar
! ! Analysis and Conclusion Questions 1) Would you use an acid or a base as a standard when titrating against a solution of soda pop? Why? A base would be used since soda pop is acidic. In order to neutralize the solution, a base would be needed to combine with the acid. 3) When selecting an appropriate chemical indicator, one should choose an indicator that will undergo a shift to a darker color. Why?
The darker the color, the easier it is to tell when a change occurs at the exact moment that it does.
! Synthesis Questions 1) When solutions approach the titration equivalence point, they demonstrate an increased sensitivity to hydronium ion equilibrium. Why? As the solution approaches the titration equivalence point, little changes in the amount of hydronium ion equilibrium will shift the solution’s pH and shift the equilibrium easier.
! Discussion
! In this titration lab, the acetic acid concentration was found, along with the percent of acetic acid in the vinegar. To do that, equilibrium, where the acid and base are neutralized, had to be created. In Part 1, The indicator of phenolthalein was looked at in different substances: two bases, two acids, and one neutral solution. The results showed the bases (NaOH and ammonia) turning pink. The acid solutions (HCl and vinegar), as well as the generally neutral solution of phenolthalein produced a somewhat clear color. By continually adding NaOH to the vinegar solution until the solution turned pink for at least 30 seconds, one could get a good idea of where the point of equilibrium was. That was the way it was known that the pH had shifted from acidic to slightly basic.
This process made it somewhat difficult to find the exact point of neutralization. The color change was so slight that it wasn’t obvious when the solution became neutral. Even when there was pink color, it often lasted only for a few short seconds until it was swirled away. The actual point of neutralization was a drop before the solution turned pink, but the results from one drop after was close enough. The volume of NaOH that was put into the vinegar was 53mL. The amount of NaOH one drop before that would have been ideal but that did not have a large affect on the data. For this lab, the process took a long time, which is why two trials did not get to be completed. The completion of only one trial is unfortunate because the use of two trials (to make experimental error more obvious) could not be put into affect. If mistakes were made in the single trial, the mistakes were not as easily picked up on. I believe the results were generally accurate, although experimental error could have been a factor by inaccurate measurement readings, as well as inaccurate judging of when the pink color was what it was suppose to be.The volume of NaOH was 53mL at a concentration of 0.5M. It seemed accurate that the mass of the acetic acid in the vinegar was 1.59g, while the mass of the NaOH was 6.59g. The comparison between those two numbers is understandable because there was much more NaOH then vinegar. The final calculation to find the percent acetic acid in vinegar made it ore obvious that a large mistake was not made. The percent found was 31.8% acetic acid in vinegar, which seems reasonable. The percent was enough that the vinegar solution was acidic enough, but not so much that it was unsafe. These results conclude to the belief that no serious mistake in the lab was made. Overall, the purpose was accomplished. The point of neutralization was found and
from that the concentration, molarity, and percent compositions of vinegar could be calculated.
