Frianeza
Chemistry Stephen 1AL
Lab 100
9/24/12
Acids in the Environment II: Indicator Titration
Abstract
The task in this experiment was to accurately determine the molarity of an HCl solution
of unknown concentration. The solution was diluted and then a titration was performed
using the base Tris(hydroxymethyl)aminomethane (TRIS) and bromocresol green
indicator. TRIS in solution with the indicator turns blue, and the dilute HCl was added
until it turned green. After three trials with less than 4% difference were performed, the
average molarity of the dilute HCl was calculated to be 0.1123 M.
Introduction
The chemical reaction taking place in this experiment is the following:
HCl(aq) + TRIS + TRISWhere TRIS = (HOCH2)3CNH2
The goal of the experiment is to accurately determine the molarity of the unknown HCl
solution to 4 significant figures. The measuring instrument with the least number of
significant figures is the buret with 4, so the calculations will be accurate to at most 4
figures. By figuring out the molarity of the solution, the potential danger of it will be
known and it can be used to obtain accurate data in other experiments. HCl will be
titrated to the basic TRIS solution, and the color shown by the indicator will be carefully
monitored for changes. The method of titration was chosen due to the accuracy of
measurements; burets are able to measure to 0.01 mL. The theory behind this experiment
is that, as an acidic solution is added to a basic solution, the hydronium and hydroxide
ions react to form water. This lowers the pH of the basic solution in the flask. At a pH of
roughly 5, the equivalence point, bromocresol green turns green, and the number of
moles of HCl and TRIS are equal. At this point, the volume of HCl used was measured
and, knowing the amount of TRIS used, the number of moles of reactants was calculated,
leading to the determination of the molarity of the unknown HCl solution.
Methods
5 mL of the concentrated HCl was pipetted into a 250.00 mL volumetric flask, and water
was added until the line. The buret was rinsed twice with the dilute HCl. Between 0.15
and 0.25g of the TRIS was measured out and then dissolved into a 250 mL Erlenmeyer
flask so that the final solution volume was 50 mL. 5 drops of bromocresol green were
added. Titration of the HCl began at a fast rate, swirling as the acid was added, and
slowed down as the equivalence point was approached. Upon reaching the equivalence
point, when the solution turned green, titration was stopped and final volume of HCl was
recorded. The following measurements were taken: mass of TRIS to the nearest 0.0001
grams, and initial and final volume of HCl to the nearest mL. 0.01 The procedure is from
the General Chemistry 1A Laboratory Manual, Michelle Douskey, Fall 2012.
Results
Trial TRIS (g)
1
0.2420
2
0.1776
3
0.1533
Initial HCL (mL)
0.60
0.88
0.82
Final HCl (mL)
18.26
14.01
12.10
Used HCl (mL)
17.66
13.13
11.28
Molarity (M)
0.1131
0.1117
0.1121
Knowing that the molar mass of TRIS is 121.14 g/mol, the following calculations were
performed to determine molarity of HCl:
For trial 1:
The calculations were repeated for Trials 2 and 3.
Average Molarity: 0.1123 M
All digits from all measurements were recorded and final number of significant figures
was determined after all calculations. In order to determine the concentrated molarity, the
following calculation was performed:
Class average for the molarity: 0.0992 M
Percent difference from class average: 13.21%
Discussion and Conclusions
With unknown concentration of an acid or base, titration is used because the method
employs very accurate measurements and it can be performed slow enough that the
process can be stopped at any point. One very important observation from this
experiment was that just one small drop was able to change the color of the solution from
blue to yellow instantly. In all trials performed, the proper green color was not obtained,
with each trial turning yellow. Inexperience in executing accurate titrations and being
unable to successfully perform partial drops played a large factor in not obtaining the
green equivalence point solution color, and the additional HCl added which changed the
solution yellow was likely the reason for the high molarity calculated and percent
difference from the class average being so great. However, the molarities calculated in all
three trials were within 1% of the average of 0.1123 M, so the titration was consistent
each time. Overshooting the equivalence point led to believing the molarity of the HCl
was higher than it should have been, in this case 0.0131 M higher. Therefore there were
fewer moles per liter than the calculations say, so a greater volume of HCl would have
been required to reach equivalence. Using the data from trial 1 and the class average for
molarity, the amount of volume that should have been used was 20.01 mL, compared to
the experimental value of 17.66 mL.
Solving for V yields:
However, this calculation does not agree with the intuition that in overshooting the
equivalence point, a lower volume must have been required to reach equivalence. Thus,
there must be other factors involved that led to the calculated average molarity being
greater than the class average. Further experience in operating and reading the laboratory
equipment should lead to more accurate experiments.