Introduction:
In this experiment, we titrated HCl against NaOH in order to find the initial concentration of HCl. We first titrated HCl without the use of a pH meter and then with a pH meter.
Alexander Krzyston Acid Titration| Alexander J Krzyston Acid Titration| Alex James Krzyston
Alex Krzyston Acid Titration| Alex J Krzyston Acid Titration| Alexander James Krzyston
NORTHWESTERN UNIVERSITY Acid Titration| EVANSTON Acid Titration| BURR RIDGE
In the titration without a pH meter, we used a phenolphthalein indicator to determine when the solution has reached the equivalence point, when the solution turned bright pink. For the titration with a pH meter, we recorded the change in volume with every change of 0.3 in pH. We then plotted ΔpH /ΔV vs. V, at the peak of the graph is the equivalence point and used this point to calculate the initial concentration of HCl.
We also titrated Acetic Acid against NaOH to find its concentration and pKa.
OVERVIEW:
1. Sometimes chemists use indicators that change color at a pH of 9 to 10 instead of a pH of 7 because not all acid-base titrations reach their equivalence pints at a pH of 7 when they involve a weak acid or a weak base. For example in the titration of a weak acid with a strong base, like the one we just did of HOAc against NaOH, the equivalence point was at a pH around 9 which is why we used phenolphthalein as an indicator because it changes from colorless to pink over a pH of 8.2 to 10.
2. In the titration of HCl against NaOH we used pH= 7 as the equivalence point because it was the titration of a strong acid against a strong base so when the two were at the equivalence point, all of the OH- from the NaOH neutralized with all of the H+ from the HCl and the concentrations are equal resulting in H2O. All of the HCl went into forming H+ and all of the NaOH went into forming OH- to react with the H+. In contrast, with the Acetic Acid, not all of the HOAc went into forming H+ ions.
This lab called for very careful measurements which was difficult to do and resulted in data that was not as accurate as it could be. Going one drop over the equivalence point had a huge impact on the titration especially since the pH changes so radically near the equivalence point. This made tracking the pH with the pH meter near the equivalence point difficult. These drastic pH changes are clearly visible on our graphs. We had some difficulty in trying to track the change in volume with each change of 0.3 in pH because the pH meter slow in reaction time in adjusting to the pH. Often times we would add NaOH until there was a 0.3 change in pH and when we then stopped to record the change in volume the pH would continue to rise. Also, because of the drastic change in pH near the equivalence point, there were huge fluctuations in the pH meter near this point (after adding only 0.10 the change in pH was 1.39 and after adding 0.02mL of NaOH, the change in pH was 1.22). Initially, there would be a huge change in pH and then as the added NaOH dispersed throughout the solution, the pH would drop. For this reason there are some inconsistencies in our data, especially with the titration of HCl and NaOH with the pH meter. The most noticeable is the change in pH from 3.43 to 6.07 in which the volume changed from 22.60 to 22.85. This had a minimal impact on the shape of the pH vs. volume graph, but did affect the shape of the ΔpH/ΔV vs. volume graph. Instead of a slow steady rise followed by a large, drastic rise and fall, the graph had a slight peak followed by a slight drop followed by the huge rise. I think we should have allowed more time for the pH meter to stabilize before continuing with the titration. However, our data for the titration of NaOH and Acetic Acid was more consistent which shows that as the experiment when on, we became better at working with the pH meter. In addition, because we were working with a weak acid and strong base which created a buffer solution, the pH did not change as rapidly as with the strong acid.
According to the ΔpH/ΔV vs. volume graph our equivalence point was at 32.02mL which is at a pH of 9.10. The equivalence pH was supposed to be at 7, when the added OH- from the NaOH neutralized all of the H+ from the HCl. This difference could result from the fact that we did not wait long enough for the pH meter to stabilize to get the final reading. Had our ability to control the addition of NaOH been better, I do not think we would have had such issues. However, the shape of the graph was correct and shows a steep rise in the region from a pH of about 6.00 to 10.00.
Procedure:
Data:
Titration of HCl against NaOH
without pH meter
Run Initial Buret Reading (mL) Final Buret Reading (mL) ΔV (mL)
1 0.00 23.25 23.25
2 0.00 23.00 23.00
3 0.00 23.00 23.00
Mean 0.00 23.08 23.08
Titration of HCl against NaOH
with pH meter
Vbase (mL) pH of Solution ΔV (mL) ΔpH ΔpH
ΔV (mL)-1
0.00 1.50
17.32 1.87 17.23 0.37 0.02
19.71 2.12 2.39 0.25 0.10
20.60 2.31 0.89 0.29 0.21
21.71 2.63 1.11 0.32 0.29
22.60 3.43 0.89 0.80 0.90
22.85 6.07 0.25 2.73 12.92
22.90 6.49 0.15 0.42 2.80
23.00 7.88 0.10 1.39 12.90
23.02 9.10 0.02 1.22 61.00
23.05 9.49 0.03 0.39 13.00
23.41 10.46 0.36 0.07 2.70
23.72 10.77 0.31 0.31 1.00
24.10 11.00 0.38 0.27 0.71
Titration of Acetic Acid against NaOH
Vbase (mL) pH of Solution ΔV (mL) ΔpH ΔpH
ΔV (mL)-1
0.00 3.08
0.91 3.37 0.91 0.29 0.32
2.00 3.67 1.09 0.30 0.28
3.95 3.99 1.95 0.32 0.11
6.83 4.30 2.88 0.31 0.11
10.28 4.59 3.45 0.29 0.08
14.20 4.90 3.92 0.31 0.08
17.85 5.26 3.65 0.36 0.90
19.32 5.48 1.47 0.22 0.15
20.90 5.80 0.77 0.32 0.42
21.53 6.07 0.63 0.27 0.43
22.00 6.40 0.47 0.33 0.70
22.21 6.67 0.21 0.27 1.30
22.60 9.10 0.39 0.33 0.85
23.70 9.75 1.10 0.65 0.59
25.00 10.37 1.30 0.62 0.48
26.20 11.00 1.20 0.63 0.53
Analysis:
pH vs. VNaOH
pH vs. VNaOH
Initial concentration of HCl based on titration with pH meter
[NaOH] = 0.2200M
Volume of NaOH to reach equivalence = 32.08mL
32.08mL × L × 0.2200mol = 0.007058mol of NaOH
1000mL L
[OH-] = [H+] = [HCl]
Volume of HCl solution = 100mL
0.007058mol = 0.07058M
0.100 L
Initial concentration of HCl based on titration with pH meter
[NaOH] = 0.2200M
Volume of NaOH to reach equivalence point = 32.02mL
32.02mL × L × 0.2200mol = 0.007044mol of NaOH
1000mL L
[OH-] = [H+] = [HCl]
Volume of HCl solution = 100mL
0.007044mol = 0.07044M
0.100 L
Vbase (mL)
2.7
ΔpH
ΔV (mL)-1
pH vs. VNaOH
OVERVIEW:
1. Sometimes chemists use indicators that change color at a pH of 9 to 10 instead of a pH of 7 because not all acid-base titrations reach their equivalence pints at a pH of 7 when they involve a weak acid or a weak base. For example in the titration of a weak acid with a strong base, like the one we just did of HOAc against NaOH, the equivalence point was at a pH around 9 which is why we used phenolphthalein as an indicator because it changes from colorless to pink over a pH of 8.2 to 10.
2. In the titration of HCl against NaOH we used pH= 7 as the equivalence point because it was the titration of a strong acid against a strong base so when the two were at the equivalence point, all of the OH- from the NaOH neutralized with all of the H+ from the HCl and the concentrations are equal resulting in H2O. All of the HCl went into forming H+ and all of the NaOH went into forming OH- to react with the H+. In contrast, with the Acetic Acid, not all of the HOAc went into forming H+ ions.
This lab called for very careful measurements which was difficult to do and resulted in data that was not as accurate as it could be. Going one drop over the equivalence point had a huge impact on the titration especially since the pH changes so radically near the equivalence point. This made tracking the pH with the pH meter near the equivalence point difficult. These drastic pH changes are clearly visible on our graphs. We had some difficulty in trying to track the change in volume with each change of 0.3 in pH because the pH meter slow in reaction time in adjusting to the pH. Often times we would add NaOH until there was a 0.3 change in pH and when we then stopped to record the change in volume the pH would continue to rise. Also, because of the drastic change in pH near the equivalence point, there were huge fluctuations in the pH meter near this point (after adding only 0.10 the change in pH was 1.39 and after adding 0.02mL of NaOH, the change in pH was 1.22). Initially, there would be a huge change in pH and then as the added NaOH dispersed throughout the solution, the pH would drop. For this reason there are some inconsistencies in our data, especially with the titration of HCl and NaOH with the pH meter. The most noticeable is the change in pH from 3.43 to 6.07 in which the volume changed from 22.60 to 22.85. This had a minimal impact on the shape of the pH vs. volume graph, but did affect the shape of the ΔpH/ΔV vs. volume graph. Instead of a slow steady rise followed by a large, drastic rise and fall, the graph had a slight peak followed by a slight drop followed by the huge rise. I think we should have allowed more time for the pH meter to stabilize before continuing with the titration. However, our data for the titration of NaOH and Acetic Acid was more consistent which shows that as the experiment when on, we became better at working with the pH meter. In addition, because we were working with a weak acid and strong base which created a buffer solution, the pH did not change as rapidly as with the strong acid.
According to the ΔpH/ΔV vs. volume graph our equivalence point was at 32.02mL which is at a pH of 9.10. The equivalence pH was supposed to be at 7, when the added OH- from the NaOH neutralized all of the H+ from the HCl. This difference could result from the fact that we did not wait long enough for the pH meter to stabilize to get the final reading. Had our ability to control the addition of NaOH been better, I do not think we would have had such issues. However, the shape of the graph was correct and shows a steep rise in the region from a pH of about 6.00 to 10.00.
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