Acidbasetitrationsupdated (1)
Download as DOC, PDF1 like542 views
The document summarizes an experiment on acid-base titrations. Sodium hydroxide was standardized and used to titrate hydrochloric acid, phosphoric acid, and an unknown acid. Titrations of hydrochloric acid showed the equivalence point was pH 7. Phosphoric acid titration showed two equivalence points, determining its two pKa values. Titration of the unknown acid found its molecular weight to be 76.09 g/mol, identifying it as tartaric acid. A final potassium hydrogen phthalate titration found the experimental curve nearly identical to the theoretical curve.
![weak acid in water (2)
. After base is added, the pH is determined by the equilibrium
between the protonated acid, and its deprotonated conjugate base. The pH of this
equilibrium is described in the Henderson-Hasselbach equation (4)
:
pH = pKa + log10 [A-]/[HA].
As the equivalence point of the titration is reached, the acid is completely converted to its
conjugate base, and the pH is determined by the hydrolysis reaction of the conjugate
base(5)
. Further addition of base past the equivalence point results in the domination of the
OH-
ion compared to all the other species in solution, so the pH is controlled by the
concentration of OH-
in the system (6)
.
The relationship between the pH and the amount of titrant added permits an
understanding of the equilibrium properties of the acid. At the half-equivalence point the
pH is the same as the pKa of the acid. This property is due to the fact that when one half
an equivalent of the base is added, the concentrations of the acid and its conjugate base
are the same. Thus, in the Henderson-Hasselbach equation the only term that determines
the pH value is the pKa term. After this determination, it is simple to calculate the
equilibrium constant of the acid.
In this lab, the base NaOH is used to titrate several different acids. This base is
standardized to a known concentration so that all the measurements are precise and
accurate. In the titrations, as the equivalence point nears, it is necessary to add very small
amounts of titrant. The reason for this is that nearly all of the acid has been deprotonated
and is in its conjugate base form. An addition of too much base could overshoot the point
at which the acid and the base are neutralized. Missing this point in the titration would
make the titration useless.
Experimental:
The details of this experiment can be found in the lab handout (1)
. The first step in
the experiment involved the standardization of the titrant, NaOH. This standardization
was performed by three titrations of the NaOH with known amounts of potassium
hydrogen phthalate (KHP). Before each of these titrations, the KHP was dissolved in
boiling water (50 ml) and phenolphthalein (3 drops) was added as the indicator. The
endpoint of the titrations was then determined by the colourometric method. The molarity
of the NaOH solution was then determined using the calculations found in Appendix 1.
The next part of this experiment was the titration of hydrochloric acid with the NaOH
stock solution, whose concentration had just been determined. To prepare it for the
titration, the stock HCl solution (12 M) was diluted to 3.0 * 10-3
mol in 25.0 ml (0.12 M).
During the titration, the pH of the solution was recorded for each addition of titrant using
a pH meter that had been standardized in a buffer solution with pH of 7.0. In addition,
bromocresol green was used as an indicator for the endpoint of the titration. A similar
titration was performed on 25.0 ml of phosphoric acid (0.075 M), except that the
colourometric method of determining the endpoint was not used: only the pH readings by
the meter were used for this determination. After this titration, three samples of an
unknown acid (ID letter “O”) were prepared in the same manner that the KHP samples
were prepared earlier. Each sample was titrated with the NaOH titrant until the endpoint
was reached, as determined by color change. From these three titrations, the identification
of the unknown acid was possible. A final titration was performed with a sample of KHP,](https://image.slidesharecdn.com/acidbasetitrationsupdated1-150424084021-conversion-gate02/85/Acidbasetitrationsupdated-1-2-320.jpg)
Acidbasetitrationsupdated (1)
- 1. Exercise 6: Acid-Base Titrations Nick Redmond Partner: Stephan Villavicencio Thurs. Section Lab date: March 19, 1998 Report date: April 6, 1998 Abstract: The equilibrium between acids and bases during a titration can be used to determine several characteristics of the acid or the base. Sodium hydroxide was standardized to 0.1035 M in three acid-base titrations of potassium hydrogen phthalate (KHP). This standardized NaOH solution was then used in a series of other titrations with acids in order to gain information about those acids. The first acid tested was hydrochloric acid. This titration showed that the equivalence point occurred at a pH of 7.0. The second acid tested was the polyprotic phosphoric acid. This titration resulted in two equivalence points, one at pH 5.01 and the second around pH 9.25. The pKa1 of this acid was 2.25 and pKa2 was 6.90, compared with the known values of 2.16 and 7.21, respectively. Three titrations with the unknown acid “O” showed that the molecular weight of the acid was 76.09 g/mol. This is approximately one-half the known molecular weight of the compound represented by “O,” tartaric acid. The final titration of KHP permitted a comparison between an experimental titration curve and a theoretical curve. These two curves were nearly identical. Introduction: Acid and bases are the means used to adjust the concentration of protons (H+ ) and hydroxide ions (OH- ) in an aqueous solution. The concentrations of these particular species are important because the reactivity of many compounds depends on whether or not the compound is protonated. The compound may be protonated when the concentration of H+ in the solution is high, due to the acid. It may be deprotonated when the concentration of OH- is high, as a result of the base. The relationship between these two ions revolves around the auto-ionization reaction of water(1) : 2H2O (l) ⇔ H3O+ (aq) + OH- (aq) This reaction is described by the equilibrium constant Kw, 1.01 x 10-14 , at 25o C. The auto- ionization equilibrium couples with the equilibrium between the acid and the base to give an aqueous acid-base system. The strength of the proton concentration, its acidity, is described in terms of pH. This value is measured using a specialized voltmeter, which is calibrated with a buffer solution of known pH (2) . The pH of a system at a particular point in a titration can also be determined by the use of an indicator, which is a buffer that changes color at a specific pH (3) . The pH of an aqueous system changes as an acid or a base is added to it. If a weak acid of known concentration is titrated with a strong base, the pH of the system rises with each addition of base. The amount that the pH rises, though, depends on how far the reaction has proceeded. Initially, the pH is determined solely by the dissociation of the
- 2. weak acid in water (2) . After base is added, the pH is determined by the equilibrium between the protonated acid, and its deprotonated conjugate base. The pH of this equilibrium is described in the Henderson-Hasselbach equation (4) : pH = pKa + log10 [A-]/[HA]. As the equivalence point of the titration is reached, the acid is completely converted to its conjugate base, and the pH is determined by the hydrolysis reaction of the conjugate base(5) . Further addition of base past the equivalence point results in the domination of the OH- ion compared to all the other species in solution, so the pH is controlled by the concentration of OH- in the system (6) . The relationship between the pH and the amount of titrant added permits an understanding of the equilibrium properties of the acid. At the half-equivalence point the pH is the same as the pKa of the acid. This property is due to the fact that when one half an equivalent of the base is added, the concentrations of the acid and its conjugate base are the same. Thus, in the Henderson-Hasselbach equation the only term that determines the pH value is the pKa term. After this determination, it is simple to calculate the equilibrium constant of the acid. In this lab, the base NaOH is used to titrate several different acids. This base is standardized to a known concentration so that all the measurements are precise and accurate. In the titrations, as the equivalence point nears, it is necessary to add very small amounts of titrant. The reason for this is that nearly all of the acid has been deprotonated and is in its conjugate base form. An addition of too much base could overshoot the point at which the acid and the base are neutralized. Missing this point in the titration would make the titration useless. Experimental: The details of this experiment can be found in the lab handout (1) . The first step in the experiment involved the standardization of the titrant, NaOH. This standardization was performed by three titrations of the NaOH with known amounts of potassium hydrogen phthalate (KHP). Before each of these titrations, the KHP was dissolved in boiling water (50 ml) and phenolphthalein (3 drops) was added as the indicator. The endpoint of the titrations was then determined by the colourometric method. The molarity of the NaOH solution was then determined using the calculations found in Appendix 1. The next part of this experiment was the titration of hydrochloric acid with the NaOH stock solution, whose concentration had just been determined. To prepare it for the titration, the stock HCl solution (12 M) was diluted to 3.0 * 10-3 mol in 25.0 ml (0.12 M). During the titration, the pH of the solution was recorded for each addition of titrant using a pH meter that had been standardized in a buffer solution with pH of 7.0. In addition, bromocresol green was used as an indicator for the endpoint of the titration. A similar titration was performed on 25.0 ml of phosphoric acid (0.075 M), except that the colourometric method of determining the endpoint was not used: only the pH readings by the meter were used for this determination. After this titration, three samples of an unknown acid (ID letter “O”) were prepared in the same manner that the KHP samples were prepared earlier. Each sample was titrated with the NaOH titrant until the endpoint was reached, as determined by color change. From these three titrations, the identification of the unknown acid was possible. A final titration was performed with a sample of KHP,
- 3. using the same method as in the phosphoric acid titration. The theoretical titration curve for KHP was then calculated and compared with the experimental titration curve for this final titration. Results and Discussion: The standardization of the NaOH involved three titrations with KHP. The endpoint was determined by the change in color from clear to pink, as caused by the phenolphthalein indicator around pH 8.0 (7) . The results from those titrations are summarized in the following table: 1st Run 2nd Run 3rd Run Mass KHP 0.8630 g 0.8054 g 0.8433 g Vol. Titrant 41.09 ml 38.09 ml 39.90 ml Molarity of NaOH 0.10285 M 0.10355 M 0.10348 M More confidence is held in the second two runs than in the first run because the color change for the first run was not as obvious as was expected. The equivalence point was reached when the slightest tinge of pink persisted in solution. On the first run, however, that slight tinge of pink was not recognized as such, and a very small amount of excess NaOH may have been added. This error was not repeated because for the second two runs because the slight tinge was expected for those runs. See Appendix 1 to view what effect this error had on the molarity determination. The concentration of the NaOH was 0.1035 M. The standard deviation for the three runs was 3.86 x 10-4 .
- 4. The titration of hydrochloric acid used 3.0 x 10-3 moles of HCl in 25.0 ml solution. The titration curve for the titration of HCl can be seen below: Figure 1: Titration Curve of HCl with NaOH (0.1035 M) as Titrant From this curve it can be seen that the equivalence point was reached when about 31 ml NaOH was added. Looking more closely at the data, the amount of titrant added at this point was determined to be 31.43 ml. This means 3.25 x 10-3 moles of NaOH were needed to react with 3.0 x 10-3 moles of HCl. Because HCl is a strong acid and NaOH is a strong base, the Ka is very high. Essentially all the protons in equilibrium come from the acid, and not from water. This means before the equivalence point, pH is determined almost exclusively by the strong acid, and after the equivalence point it is determined by the strong base. At the equivalence point, the two cancel each other out, and the pH is a neutral 7.0. There is no use trying to find a half-equivalence point to find the pKa of the acid because the acid’s Ka is too large. The experimental data supports these qualities of a strong base. Titration Curve 0 5 10 15 0 10 20 30 40 50 Amount of Titrant Added (ml) pH
- 5. The curve for the titration of phosphoric acid follows: Figure 2: Titration Curve of H3PO4 with NaOH (0.1035 M) as Titrant The curve above shows two different equivalence points, the first after the addition of 19.27 ml NaOH and the second after the addition of 38.59 ml. This is seen because phosphoric acid is a polyprotic acid that has three different pKa values, one for the removal of each of its three protons. In this experiment, only two of the three are removed. The removal of each successive proton becomes more difficult. The solution begins with H3PO4. As base is added, H2PO4 - and H3PO4 are found in the solution. Eventually, only H2PO4 - can be found in the solution, and the first equivalence point is reached. As shown by the Henderson-Hasselbach equation, the half equivalence point determines the pKa for the removal of this first proton. This pKa was determined to be 2.25, compared with a literature value of 2.16 (8) . Additional base does not result in an overwhelming of the system with OH- ions, as it does in monoprotic acids. The hydroxide ion is instead used to remove the second proton. This is more difficult than the first removal. The addition of base creates in the solution both H2PO4 - and HPO4 2- . Eventually, the solution only contains this latter ion, and the second equivalence point is reached. The amount of base needed to go from the first equivalence point to this second determines the new half equivalence point where H2PO4 - and HPO4 2- can be found in equal concentrations. The pH at that point is equal to the second pKa. Experimentally this value was 6.90, as compared with the known pKa2 of 7.21. The addition of more base is used to remove the third proton, but in this experiment that was not attempted for time considerations. Titration Curve 0 5 10 15 0 20 40 60 80 Amount of Titrant Added (ml) pH
- 6. The endpoints of the three titrations of the unknown acid “O” were determined by the color change due to the indicator phenolphthalein. The data from these three titrations is summarized in the following table: 1st Run 2nd Run 3rd Run Mass Unknown 0.3459 g 0.3569 g 0.3770 g Vol. Titrant 43.85 ml 45.38 ml 47.88 ml Molarity of NaOH 0.1035 M 0.1035 M 0.1035 M M.W. of Unknown 76.22 g/mol 75.99 g/mol 76.08 g/mol The average molecular weight from these three runs is 76.09 g/mol. The unknown acid “O” is known to be tartaric acid, which has a molecular weight of 150.09 g/mol, almost twice the value obtained experimentally (an experimental error of 49.3%). This difference may be explained by a study of the structure of tartaric acid. In its ionic form, it seems possible that the structure could be cleaved into two molecules each of molecular weight about 76 g/mol. This seems to be the only possible explanation for the discrepancy.
- 7. In the final titration of KHP, 9.065 x 10- 3 moles of KHP were diluted in 36.5 ml. The spreadsheet in Appendix 2 contains all the data necessary for the theoretical titration, as well as the necessary experimental data. The graph below is the experimental and theoretical titration curves for the final titration of KHP: Figure 3: Experimental and Theoretical Titration Curves of KHP with NaOH (0.1035 M) as Titrant From this graph it is obvious that the two are almost identical. For the theoretical plot, the concentration of the acid was 0.11337 M, determined from the equivalence point of the actual titration. The volume of acid used was 36.5 ml. The Ka of the acid was determined from the half-equivalence point of the actual titration, where the pH and pKa were 5.05. The Ka was therefore 8.91 x 10-6 . The concentration of the base was 0.1035 M. The similarities in the two curves confirms the success of the experimental titration of KHP, and all the methods used in that titration. Conclusion: The acid base titrations provided key insights into the behavior of both the acid and the base in their equilibria with each other and with water. It was observed that when a strong acid and a strong base react with each other in equivalent concentrations, the pH of the solution is neutral. It was also seen that the combination of weak acids with a base results in a buffered system whose pH only gradually changes with the addition of base. It was observed that this buffered system loses that quality as the protonated acid runs Experimental and Theoretical Titration Curves 0 5 10 15 0 20 40 60 Volume Titrant Added (ml) pH Series1 Series2
- 8. out. From the titration curves of these systems, the pKa of the acid was determinable. It was also shown that the molecular weight of an unknown acid could be determined if a few facts were known about the acid-base system. Finally, the theoretical titration curve could be computed using a complicated formula that incorporated all the aspects of the acid-base equilibrium. These observations and accomplishments demonstrate important qualities of the acid-base equilibrium. These qualities have real affects on real systems in the world. The ability to predict Ka, pH, molecular weight, or any of the other values that were used in this experiment gives great power and understanding to the person trying to understand the chemical systems around him. References: 1. Chem 282 Lab Handout: “Acid-Base Titrations,” p1. 2. Chem 282 Lab Handout: “Acid-Base Titrations,” p2. 3. Harman, W. Dean, Chem 282 Lecture 4. Chem 282 Lab Handout: “Acid-Base Titrations,” p3. 5. Chem 282 Lab Handout: “Acid-Base Titrations,” p4. 6. Chem 282 Lab Handout: “Acid-Base Titrations,” p5. 7. Olmstead, Chemistry: The Molecular Science, p804. 8. Olmstead, Chemistry: The Molecular Science, pA-27.