Finding The Pka of An Unknown Acid Indicator Using Spectrophotometry
The goal of this experimental process was finding the pKa of an unknown acid indicator using spectrophotometry. Through a process of calculations the pKa of the unknown acid indicator was found. Using wavelengths from 350 nm to 650 nm absorbances were found for low, medium and high pH ranges. With known data about the unknown indicator that the pKa was found using two different methods of calculations.
Cramer's Rule was accessed, and the average pKa was determined to be _#.3610 with a standard deviation of _ at a 95% confidence interval. This method produced an error of _. Furthermore, assessment with Beer's Law with the same indicator produced a __ with a __ error at the same confidence level.
The _ method produced a more accurate pKa value of ___. Finally, it was concluded that the unknown acid indicator was _, with a literature value pKa of_, because the pKa value of_ matched the color change of the concluded indicator.
The unknown indicator solution( ) to be experimented with was retrieved from the stock room, along with the proper devices and materials for the experiment. To measure pH effectively the pH meter underwent a two point calibration process, making sure that it was reading the correct pH for the next steps. pH of 4 and 10 was used to calibrate the device. A 100 mL beaker was set up with 1 mL of the unknown indicator solution, 25mL of acetic acid (HC2H3O2) and a magnetic stirrer. An identical solution was also made to be used as a visual control.
The solution beaker to be changed was placed on a magnetic stirring device at a relatively low rate of stir. A pH meter was placed into the solution. 1M NaOH was added drop by drop to the solution until the color and the pH of the solution was recorded. The color was recorded for every 0.1 change in pH until the pH of the solution reached 11. The pH level where color change was very noticeable was recorded.
Three volumetric flasks were retrieved. Using the important values noted in the color change recording three buffered solutions were made, a low pH, a high pH, and an intermediate pH solution. For each buffer solution 50 mL of monopotassium phosphate KH2Po4 was added along with either 1M HCl or 1M NaOh to get to the pH required from the important pH values.
In separate but identical procedures, each buffer solution was transferred to a 100 mL volumetric flask, and 3.00 mL of the original indicator solution was added. The flask was filled to the 100 mL mark with deionized water.
eionized water was added until the solution was slightly below the 100 mL mark. To ensure that the solution is exactly 100 mL, a medicine dropper was used to add more deionized water. Each time the solution was made to a 100 mL it was swirled. Each buffer was transferred into a 250 mL erlenmeyer flask. Final pH readings were recorded and a drop of CCl4 was dropped in to prevent bacteria and fungal growth.
Next week, cuvettes and a spectrophotometer were prepared. Two cuvettes were both lined twice with deionized water, one was filled with deionized water, the other was lined twice with the first buffer solution to be tested and filled with the solution. Before placing the cuvettes into the machine they were wiped with kim-wipes to make sure the glass was as clear as possible. The spectrophotometer was zeroed using the solution of just water at 350 nm wavelength. The cuvette with the solution was then placed into the holder. The absorbance was measured using a spectrophotometer as a function of the wavelength from 350 to 650 nm at 25 nm intervals. Using the maximum absorbances found a more precise scan was done around the maximas at 5nm intervals. The wavelength of the absolute maximas of absorbance were recorded. This process was repeated again for the other two buffer solutions.
By diluting the solutions by their concentration of indicator at 1c, 0.8c, 0.6c. 0.4c, and 0.2c, absorbance of each was measured. These measurements were calculated and graphed to notice slope and any deviations. This was done for only the low and high buffer solutions only.
he data and processes established in this experiment led to the calculation of the pka of the unknown indicator to be ____. Although data was collected in similar ways the final calculations came down to two different methods, Cramer’s rule and Beer’s law. Cramer’s law produced a pKa value of ____with a standard deviation of __and a ___ error at a 95% confidence interval. Beer's Law produced an average pKa value of __with a ___error at the same confidence interval. Although both proccessizes exported similar results ________ is perhaps more accurate.
“Using the data corresponding to group 6, the pKa of the unknown indicator was determined to be 4.10 through the calculation of two methods. Using method one, Cramer's Rule, produced a pKa value of 4.518 1.3610 with a standard deviation of 0.379 and a 10.20% error at a 90% confidence interval. Method two, Beer's Law, produced an average pKa value of 3.733 with a 9.02% error. Comparing the calculated pKa values to the pKa of the determined pKa of 4.10, the second method was found to be more accurate. The second method should be more accurate as it takes the slope of the charts graphed using Beer's Law (Figures 2 and 3), which has a high coefficient of determination value (0.9995 and 0.9996 respectively). The coefficient of determination is a statistical measure that represents the proportion of variance for the dependent variable, absorbance, that's explained by the independent variable, concentration. in a model. The closer the coefficient of determination is to 1, the more accurate the graph is. Although there could be systematic errors that could affect the results of the experiment, the experiment proved to have similar results in multiple methods.”
2. Use Le Chatelier’s principle to explain why each form of the indicator is dominant at relatively low pH or at high pH.
Because the addition of acids or bases change the equilibrium of a system Le Chatelier’s principle is applicable to this scenario. For example, when HCl was added into the indicator solution in this experiment, which is an acid due to its H+ content, HIn will have more dominance in the solution. Conversely, when NaOH was added to the solution in this experiment In- became more dominant. This is because of the lack of hydrogen ions. When there is a loss or gain of hydrogen ions the equilibrium is shifted and that is what causes the changes of pH. Thus, at low or high pH there will be a greater equilibrium shift which results in a more profound color change.
3. Why doesn’t moderate dilution of a buffered solution change its pH?
A moderate dilution will not have much of an effect on the pH of the solution due to the fact that a buffered solution involves the slight back and forth between acids and bases and their conjugates. The ratio of conjugate base to weak acid in this solution would not have much of a change thus pKa will remain constant.
4. Clearly explain the existence and significance of the isobestic point
The isobestic point helps define a unique relationship between the absorbance and wavelength of a substance. The point itself is the point in which the total absorbance does not change, it will be at a specific wavelength. It occurs when there is a ratio of whole number integers between the concentrations of the different substances. This is because the same absorbance is directly linked to the same concentration.
5. Why don’t we have to find the value of b?
This experiment involved the same glassware for the three different solutions. Because the same cuvette was washed and lined each time another cuvette was not required. In addition, this cuvette’s size remained the same thus the value of b, the diameter of the cuvette, was constant.
Some student investigators have noted that they have obtained zero, negative, or indeterminate values for the concentration ratio at certain wavelengths. Why does this occur? Should these values be used in determining the average pKa? Why, or why not? If you found this to be the case for your own experiment, discuss the reasons why this occurred.
Generating buffered solutions of varied pH ranging from low, intermediate, and high pH values, was crucial to this experiment's success. The pH’s to be used in the buffered solutions was determined by using the unknown indicator and a pH meter along with basic solution NaOH to increase the pH. By checking when the indicator had a major color change due to the pH increase were also the points to be used in the buffering solutions. This simple process does have a layer of error. The human eye is not perfect and seeing color changes in solutions is not always an easy task. Also, the pH increments were pretty general, for a more exact pH range perhaps a smaller pH increment would help.
In this experiment the buffered solutions are used to calculate absorbances at certain wavelengths. That is why it is crucial that the pH range is somewhat accurate so accurate results can be produced. Therefore, a two point calibration process was used in order to achieve more exact results. The two point calibration process helps make sure the machine is outputting the right data. It was found that the values were slightly different before and after the two point calibration. It was also found that the calculated pH and the measured pH of the buffered solutions were slightly off for the low and intermediate buffer solutions. This is important to recognize but its effect on throwing off the experiment is minimal because pH is a range and this indicator’s range is unique. General trends will still be able to be seen with the absorbances.
After measuring absorbances at different wavelengths for all 3 buffered solutions, a key step was taken. A fine scan around the areas that had maximas was taken. This helped narrow down the actual wavelengths that the maximum absorbances were occurring. pKa of the indicator was found using two different methods. Method one, Cramer’s Rule, utilized absorbances rather than the concentrations of the solutions and resulted in a pKa value of ________ at a 95% confidence interval. Beer’s Law, the second method, was more scoped in on the concentrations of the solutions. After taking data on the absorbances of the high and low solutions at different dilutions with water, pKa can be determined graphically. Beer’s Law produced an average pKa value of _______ at a 95% confidence interval. ( ). The literature pka value was determined to be ____. The approach taken in Cramer’s Rule (method one) relies heavily on the experimental calculations. This method uses something similar to a system of buffers to conclude the [HIN] and [In-].mAs it suggests, the values that it gets are quite experimental and heavily based on what steps were taken by the user. On the other hand, Beer’s Law is less experimental due to the fact that concentrations come into play to check absorbances and create a functional graph. When it is known how far the light has traveled through the sample then the concentrations are directly correlated to the absorbance.
Although both methods are successful in determining values close to the literature pKa values, method __ proved more successful. The lack of concentration diversity in method one may have resulted in the different results. Method two overall has more data points and allows any outliers to be rooted out/ have less effect on the results. Overall, method one’s pKa value was about ____ units higher/lower than the literature value while method two’s pKa value was about ___ units lower/higher.
There are multiple reasons that could have resulted in the standard deviations for both methods as well as the reasoning behind the error from the literature pKa value. Specifically for Beer’s method, dilution of the solutions may have created a systematic error. It may have become more acidic or basic after each time. Thus, there may be slight changes in measurement of the absorbances depending on how badly they were altered. Despite this, moderate dilution of these solutions does not really change the pH of the solutions due to the fact that they are buffered solutions. Because these buffered solutions involve a weak acid and its conjugate base when deionized water is added, the acidic and basic parts of the solution change equally. This is due to the fact that water is made up of hydrogen ions and hydroxide, the same as the respective acid and the base are in the solution. It should be noted that the pH values were most likely correct due to the fact that there were not any outliers. This does not suggest perfection but it does suggest that the procedure is unique and appropriate for achieving pKa values of an unknown indicator.
In addition, experimental error may have come from other factors. Measuring absorbance using cuvettes in general is an error bound experiment. This is due to the fact that the cuvettes are extremely sensitive in the machine. Data may have been skewed for multiple reasons due to the cuvettes: there may have been a scratch or a fingerprint on the glass that is difficult to see; or when the the cuvettes were being lined or washed it may have been done improperly. Changes to the lower portions of the cuvette can have big impacts on its optical usage in the spectrophotometer. In addition it may have been that some fungal or bacterial aspects developed in the solution during the times it was left unattended for a week. A drop of CCl4 was dropped into the solutions to prevent this but these long gaps may have altered the solutions thus altering the measurements of their absorbances. However, the consistency of the data suggests these errors were most likely minimal. Measurement errors produced random errors too. The spectrophotometer may not always be accurate, there may be minor errors in the concentrations for each solution, and inaccuracy in measurement of pH.
Both methods proved useful in finding an accurate naming of this unknown indicator. This indicator, by the processes described here is ________ with pKa of ______. By comparing the data in this experiment to those of the literature values it is clear that the experiment was overall a success despite some of the errors mentioned taking roles in the deviations.
Finding The Pka of An Unknown Acid Indicator Using Spectrophotometry
The solution to the unknown indicator solution ( ) was found in the stock room, and the corresponding equipment and materials for the experiment. The pH meter was calibrated to two points during this procedure, ensuring reading the proper pH for the following procedures. A pH of 4 and 10 were chosen for calibration of the device. A 100 mL beaker was prepared with 1 mL of unknown indicator solution, 25mL of acetic acid (HC2H3O2), and a magnetic stirrer. An identical solution was also produced to serve as a visible control.
The solution beaker to be changed at a slow speed was put on a magnetic stirring device. The pH of the solution was measured with a pH meter. 1M NaOH was added to the solution until its color and pH were documented. The color was noted for every 0.1 change in pH until the solution’s pH reached 11. The pH level at which color changes were most apparent was recorded.
The second step in the process was to remove three volumetric flasks. Three buffered solutions were produced using the crucial values noted in the color change recording, including a low pH, high pH, and intermediate pH solutions. 50 mL of monopotassium phosphate KH2Po4 was added to each buffer solution with either 1M HCl or 1M NaOH to reach the required pH values.
The buffer solution was transferred to a 100 mL volumetric flask in two separate operations. 3.00 mL of the original indicator solution was added to each buffer solution, which had been transferred similarly to a 100 mL volumetric flask. The water level of the flask was brought up to 100 mL with deionized water.
Deionized water was gradually added until the solution reached approximately 90 mL. A medicine dropper was used to add more deionized water to increase the volume to exactly 100 mL. The solution was mixed for 100 mL each time. Each buffer was poured into a 250 mL Erlenmeyer flask. Final pH readings were taken, and a drop of CCl4 was added to prevent bacterial and fungal development.
A preparation of cuvettes and a spectrophotometer was done the following week. One cuvette was filled twice with deionized water, and one was lined twice with the first buffer solution to be tested, after which it was filled again with the solution. To ensure that the cuvettes were as clean as possible, they were cleaned with kim-wipes before being placed into the machine. The solution of only water at 350 nm wavelength was used to zero the spectrophotometer. The cuvette with the solution was put into the holder. The absorbance was determined using a spectrophotometer as a function of wavelength from 350 to 650 nm at 25-nanometer intervals. The greatest absorbances were identified and used in a more precise scan around the maximums at 5nm distances. The wavelength of the absolute maxima of absorbance was recorded. This procedure was repeated for the other two buffer solutions one last time.
To test the accuracy of the absorbance readings, 0.1ml of each solution was diluted with 10x concentrated indicator to make 1c, 0.8c, and 0 percent concentration indicators were prepared by diluting them by their concentration of indicator at 1c, and so on. Absorbance was measured for each after being diluted with increasing concentrations.
This experiment required generating buffered solutions of multiple pH levels, ranging from low to high. The appropriate pH values were determined for the buffered solutions using the unknown indicator, a pH meter, and standard solution NaOH to raise the pH. The points used in the buffering solutions were determined by observing when the indicator’s major color change was due to a pH increase. This straightforward procedure has some room for error. The human eye is not flawless, and detecting color changes in solutions isn’t always straightforward. Furthermore, the pH increments were rather wide-ranging; a smaller pH increment may have helped for a more precise pH range.
The buffered solutions are utilized to determine absorbance at particular wavelengths in this experiment. As a result, the pH range must be correct for good results to emerge. As a result, a two-point calibration approach was used to get more precise outcomes. The two-point calibration procedure ensures that the machine is sending the correct data. It was discovered that the values were somewhat different before and after the two-point calibration. The low and intermediate buffer solutions had a measured pH that was slightly higher than their calculated pH, while the high buffer solution had a measured pH that was significantly lower than its calculated pH. This is significant to recognize, but the impact on the experiment is minor since pH is a range and this indicator’s range is distinct. With the absorbances, general tendencies will still be visible.
After taking absorbances at various measurements for all three buffered solutions, a crucial phase was completed. A fine scan of the regions with maximums was undertaken. This assisted us in determining the actual wavelengths where the maximum absorbances were found. The pKa of the indicator was determined using two distinct approaches. The first approach, Cramer’s Rule, used absorbances rather than solution concentrations and produced a pKa value of ________ with a 95% confidence range. The second approach, known as Beer’s Law, focused more on the solutions’ concentrations. The pKa of a solution can be determined graphically after taking absorbance data for high and low solutions at various dilutions with water. At a 95% confidence interval, the average pKa value of ‘Beer’s Law’ was ________. ( ). The literature pka value was found to be ____. Cramer’s Rule (method one) utilizes experimental calculations significantly. The [HIN] and [In-].mAs the name implies, when Cramer’s Rule is used, it works with a system of reservoirs to determine the [HIN] and [In-]. The values it arrives at are quite experimental, as they are heavily influenced by what the user took actions. The Beer’s Law equation (Equation 7) is more experimental than Beer’s Theory since it involves concentrations to check absorbances and generate a functional graph. When the distances traveled by the light through the sample are known, the concentrations are directly linked to the absorbance.
Both techniques effectively estimate pKa values that are close to those found in the literature, but method __ was more successful. The variety of concentrations in method one may have produced varied outcomes. Overall, method two has more data points and is better suited to removing outliers or reducing their impact on the results. Overall, the pKa value of method one was about ____ units greater/lesser than the literature value, whereas the pKa value of method two was about ____ lower/higher.
Several factors might have caused the standard deviations for both methods and the rationale behind the error from the literature pKa value. Dilution of solutions may have generated a systematic error in Beer’s method, for example. Each time exposed to air, it may have gotten more acidic or alkaline. As a result, the absorbances might vary somewhat depending on how much they were altered. Even though these treatments are buffered, moderate dilution of them does not significantly alter the pH of the solutions since they are buffered solutions. Because deionized water is added to buffered solutions, the acidic and basic portions of the solution change equally. This is a result of water being composed of hydrogen ions and hydroxide, just like the acid and base in the solution. The pH values were most likely correct since there were no outliers. This does not imply that the procedure is perfect, but it does indicate that it is distinctive and suitable for achieving pKa values of an unknown indicator.
Other sources of error are conceivable. Absorbance measurements with cuvettes are subject to experimental error, so the experiment was restricted. This is since cuvettes are extremely sensitive in the laboratory. Due to the cuvettes, data may have been skewed for various reasons: there could be a scratch or a fingerprint on the glass that is difficult to detect, or when the cuvettes were being lined or washed, they could have been cleaned incorrectly. Changes to the cuvette’s lower portions may significantly affect its optical usage in the spectrophotometer. It’s conceivable that some fungal or bacterial components arose in the solution during the week when it was left unattended. Due to the long pauses in the data that occurred because of this, it’s difficult to determine whether errors influenced their results. However, careful statistical analysis suggests that these mistakes were most likely minor. Second, the pH value inaccurately measured by the spectrophotometer may result in random errors. There could be minor variations in solution concentrations, and pH accuracy might be an issue.
The pKa value of each of the samples was determined, and then they were submitted to additional tests. Both approaches proved beneficial in determining an accurate name for this unknown indicator. This indicator has a pKa of ____ and is______ with respect based on the methods described here. Despite some of the mistakes discussed causing issues, the pKa values determined in this experiment were close to the literature value.
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