Hypothesized Correlation Between Ionization And Ionic Acid
The purpose of this experiment was to observe the hypothesized correlation between ionic strength and degree of ionization of a weak acid, while determining the factors that would impact the degree of ionization using measurements of pH. It was initially required to prepare standard 20-mL solutions of varying concentration of acetic acid for use in the rest of the experiment. One solution of CH3COOH was supposedly developed to be 2.00M, however it is highly possible that the solutions of CH3COOH were not exactly the expected concentration, evident by the differences in the percent ionization of the no salt added solutions.
In order to calculate the degree of ionization of the weak acid, the varyingly concentrated acetic acid solutions and …show more content…
pH measurements were used. An uncertainty of 0.01g was appropriate for the mass measurements because the precision balance measured up to the second decimal place. An uncertainity of 0.10 mL was appropriate for the volume measurements because the buret measured up to the hundreths place and the initial and final measurements of liquid in the buret added both 0.05 mL uncertainties to be 0.10 mL. An uncertainity of 0.01 was appropriate for the pH measurements because the pH probe used measured up to the second decimal place.
In performing the procedure, the results developed a consistent trend as evident within the graphs.
In the outcome of the experiment, when the ionic strength of the acetic acid solution increased with respect to concentration, the pH relatively decreased and percent of ionization decreased. However the data points had a low correlation to the line of best fit with an R2 of 0.6762 and resembled an inverse curve. The most directly proportional relationship between the ionic strength and percent ionization was determined to be an inverse of the square root of ionic strength: α∝1/√I, where α is percent ionization and I is ionic strength (M). The graph between percent ionization and the inverse square root of ionic strength has a nearly perfect R2 value of 0.989, suggesting that this relationship is the most accurate. The observed relationship between percent ionization and ionic strength is confirmed by Ostwald's dilution law, which suggests that α=√(K_a/C), where percent ionization has a directly proportional relationship with the inverse square root of concentration. Since concentration was the primary variant in this first no salt added experiment, the data collected is quite accurate and helps confirm the correlation explained in Ostwald's law between ionic strength and percent
ionization.
In the outcome of the second experiment, when the ionic strength of the acetic acid electrolyte was increased due to added salt over time, the pH decreased and percent of ionization increased as higher concentrations of salt were added. As evident in the graphs of a 2.00M CH3COOH solution, there is a strong, positive correlation between ionic strength and percent ionization, as the lines of best fit have positive slopes and relatively high correlation coefficient values. The R2 values of both trendlines, 0.9837 and 0.9593, suggest that the data points measured have a very strong correlation to the line of best fit, confirming that a linear fit is the best approximation of this evident trend regarding salts.
Since there was no value discovered in the literature regarding the degree of ionization with the addition of neutral salts, the accuracy of the data cannot be confirmed and percent error cannot be calculated. However, due to the consistency of the trends within each change of the ionic strength, it can be concluded that increases in the ionic strength lead to increases in the degree of ionization. In order to determine which salt added had the greatest impact on the degree of ionization, the slopes of the lines of best fit must be compared. In comparing the slopes of the lines of best fit, it can be determined which change in ionic strength incurred the greatest change in degree of ionization. The line of best fit at 2.00M CH3COOH for varying KCl concentration had a 0.0851 slope and for varying CaCl2 concentration had a 0.694 slope. The comparison of slopes suggests that the salt with a higher cation charge results in a larger degree of ionization. Although no literature value could be compared with the experimental value, it is still possible that the experiment contained systematic and random error.
Evaluation of Procedure:
Overall, the procedure was well-followed, however there were some possible errors within the experiment. One possible source of error was that there were particles left inside in the beakers used in the experiment. These possibly could have altered the measurement of the pH using the probe as any excess ions in the solution would affect the value recorded. In addition, any water residue in the flasks could have affected the molarity of the acetic acid, altering the results of the degree of ionization and pH. Another possible source of error arose from the instruments used and their random error. Primarily, the pH probe showed an extensive amount of error. Although it recorded to two decimal places, it was extremely challenging to obtain a precise measurement due to the continued fluctation of values on the pH probe reader. After calibrating the pH probe, it worked for a short period of time until it lost its calibration. This could mean that all the pH values calculated were slightly incorrect. In addition to the probe, there was random error in the analytical balances which might have led to discrepancies between the actual value and measured value. Finally, a source of uncontrollable error was the existence of a large period of time elapsed in performing the experiment. Some of the acetic acid solutions were produced in a laboratory period one day prior to the pH measurements. This period of time elapsed could have allowed evaporation or a change in the concentration of the acids, skewing the pH measurements.
Improving the Investigation:
One possible improvement to the experiment would be to use more precise instruments. Using a more precise pH probe would help eliminate the constant flux experienced when measuring the pH of the solutions. In addition, using a more precise balance would have provided more accurate measurements of the salts added to the solutions, to ensure an accurate concentration.
Another possible improvement would be to increase the number of trials performed. In the experiment, only one trial of each acid and salt concentration pair was performed, resulting in the inability to determine precision of values. With a more consistent and longer time period to perform the experiment, more trials could be completed and result in more consistent data points.
One final possible improvement would be to work towards better controlling the controlled variables during the entire experiment to ensure the environmental conditions were exactly the same. By maintaining constant temperature, pressure, and other factors, the experiment might have more accurate results.
In order to calculate the degree of ionization of the weak acid, the varyingly concentrated acetic acid solutions and …show more content…
pH measurements were used. An uncertainty of 0.01g was appropriate for the mass measurements because the precision balance measured up to the second decimal place. An uncertainity of 0.10 mL was appropriate for the volume measurements because the buret measured up to the hundreths place and the initial and final measurements of liquid in the buret added both 0.05 mL uncertainties to be 0.10 mL. An uncertainity of 0.01 was appropriate for the pH measurements because the pH probe used measured up to the second decimal place.
In performing the procedure, the results developed a consistent trend as evident within the graphs.
In the outcome of the experiment, when the ionic strength of the acetic acid solution increased with respect to concentration, the pH relatively decreased and percent of ionization decreased. However the data points had a low correlation to the line of best fit with an R2 of 0.6762 and resembled an inverse curve. The most directly proportional relationship between the ionic strength and percent ionization was determined to be an inverse of the square root of ionic strength: α∝1/√I, where α is percent ionization and I is ionic strength (M). The graph between percent ionization and the inverse square root of ionic strength has a nearly perfect R2 value of 0.989, suggesting that this relationship is the most accurate. The observed relationship between percent ionization and ionic strength is confirmed by Ostwald's dilution law, which suggests that α=√(K_a/C), where percent ionization has a directly proportional relationship with the inverse square root of concentration. Since concentration was the primary variant in this first no salt added experiment, the data collected is quite accurate and helps confirm the correlation explained in Ostwald's law between ionic strength and percent
ionization.
In the outcome of the second experiment, when the ionic strength of the acetic acid electrolyte was increased due to added salt over time, the pH decreased and percent of ionization increased as higher concentrations of salt were added. As evident in the graphs of a 2.00M CH3COOH solution, there is a strong, positive correlation between ionic strength and percent ionization, as the lines of best fit have positive slopes and relatively high correlation coefficient values. The R2 values of both trendlines, 0.9837 and 0.9593, suggest that the data points measured have a very strong correlation to the line of best fit, confirming that a linear fit is the best approximation of this evident trend regarding salts.
Since there was no value discovered in the literature regarding the degree of ionization with the addition of neutral salts, the accuracy of the data cannot be confirmed and percent error cannot be calculated. However, due to the consistency of the trends within each change of the ionic strength, it can be concluded that increases in the ionic strength lead to increases in the degree of ionization. In order to determine which salt added had the greatest impact on the degree of ionization, the slopes of the lines of best fit must be compared. In comparing the slopes of the lines of best fit, it can be determined which change in ionic strength incurred the greatest change in degree of ionization. The line of best fit at 2.00M CH3COOH for varying KCl concentration had a 0.0851 slope and for varying CaCl2 concentration had a 0.694 slope. The comparison of slopes suggests that the salt with a higher cation charge results in a larger degree of ionization. Although no literature value could be compared with the experimental value, it is still possible that the experiment contained systematic and random error.
Evaluation of Procedure:
Overall, the procedure was well-followed, however there were some possible errors within the experiment. One possible source of error was that there were particles left inside in the beakers used in the experiment. These possibly could have altered the measurement of the pH using the probe as any excess ions in the solution would affect the value recorded. In addition, any water residue in the flasks could have affected the molarity of the acetic acid, altering the results of the degree of ionization and pH. Another possible source of error arose from the instruments used and their random error. Primarily, the pH probe showed an extensive amount of error. Although it recorded to two decimal places, it was extremely challenging to obtain a precise measurement due to the continued fluctation of values on the pH probe reader. After calibrating the pH probe, it worked for a short period of time until it lost its calibration. This could mean that all the pH values calculated were slightly incorrect. In addition to the probe, there was random error in the analytical balances which might have led to discrepancies between the actual value and measured value. Finally, a source of uncontrollable error was the existence of a large period of time elapsed in performing the experiment. Some of the acetic acid solutions were produced in a laboratory period one day prior to the pH measurements. This period of time elapsed could have allowed evaporation or a change in the concentration of the acids, skewing the pH measurements.
Improving the Investigation:
One possible improvement to the experiment would be to use more precise instruments. Using a more precise pH probe would help eliminate the constant flux experienced when measuring the pH of the solutions. In addition, using a more precise balance would have provided more accurate measurements of the salts added to the solutions, to ensure an accurate concentration.
Another possible improvement would be to increase the number of trials performed. In the experiment, only one trial of each acid and salt concentration pair was performed, resulting in the inability to determine precision of values. With a more consistent and longer time period to perform the experiment, more trials could be completed and result in more consistent data points.
One final possible improvement would be to work towards better controlling the controlled variables during the entire experiment to ensure the environmental conditions were exactly the same. By maintaining constant temperature, pressure, and other factors, the experiment might have more accurate results.