Chemistry
Chemistry
10th Edition
ISBN: 9781305957404
Author: Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Publisher: Cengage Learning
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Hello, based on this graph and table, 

how can i solve for to find experimental pKa using the half equivalence point? and Ka of acetic acid?  

 

This table presents data on the addition of sodium hydroxide (NaOH) to a solution and the corresponding changes in pH and delta pH (dpH). ### Table: NaOH Addition and pH Measurements | NaOH added (mL) | pH | dpH | |-----------------|----|-----| | 0.5 | 3.9| null| | 1.0 | 4.1| 0.2 | | 1.5 | 4.2| 0.1 | | 2.0 | 4.3| 0.1 | | 2.5 | 4.6| 0.3 | | 3.0 | 5.1| 0.5 | | 3.5 | 5.3| 0.2 | | 4.0 | 6.2| 0.9 | | 4.2 |10.5| 4.3 | | 4.4 |11.0| 0.5 | | 4.6 |11.6| 0.6 | | 4.8 |11.7| 0.1 | | 5.0 |11.7| 0.0 | ### Explanation: - **NaOH added (mL)**: This column lists the volume of sodium hydroxide in milliliters that has been added to the solution. - **pH**: This column lists the corresponding pH values of the solution after the addition of NaOH. The pH is a measure of the acidity or basicity of the solution, with values below 7 indicating an acidic solution, and values above 7 indicating a basic solution. - **dpH**: This column lists the change in pH between successive measurements. The value of "null" in the first row indicates that there is no previous value to compare to for a change in pH. The data shows a trend of increasing pH with the addition of NaOH, which is expected as NaOH is a strong base. There is a notable jump in pH from 6.2 to 10.5 when 4.2 mL of NaOH is added, indicating a sharp rise in alkalinity. This may suggest the completion of neutralization of any present acidic components around that
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Transcribed Image Text:This table presents data on the addition of sodium hydroxide (NaOH) to a solution and the corresponding changes in pH and delta pH (dpH). ### Table: NaOH Addition and pH Measurements | NaOH added (mL) | pH | dpH | |-----------------|----|-----| | 0.5 | 3.9| null| | 1.0 | 4.1| 0.2 | | 1.5 | 4.2| 0.1 | | 2.0 | 4.3| 0.1 | | 2.5 | 4.6| 0.3 | | 3.0 | 5.1| 0.5 | | 3.5 | 5.3| 0.2 | | 4.0 | 6.2| 0.9 | | 4.2 |10.5| 4.3 | | 4.4 |11.0| 0.5 | | 4.6 |11.6| 0.6 | | 4.8 |11.7| 0.1 | | 5.0 |11.7| 0.0 | ### Explanation: - **NaOH added (mL)**: This column lists the volume of sodium hydroxide in milliliters that has been added to the solution. - **pH**: This column lists the corresponding pH values of the solution after the addition of NaOH. The pH is a measure of the acidity or basicity of the solution, with values below 7 indicating an acidic solution, and values above 7 indicating a basic solution. - **dpH**: This column lists the change in pH between successive measurements. The value of "null" in the first row indicates that there is no previous value to compare to for a change in pH. The data shows a trend of increasing pH with the addition of NaOH, which is expected as NaOH is a strong base. There is a notable jump in pH from 6.2 to 10.5 when 4.2 mL of NaOH is added, indicating a sharp rise in alkalinity. This may suggest the completion of neutralization of any present acidic components around that
## Equilibrium Constant Experiment Data: NaOH and CH3COOH ### Experiment Overview This graph represents data collected from an experiment to determine the equilibrium constant of a reaction between sodium hydroxide (NaOH) and acetic acid (CH3COOH). ### Graph Description #### Title: Equilibrium Constant Experiment Data: NaOH and CH3COOH #### Axes: - **X-axis:** Volume of NaOH added (in mL) - **Y-axis:** pH Reading #### Data Points: The graph illustrates the relationship between the volume of NaOH added to a solution of acetic acid and the resulting pH of the solution. #### Interpretation: - At the beginning of the experiment (0.5 to around 3.5 mL of NaOH added), the pH increases gradually from approximately 4 to around 6. This phase represents the buffering region of the acetic acid solution. - A sharp increase in pH is observed between 3.5 mL and 4.5 mL of NaOH added, where the pH rises steeply from about 6 to 12.5. This significant increase indicates the equivalence point where the amount of NaOH added is stoichiometrically equivalent to the acetic acid present in the solution. - Beyond the equivalence point (greater than 4.5 mL), the pH levels off, stabilizing around 12 to 12.5, reflecting the excess NaOH in the solution. ### Conclusion: The equilibrium point in the graph is determined by the steepest rise in the pH values, indicating the reaction completion between NaOH and CH3COOH. Understanding this data is crucial for calculating the equilibrium constant and understanding the acid-base neutralization dynamics.
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Transcribed Image Text:## Equilibrium Constant Experiment Data: NaOH and CH3COOH ### Experiment Overview This graph represents data collected from an experiment to determine the equilibrium constant of a reaction between sodium hydroxide (NaOH) and acetic acid (CH3COOH). ### Graph Description #### Title: Equilibrium Constant Experiment Data: NaOH and CH3COOH #### Axes: - **X-axis:** Volume of NaOH added (in mL) - **Y-axis:** pH Reading #### Data Points: The graph illustrates the relationship between the volume of NaOH added to a solution of acetic acid and the resulting pH of the solution. #### Interpretation: - At the beginning of the experiment (0.5 to around 3.5 mL of NaOH added), the pH increases gradually from approximately 4 to around 6. This phase represents the buffering region of the acetic acid solution. - A sharp increase in pH is observed between 3.5 mL and 4.5 mL of NaOH added, where the pH rises steeply from about 6 to 12.5. This significant increase indicates the equivalence point where the amount of NaOH added is stoichiometrically equivalent to the acetic acid present in the solution. - Beyond the equivalence point (greater than 4.5 mL), the pH levels off, stabilizing around 12 to 12.5, reflecting the excess NaOH in the solution. ### Conclusion: The equilibrium point in the graph is determined by the steepest rise in the pH values, indicating the reaction completion between NaOH and CH3COOH. Understanding this data is crucial for calculating the equilibrium constant and understanding the acid-base neutralization dynamics.
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