Exp. 8 - Determining the Equivalent mass and Ka of an unknown acid PDF

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Experiment 8

Determining the Equivalent mass and Ka of an unknown Acid Introduction

Acids undergo dissociation in water and establish an equilibrium according to the following general equation: HA(aq) + H2O (l)

H3O+ (aq) + A- (aq)

For acetic acid, the equilibrium equation is HC2H3O2 (aq) + H2O (l)

H3O+ (aq) + C2H3O2- (aq)

The equilibrium expression for an aqueous acid dissociation reaction is called the acid dissociation constant, Ka. ฀฀฀ ฀ =

[฀฀3 ฀฀+ ][฀฀− ] for [฀฀฀฀]

฀฀฀ ฀ =

[฀฀3 ฀฀+ ][฀฀2 ฀฀3 ฀฀−2] [฀฀฀฀2 ฀฀3 ฀฀2 ]

the general equilibrium reaction for the acetic acid equilibrium

Ka is a constant for an acid at a given temperature. Acids with large Ka values dissociate to a large extent in solution and are considered strong acids. Weak acids, on the other hand, have small Ka values and dissociate by a small amount remaining predominantly in the protonated state. The degree to which they dissociate can be measured by pH, which is the negative log of the hydronium ion concentration. pH = -log[H3O+] Kas can be determined a variety ways experimentally; one of which, the method of titration, will be used in this laboratory exercise. Acid-Base Titrations In a titration of an acid with a base, one starts out with an acid or base in a flask and slowly adds the other component (base or acid) until an equivalence point is reached. In an acid base titration, the equivalence point is the point at which the moles of acid present is stoichiometrically equivalent to the moles of base added. If one knows the concentration of either the acid or the base, determining the concentration of the other is a straight forward exercise. In titrations of acids and bases, the equivalence point of the titration must be accurately determined either with an indicator or with a pH meter. While an indicator is useful for a quick and easy equivalence point determination, a pH meter provides significantly more information over the course of the titration. A typical titration curve of a weak acid titrated with a strong base is presented below.

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14 12

pH of solution

10 8 6 4 2 0 0

5

10 15 Volume of NaOH / mL

20

25

The star at the inflection point represents the equivalence point of the titration. As the titration passes through the equivalence point there is a rapid change in the pH of the solution, and if the pH is carefully measured and the titrant added slowly enough, an accurate equivalence point can be readily determined. Molecular Mass determination: If the acid present in a titration is monoprotic (contains one acidic hydrogen), at the equivalence point the number of moles of acid present will be equivalent to the number of moles of sodium hydroxide added. molesNaOH = molesHA for monoprotic acids From this, the amount of acid (number of moles) in an unknown solution can be determined. If the mass of the acid is known, then dividing the mass by the number of moles will give the molar mass of the acid. ฀฀฀฀฀฀฀฀ ฀฀฀฀฀฀฀฀฀฀ ฀฀฀฀฀฀฀฀ = ฀฀฀฀฀฀฀฀฀฀

Ka determination: In the equation: ฀฀฀ ฀ =

[฀฀3 ฀฀+ ][฀฀− ] [฀฀฀฀]

If a situation can be generated where the [A-] = [HA], then Ka = [H3O+]. This is accomplished when one half of the amount of NaOH that it takes to reach the equivalent point is added to the acid solution. The [H3O+] can then be measured by measuring the pH of the solution. In the titration curve below, the happy face represents when one half of the equivalent amount of sodium hydroxide had been added to the acid solution. Determining the pH at this point will allow us to know the hydronium ion concentration and thus the Ka. The following equation shows the relationships ฀฀฀ ฀ =

[฀฀3 ฀฀+ ][฀฀− ] = 10−฀฀฀฀ [฀฀฀฀]

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14 12

pH of solution

10 8 6 4 2

0 0

5

10 15 Volume of NaOH / mL

20

25

Standardization of NaOH Sodium hydroxide is a hygroscopic solid, meaning it will absorb moisture from the atmosphere. For this reason, it is difficult to accurately produce a solution of sodium hydroxide, as the solid sodium hydroxide will often have an unknown amount of water contained in it. Even if you dry the solid sodium hydroxide, in the time it takes to measure it on a balance, it can absorb a significant amount of water. It is, therefore, necessary to test a sodium hydroxide solution to determine its concentration. This method is called standardization. Sodium hydroxide will be standardized with a solid acid, potassium hydrogen phthalate (KHP). Solid acids are typically used in standardizations as they can easily be dried in an oven and their masses can be very accurately determined using a balance.

Potassium hydrogen phthalate (KHP) - KC8H5O4 In the standardization, a known mass of potassium hydrogen phthalate is dissolved in water. An indicator is added to the solution to show when the solution reaches the equivalence point. Sodium hydroxide is then added slowly from a burette until the color of the solution changes indicating that the sodium hydroxide has reacted with all of the potassium hydrogen phthalate. Since they react in a 1:1 molar ratio, converting the mass of potassium hydrogen phthalate to moles will give Page 3 of 7

the number of moles of sodium hydroxide present in the volume of solution added from the burette to reach the equivalence point. Dividing the moles of sodium hydroxide by the volume of sodium hydroxide solution released from the burette, will give the concentration of the solution. Moles of KHP analyte = Moles NaOH titrant (Moles NaOH titrant) / (Volume NaOH titrant) = Molarity of NaOH solution

Procedure

Standardization of NaOH 1. In a large beaker, prepare approximately 500 mL of a 0.1M NaOH solution. You will standardize this solution, so do not worry about accuracy. Ensure that you mix your solution well. 2. Prepare a buret with your solution 3. Accurately measure approximately 0.7g of potassium hydrogen phthalate (KHP), dissolve it in 50 mL of water in an Erlenmeyer flask and add a few drops of phenolphthalein. 4. Titrate the KHP solution with the sodium hydroxide. 5. Repeat to ensure accuracy.

Unknown analysis 1. Prepare 3 - 250 mL beakers with your unknown acid: Accurately measure ~0.2 g of the unknown acid, dissolve it in 50 mL of water in a beaker and add a few drops of phenolphthalein. 2. In the first beaker, run a fast titration to determine the approximate amount of NaOH required to reach the equivalence point (when the solution turns pink). 3. Fort eh second and third beaker, do the following: 4. Plug the pH meter into the labquest and turn it on. 5. Check the calibration of the pH meter in the buffer solutions. If the pH is off from the listed buffer pH by more than 0.10, perform a calibration of the pH meter. TA will have the instructions. If the pH reading is good, proceed to the titration. Do not put the pH meter back into the storage solution until the end of the experiment. You can place it in deionized water in between trials and rinse it with the deionized water wash bottle. 6. Place the beaker on top of a magnetic stirrer, add a stir bar, and begin stirring slowly. Place the pH meter in the solution. Record the pH of the solution once the reading has stabilized. 7. Begin adding the NaOH solution 1 mL at a time and record the pH after each mL has been added up to within 1 mL of the equivalence point (as determined in the fast titration with the first beaker). 8. Within 1 mL of the anticipated equivalence point, record the pH after the addition of every 0.1 mL of sodium hydroxide until you have added 1 mL past the point when the solution turns pink. Mark your data to indicate the point (pH and volume) at which it turns and stays pink. 9. Continue to add the NaOH in 1 mL increments recording the pH for an additional 5 mL. 10. After the final titration, rinse the tip of the pH meter with deionized water and return it to its storage bottle. 11. All waste can be poured down the drain with water. Page 4 of 7

Data Record all data with appropriate labels and units. The following data tables need to be transcribed into your lab notebook for data collection during lab. Unknown Acid (enter letter) _____________

Part 1: NaOH Solution: Sample Mass of NaOH added to 500 mL of DI water Standardization of NaOH: Trial #1

#2

Sample Mass of KHP added to 50.0 mL DI water Initial burette volume of NaOH Final burette volume of NaOH Mass of KHP added to 50.0 mL DI water Initial burette volume of NaOH Final burette volume of NaOH

Titration of Unknown acid (sample preparation): Trial #1 #2 #3

Sample Mass of Unknown Mass of Unknown Mass of Unknown

Part 2: Trial #1: Sample Initial burette volume of NaOH Final burette volume of NaOH Trial #2 and #3: Volume of NaOH (mL)

pH

Volume of NaOH (mL)

pH

Volume of NaOH (mL)

pH

Analysis 1. Determine the concentration of sodium hydroxide using the data from the standardization experiment. 2. For each trial with the unknown acid, calculate the following Trial 2: a. Moles of NaOH required to reach the equivalence point b. Moles of unknown acid in the beaker c. Molar mass of unknown acid Trial 3: a. Moles of NaOH required to reach the equivalence point b. Moles of unknown acid in the beaker c. Molar mass of unknown acid 3.

From the data for each trial, determine the pKa and calculate the Ka of the unknown acid.

Trial 2 Trial 3 4.

Determine the average molar mass and Ka of the acid.

5.

Compare your results to the unknown acid choices listed in “Possible Unknown Acids for Exp. 8” on Blackboard. a. What is the identity of the unknown acid? b. Determine the % error in the experimentally determined molar mass and Ka.

6.

Using excel, make titration curves for each of your two unknown trials. The titration curve will be a plot of pH on the Y axis and volume of sodium hydroxide added on the X axis. These should be scatter plots and can be connected with a “smooth line”. As always, ensure that the scales are appropriate and that the axes are labeled. On the plot, label the equivalence point and thepoint at which the pKa is equal to pH. Attach the plots to the end of this report....