EXPERIMENT 14: DETERMINATION OF THE EQUILIBRIUM CONSTANT OF A WEAK ACID

The following preparatory questions should be answered before coming to class. They are intended to introduce you to several ideas important to aspects of the experiment. You must turn-in your work to your instructor before you will be allowed to begin the experiment. Be sure to bring a calculator and paper to laboratory.

1. Potassium acid phthalate, KHP (KHC8H4O4), is a primary standard reagent used to determine exactly the concentration of a solution of base, such as NaOH, whose concentration is known approximately. (a) What is a primary standard reagent ?

(b) Knowing that KHP is a weak acid, write the neutralization reaction between KHP and NaOH.

2. When 0.317 grams of KHP were dissolved in 50.0 mL of water and titrated with a solution of NaOH, 12.5 mL were required to reach the endpoint of the titration.

(a) How many moles of the acid, KHP, were originally present before the addition of the base?

(b) How many moles of NaOH were required to neutralize the amount of KHP in Part (a)?

(c) What is the concentration of the NaOH solution?

(d) What is the pH of the solution at the endpoint? (Ka{KHP} = 3.89 x 10-6)

(e) Given the following table of indicators, pH ranges and colors, which would be the best indicator for detecting the endpoint of the titration.

3. A weak acid with the general formula of HA will react with a base, such as NaOH. Write the neutralization equation which describes the reaction.

4. If Ka for the weak acid is 1.6 x 10-6, calculate the magnitude of the equilibrium constant for the equation in Problem #4. (Note: First write the net ionic equation for the equation in Problem #4.)

5. The student is now given a 60.0 mL sample of the weak acid in Problem #4. The known concentration of the weak acid is 0.116 M. After preparing a buret the student delivers 20.0 mL of the unknown acid into a 100 mL beaker. This solution is titrated using a solution of NaOH which has a concentration of 0.126 M.

(a) How many moles of the weak acid were added to the beaker?

(b) How many mole of NaOH are required to neutralize (reach the equivalence point) the sample of weak acid?

(c) How many milliliters of the NaOH are required to neutralize the sample of weak acid?

(d) How many moles of NaOH have been added at one half of the volume in part 'c' (volume at the half equivalence point)?

(e) How many moles of the weak acid have reacted at the half equivalence point?

(f) Calculate the pH of the solution at the half equivalence point. (See your text for additional information. Most standard chemistry texts include a discussion of pH curves and pH at half equivalence point.)

(g) Explain how the pH at the half equivalence point is related to Ka for the weak acid.

pHep Theory and Operation

In this experiment you may be using a pHep (pH electronic paper). This small hand held device is an inexpensive instrument commonly used in the chemistry laboratory to accurately measure the pH of a solution. While a complete understanding of how a pH meter functions and the chemistry which takes place during a pH measurement are beyond the scope of this course, a brief discussion regarding the operation of the pHep is in order.

The pHep consists of a liquid crystal display, calibrated in pH units, and a combination electrode. The combination electrode is immersed in a solution and the pH of the solution is displayed in the liquid crystal window. The combination electrode consists of a reference electrode and a sensing electrode. The reference electrode has a known potential which is constant and independent of the solution in which it is immersed. The sensing electrode is sensitive to the [H3O+] in the solution. The potential of this electrode depends upon the concentration of H3O+ in the solution. When the combination electrode is immersed into a solution, a potential difference develops between the reference and sensing electrodes. The potential difference is directly dependent on the [H3O+] in the solution. This potential difference is measured by the pHep and displayed by the meter as a pH value.

To be accurate, the pHep must be calibrated with a solution of known pH. This is accomplished using a carefully prepared buffer solution of known pH. Such solutions can be prepared in the laboratory or purchased from a chemical supply company. When the combination electrode is immersed in the buffer solution, the meter will display a measured pH. If the measured pH is not equal to the known pH of the buffer, then the displayed reading must be adjusted by turning the calibration screw. When the pHep reading equals the pH of the buffer and the reading has become steady, the pHep is calibrated and ready to measure the pH of unknown solutions.

Conditioning and calibrating the pHep

1. Conditioning the pHep: Remove the cap and soak the electrode for at least 30 minutes in either a pH 4.0 or 7.0 buffer solution. Immerse the electrode end in the buffer solution to a depth of one and a half inches.

2. Calibrating the pHep: This can be done by using a buffer solution having a pH close to that of the solution you are measuring. For a pH titration, use a pH 7.0 buffer. Immerse the pHep in the buffer to a depth of one and a half inches and stir gently for a few seconds. Allow the reading to stabilize. If the display does not agree with the known pH of buffer solution turn the screw located on the back of the pHep until the digital display shows the correct pH value. Since the electrode response changes with time, you should periodically recalibrate the pHep.

3. To measure the pH of a solution, immerse the pHep to a depth of one and a half inches above the electrode. Stir gently for a few seconds before recording the reading.

4. After all the readings have been recorded, rinse the pHep with deionized water to remove residue from the electrode. Wet the base of the cap with the buffer solution used to condition the pHep and then close the pHep cap firmly. This will help to prolong the life of the pHep electrode.

EXPERIMENT 14: DETERMINATION OF THE EQUILIBRIUM CONSTANT OF A WEAK ACID

EQUIPMENT:

PART 1: Standardization of the NaOH Solution with KHP

Weigh out approximately 0.300 g of potassium acid phthalate (KHP) using a laboratory balance. Weigh the sample using an aluminum weighing pan, watch glass or weighing paper and carefully transfer the sample to a 125 mL Erlenmeyer flask. Dissolve the solid in 75 - 100 mL of deionized water. Make sure that all of the KHP has desolved before continuing. Add 4 or 5 drops of phenolphthalein indicator solution.

Obs. #1

Obtain 100 mL of NaOH solution. Rinse a buret with deionized water and then with a 5 mL sample of the NaOH solution. Be sure that both the deionized water and the NaOH solution run through the buret tip. Discard both of these solutions. Attach a buret clamp to a ring stand. Insert the cleaned buret into the buret clamp and fill the buret with NaOH solution. Open the stopcock to allow a few mL of the solution to drain from the buret . As it is draining, check the buret tip to be sure no bubbles of air are trapped in the tip. Refill the buret with the NaOH solution to a volume just between 0.00 mL and 1.00 mL. Read the initial volume of the buret to ±0.01 mL and record the initial volume of NaOH in Obs. #2.

Obs. #2

Place the flask containing the dissolved KHP under the buret and adjust the height of the clamp so the buret tip is below the lip of the flask. Place a sheet of white paper under the container to make the color change of the indicator more noticeable. Begin adding the NaOH solution from the buret slowly, swirling the solution in the beaker or flask to insure proper mixing. The point at which the NaOH solution contacts the KHP solution will show a pink color which will disappear upon swirling. As you approach endpoint, greater portions of the solution will be pink. Titrate drop-by-drop so as not to miss the endpoint. Rinse the sides of the flask with small amounts of deionized water to insure that no drops of NaOH are left clinging to the walls of the container. Remember the endpoint of the titration occurs when a pale pink color persists throughout the solution for 45 - 60 seconds. Record the volume of NaOH required to reach the endpoint in Obs. #2.

Calculate the concentration of the NaOH from your data. Show your work in the space provided. Repeat the titration with two more samples of KHP. Before doing the each of the next two titrations, calculate the volume of the NaOH solution required to neutralize the mass of KHP. This way you will know approximately what volume of base will need to reach the endpoint before beginning the titration.

Calc. #1

Average the three measured NaOH concentrations. Use this average value as the concentration of NaOH in the unknown acid titration calculations.

Calc. #2

EXPERIMENT 14: DETERMINATION OF THE EQUILIBRIUM CONSTANT OF A WEAK ACID

EQUIPMENT:

PART II: pH Titration of an Unknown Acid

Obtain 60 mL of the unknown acid solution assigned to you by your instructor. This volume should be enough for three 15 mL samples and a buret rinse. Rinse the buret with deionized water and then with a 5 mL sample of your unknown acid solution. Be sure that both the deionized water and the unknown acid solution run through the buret tip. Discard both of these solutions. Insert the cleaned buret into the other side of the buret clamp and fill the buret with the unknown acid solution. Open the stopcock to allow a few mL of the solution to drain from the buret . As it is draining, check the buret tip to be sure no bubbles of air are trapped in the tip. Refill the buret with the unknown acid solution to a volume just between 0.00 mL and 1.00 mL. Read the initial volume of the buret to ±0.01 mL and record the initial volume in the table below. Add about 15 mL of the acid to a 250 mL beaker. Read the volume of the buret after the addition to ±0.01 mL and record the result in the table. Subtract the initial buret reading from the final buret reading to find the volume of unknown acid in the beaker. (Remember, this number should have 4 significant digits!)

{CONTINUED)

After collecting the data, plot a titration curve for each sample. Use a full sheet of graph paper for each curve. Place pH on the y-axis, and mL of base added on the x-axis. Draw a smooth curve through the data points. Your curves should be almost "S"- shaped and resemble pH curves found in your chemisty text.

Locate the area of the curve where the pH is changing most rapidly with small changes in the amount of base added. This area of the curve should be almost vertical. The equivalence point is the point in the exact center of the vertical section of the curve. Mark the equivalence point on each of your titration curves. Complete the Obs. #5.

Calculate the concentration of each of the acid samples and average the results. Show your work!

Calc. #3

Complete the table below. (Recall that the pH pf the solution will equal the pKa of the weak acid at a point half way to the equivalence point.)

Turn these problems in with the laboratory write-up. SHOW YOUR WORK.

1) The following data was collected when a 25.0 mL sample of an unknown acid was titrated with a 0.100 M NaOH solution.

Determine Ka for the unknown acid.

2. Sketch the following pH curves in the space below. (Put all three curves on the same graph, but use different colored ink (or pencil) to distinguish the solutions.)

(a) a 50.0 mL sample of a 0.200 M solution of strong acid is titrated with a 0.200 M solution of a strong base

(b) a 50.0 mL sample of a 0.0200 M solution of strong acid is titrated with a 0.0200 M solution of a strong base

(c) a 50.0 mL sample of a 0.00200 M solution of strong acid is titrated with a 0.00200 M solution of a strong base.