TITRATION CURVES, INDICATORS, AND ACID DISSOCIATION CONSTANTS

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TITRATION CURVES, INDICATORS, AND ACID DISSOCIATION CONSTANTS

Adapted from "Chemistry with Computers"

Vernier Software, Portland OR, 1997

INTRODUCTION

Titration is the volumetric measurement of a solution of known concentration when it reacts completely with a measured volume or mass of another substance. In this experiment, the titrant, which is placed in the buret, will always be a base (NaOH or NH

OH). (Ammonia (NH

) reacts with water to form NH

OH.) The analyte, which is placed in a beaker below the buret, will always be an acid (KHP, HCl, or HC

H

O

).

Titration Curves

The differences in shapes of titration curves when various strengths of acids and bases are combined will be observed. In this experiment you will react the following combinations of strong and weak acids and bases (all solutions are approximately 0.10 M).

• Titration #1: Hydrochloric acid, HCl, with sodium hydroxide, NaOH.

• Titration #2: Acetic acid, HC

H

O

, with sodium hydroxide, NaOH.

• Titration #3: Standardization of NaOH:

Potassium Hydrogen Phthalate, KHP, with sodium hydroxide, NaOH.

• Titration #4: Hydrochloric acid, HCl, with ammonium hydroxide, NH

OH.

• Titration #5: Acetic acid, HC

H

O

, with ammonium hydroxide, NH

OH.

A pH electrode will be placed in one of the acid solutions and a solution of one of the bases will

slowly drip from a buret into the acid solution at a constant rate. A titration curve is normally a

plot of pH versus volume of titrant. In this experiment, however, we will monitor and plot pH

versus time, and assume that time is proportional to volume of base. The volume being delivered

by the buret per unit time should be nearly constant.

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As base is added to acid, a gradual increase in pH will occur until the solution gets close to the equivalence point. Near the equivalence point, a rapid change in pH occurs. At the equivalence point, equal numbers of moles of acid and base have been added and the pH will reflect which species are present. The titration curve will be sigmoidal with the inflection point (the point where the curvature changes direction) is the equivalence point. Beyond the equivalence point, where more base has been added than acid, more gradual increases in pH are observed.

Indicators

Indicators are weak organic acids (HIn) that change color when deprotonated (In

). A few drops of indicator added to the analyte solution before the beginning an acid-base titration. When enough base titrant is added to the analyte solution the equilibrium expressed in equation 1 will shift toward products.

(1) HIn(aq) D In

(aq) + H

(aq)

The result is the formation of more of the deprotonated indicator (In

) and a corresponding color change of the analyte solution (the endpoint). A good indicator for a specific acid-base titration has an endpoint with a pH at or near the pH of the equivalence point. In this experiment,

phenolphthalein indicator will be used for each titration. The pH range of the color change will be observed and compared with the pH of the equivalence point to determine if the indicator is an appropriate choice for each titration. As seen in the Enthalpy and Entropy of Borax

Dissolution experiment, bromcresol green is a good indicator and can be used in place of a pH probe to provide an end point, which is an estimate of the equivalence point.

Standardization of NaOH

Solutions of NaOH can be prepared by either dissolving solid NaOH pellets in water or by

diluting a concentrated solution of NaOH. However, the exact concentration of the solution

prepared by these methods cannot be calculated from the weighed mass or using the dilution

equation for two reasons: 1) Solid sodium hydroxide is hygroscopic ("water-loving"). Pellets of

NaOH exposed to air will increase in mass as they become hydrated so the actual mass of pure

NaOH is not accurately known. 2) Sodium hydroxide in solution reacts with carbonic acid (1)

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and its concentration decreases over time. The acid is formed when small amounts of CO

gas (which is always present in air) dissolves in solution.

(1) H

CO

(aq) + NaOH(aq) D H

O(l) + Na

(aq) + HCO

(aq)

The water used to make the NaOH solution can be boiled to expel the dissolved CO

gas but this time consuming procedure is often not possible in a short laboratory period. A stock solution of NaOH can be made in advance with boiled water but will reabsorb CO

over a period of time unless stored in airtight containers. Therefore, to know the exact concentration of a freshly made NaOH solution or one that has stood in air for some time, the solution needs to be standardized.

That is, the exact concentration must be determined by titrating a known mass of a primary standard acid with the NaOH solution.

A primary standard is a substance used to determine the concentration of a solution. A primary standard should be available in pure form at reasonable cost, have a high equivalent weight to minimize weighing errors, be stable at room temperature, easy to dry, and hygrophobic (should not easily absorb water when exposed to air). And most important, the primary standard should react with the solute of the solution being standardized in a simple straightforward way (i.e. a balanced chemical equation can be written). The primary standard reagent commonly used to standardize NaOH is potassium hydrogen phthalate ("KHP", KHC

H

O

). A monoprotic acid with a molecular weight of 204.22 g/mol, 1 mole of KHP reacts with 1 mole of NaOH.

(2) KHC

H

O

(aq) + NaOH(aq) à H

O(l) + Na

(aq) + K

(aq) + C

H

O

(aq)

To remove any loosely bound waters of hydration, KHP is normally heated at 110°C for one

hour then cooled in a desiccator before use. The exact mass (and number of moles of acid) is

determined by weighing the dried acid on an analytical balance. As with the other titrations

performed in this experiment, the inflection point on the titration curve is taken as the

equivalence point.

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The exact concentration of NaOH is calculated by using the stoichiometry from reaction (5) to convert the number of moles of KHP used to moles of NaOH and then dividing by the volume of NaOH used to reach the equivalence point of the reaction.

Acid Dissociation Constant, K

Acetic acid dissociates in water according to equation (1):

(2) HC

H

O

(aq) + H

O(l) D H

O

(aq) + C

H

O

(aq)

The acid dissociation expression is:

(3) K

= [H

O

] [C

H

O

] / [HC

H

O

]

The acid dissociation constant, K

, is a measure of an acid’s strength. For weak acids these values are less than 1 and typically so small that they are expressed with scientific notation.

Taking the negative log of the K

results in more easily expressed pK

values ranging from 0 to 14 for weak acids.

From the expression for K

we can derive the Henderson-Hasselbach equation for acetic acid:

(4) pH = pK

+ log ([C

H

O

] / [HC

H

O

])

The reaction between a base (such as NaOH) and acetic acid is given below:

(5) HC

H

O

(aq) + OH

(aq) D C

H

O

(aq) + H

O(l)

At the "half-way to equivalence point" exactly one half of the acetic acid has reacted with

hydroxide. At this point, the amount of acetic acid remaining is exactly equal to the amount of

acetate ion, C

H

O

, formed. The ratio [C

H

O

] / [HC

H

O

] in equation (4) is equal to one

and log (1) is equal to zero. So equation (4) becomes:

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(6) pH = pK

(at the half-way to equivalence point)

In today's experiment in Trials #2 and 4 (acetic acid + base) we will pinpoint the time it takes to reach the equivalence point, divide that time by 2 to determine the time it takes to reach the half- way to equivalence point. At the half-way point we will use the Examine function on our titration plot to determine pH and therefore pK

of acetic acid.

SAFETY

Wear safety goggles and lab aprons at all times in lab. Acids and bases are caustic solutions and should be handled with care. If spills occur, wash affected areas immediately and inform your TA or course instructor.

Before starting the experiment, the TA will asks one of you to do a quick demonstration or talk- through one of the following:

1) How to use the pH probe

2) How to read the volume on a buret

PROCEDURE

One concise data table should be created to report the data from all 4 titrations. It is your responsibility to determine the content of that table from the procedure steps below.

1. Work in pairs. Wear safety goggles and lab apron at all time in the lab. CAUTION: The solutions of acid and base are 0.10 M. Even though they are dilute, they are still caustic solutions and can cause painful burns. Wash affected areas thoroughly with cold water.

2. Connect the pH probe to Channel 1 of the LabQuest2. The vertical axis of the graph has pH

scaled from 0 to 14 units. The horizontal axis has time scaled from 0 to 120 seconds. In

Duration, change the collection time from 120 to 600.

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3. Calibrate the pH probe: Tap on the “CH1: pH…” red box, select calibrate in the drop down menu. Click on “Calibrate Now”. Rinse the pH meter with copious amounts of deionized water. Carefully blot dry. Obtain ~10 mL of pH 4, pH 7, and pH 10 buffers in 50 mL beakers. Dip the pH meter in the pH 4 buffer. When the voltage reading stabilizes, enter the pH value of the buffer in the field beneath “Enter Value”. Click Keep. Repeat the process with the pH 10 buffer. Click OK. To check the calibration measure the pH of the pH 7 buffer.

4. Trial #1: Place exactly 100-mL of DI water into a 250-mL beaker. Add 8.0 mL of 0.10 M HCl solution and 3 drops of phenothalein indicator. (Note: All trials (#1-4) will use 8.0-mL of acid in 100-mL of DI water and 3 drops of indicator.)

5. Place the beaker onto a magnetic stirrer and add a small teflon stir bar. Adjust the stirrer to a slow stirring speed. (The magnetic stirrer may cause the computer monitor to flicker.)

6. Use a utility clamp to suspend a pH electrode on a ring stand as shown in Figure 1 below.

Place the pH electrode in the HCl solution and adjust its position toward the side of the beaker so the stir bar does not collide with the fragile electrode. The live Meter window should show a pH value between 2.0 and 3.0 for the HCl solution.

Figure 1

7. Rinse a 50 mL buret with a few mL of the ~0.10 M NaOH solution. Use a funnel and

carefully fill the buret. Place a waste beaker beneath the buret and practice adjusting the drip

rate to 1 and 2 drops per second. (Note: Trials #1, 2, & 3 will use NaOH in the buret and

Trials #4 & 5 will use NH

.)

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8. You are now ready to begin monitoring data. Record the initial volume of titrant in the buret.

Go to the graphing window , and Click . Carefully open the buret stopcock to provide a constant drip rate of about 1 drop per second. Try to line up the NaOH drip so it drops into the spinning vortex of acid solution. (If the drop falls too near the pH probe without mixing with the acid, it will create a "spike" with a high pH value on your graph.)

9. Record the time, pH, and titrant volume of the end point and the equivalence point. Did it change color before, at the same time, or after the rapid change in pH at the equivalence point? Time is displayed in the live Meter window. (Note: A suitable indicator should change color at about the same time as the large jump in pH occurs - near the equivalence point.) When data collection has ended, turn off the buret stopcock to stop the flow of NaOH titrant. Record the final volume of titrant in the buret.

10. Email the data file to your ELN, make sure to add a graph title, axis labels, and change the minimum and maximum values for the x & y axis so that the curve completely fills the data window.

11. Read pH and time values along the pH curve on the LabQuest2 by pressing down on the screen and tracing your finger along graph. Determine and record the approximate time for the equivalence point. Rinse the pH electrode with DI water and store it temporarily in the electrode storage solution (pH 7 buffer) in a beaker. Pour the titrated solution from the 250 mL beaker into a 1 L beaker and set it aside to be neutralized later. Clean and dry the 250 mL beaker for the next trial.

12. Trial #2: Repeat the procedures using NaOH titrant and acetic acid (HC

H

O

) except set

the drip rate of the buret to two drops / second. (It will not be necessary to refill the buret.)

Add one item to Step 10: After determining the number of seconds to reach the equivalence

point, divide that number by two for the number of seconds it took to reach "half-way" to

equivalence point. Determine the pH value at the half-way point. Record this value as the

pK

of acetic acid. Pour the titrated solutions into the 1 L beaker that you set aside.

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13. Trial #3: Repeat the procedure used for Trial 2 for NaOH titrant and potassium hydrogen phthalate (KHP). Weigh ~0.20 g of KHP and record the mass to the nearest milligram.

(KHP is kept in a desiccator to keep it dry, be careful sliding the desiccator lid off and make sure that you put it back on.) In a 250-mL beaker, dissolve the KHP in 100 mL of DI water and add 2-3 drops of phenolphthalein indicator. Insert a note with the mass of KHP and the initial and final volumes for this titration into the LabQuest2. The experiment file for this trial should be emailed to your ELN and uploaded to the EEE DropBox your TA will create.

(An average concentration value will need to be calculated.)

Pour the titrated solutions into the 1 L beaker that you set aside. Drain the remaining NaOH from the buret into the 1 L beaker containing the used titrated solutions.

14. Trial #4: Repeat the procedure using NH

as the titrant and HCl solution. Rinse the buret several times with a few mL of the 0.10 M NH

solution, then fill the buret with 0.10 M NH

.

15. Trial #5: Repeat the procedure using NH

as the titrant and acetic acid (HC

H

O

) except set the drip rate of the buret to two drops / second. The x-axis scaling used in Trial #3 should be sufficient. (It will not be necessary to refill the buret with NH

.) Add one item to Step 10:

After determining the number of seconds it took to reach the equivalence point, find the number of seconds it took to reach half-way to equivalence point. Determine the pH value at the half-way point. Record this value as the pK

of acetic acid. Pour the titrated solution and the remaining NH

from the buret into the 1 L beaker.

16. Rinse and reattach the small plastic bottle containing electrolyte solution to the pH probe.

(The electrolyte solution should cover the tip of the probe, if not ask your TA for a refill.)

17. The large 1 L beaker now contains titrated HCl and acetic acid solutions and excess NaOH and NH

from the buret. Add 1 M HCl, with stirring, until the ammonia smell is gone.

Check the solution with pH paper, if acidic add NaHCO

with stirring until the solution is

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neutral (pH 7). The solution will bubble as CO

gas is released. When the solution is neutralized, pour it down the drain.

18. Your TA will create an EEE Dropbox or Google Drive. Your KHP titration curve should be uploaded for postlab questions.

Make sure to clear your email address and password of the LabQuest2 so others can’t access your email account. Shutdown the LabQuest2 and not simply put it to sleep. To shutdown the LabQuest2: press the home key, select System à Shut Down à OK.

CALCULATION AND DISCUSSION

(1) Label the following on titration curves 1, 2, 4, & 5, if applicable:

- Equivalence point

- pH range of the buffering region - pH = pK

- pH range of phenolphthalein

(2) For which trial(s) was phenolphthalein a good indicator? For which trial(s) was it a poor indicator? For those trials in which it was a poor indicator, suggest a better indicator and indicate its pH range on the titration curve.

(3) Calculate an average concentration of NaOH from the data collected in Trial 3. Use at least four pH curves from the EEE DropBox. Show pertinent data in a table. Perform a Q-test on any of the values appearing to be outliers.

(4) Using the average concentration of NaOH found above, calculate the theoretical pH after

2.50 mL and 9.50 mL of NaOH has been added in both the titration of HCl and of HC

H

O

.

Indicate if the volume of NaOH is before or after the equivalence point. Compare these

theoretical values with the actual values found on the titration curves created in lab. (Hint:

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Use total volume of titrant (the initial and final volume on the buret) in conjunction with the time required to deliver that volume to convert the volumes given above to time.) Show a sample of all necessary calculations. Calculate the percent error from the theoretical values.

(5) Why does the equivalence point occur at different pH values for the four titrations studied?

(Hint: What species are present at the equivalence point for each trial?)

(6) Calculate the K

of the acetic acid from the two titration curves involving acetic acid.

Report the % error for each.

(7) Using the pH curve created in Trial 3, calculate the K

of KHP. Report the % error.

QUALITATIVE ERROR ANALYSIS

1. What modifications could be made to the procedure to better account for random (indeterminate) errors?

2. List three potential systematic (instrumental, methodological, or personal) errors that could be made in this experiment. (Note: Be specific, systematic errors are in the details. For example, losing your solution because you knocked over the cuvette is not a systematic error – it’s a gross one.)

3. Did any gross errors occur? Did you mess up? Did the equipment or instrumentation fail?