Lab #14B: Acid-Base Titration Curves

 

Purpose

 

In this experiment you will use a volumetric titration to determine the titration curve for a weak acid. You will determine the equivalence point for the acid and calculate the molar concentration of the acid solution.

 

 

 

Lab #15

8

 

 

Materials

Chemicals

· Ammonia solution- NH3 (1M) ( SDS)

· Vinegar

 

Equipment from Home

· Distilled water

· Cell phone camera

· Clear tape to secure syringe

· Microsoft Excel or another spreadsheet program

 

 

 

Equipment from Kit

· Goggles & gloves

· 50 mL & 10 mL graduated cylinders

· 250 mL & 150 mL beakers

· pH meter & buffer solutions

· syringe and stopcock

· funnel

· pipettes

 

 

Introduction

 

Acid-Base titration curves are graphs that show the successive pH values that occur during the titration of a base with an acid or an acid with a base. The general purpose of a titration is to determine the amount of a particular substance in a sample. An acid-base titration curve can be used to find the equivalence point, which is the point where equal stoichiometric moles of been added. The volume of base added (titrant) at the equivalence point can be used to find the moles of acid in an unknown solution.

 

An Acid-Base titration curve will vary depending on the strength of the acids and bases. When a strong acid is titrated with a strong base, the curve will rise sharply and have an equivalence point at pH = 7. If a weak acid is titrated with a strong base or a weak base is titrated with a strong acid the titration curve is unique because one of the ions in the salt formed will often hydrolyze in water. For a weak acid, the equivalence point will be at a pH higher than 7. For a weak base it will be lower than 7.

 

Volumetric Titration

 

A titration is a process used to determine the volume of a standard solution that is needed to react with a given amount of another substance. In this experiment, your goal is to determine the molar concentration of an acid solution by conducting titrations with a base of known concentration. You will be testing a weak acid found in vinegar, acetic acid (HC2H3O2). You will use ammonia (NH3) solution as your base of known concentration. The reaction equation is shown below in net ionic form.

HC2H3O2(aq) + NH3 (aq) ↔ NH4+ (aq) + C2H3O2– (aq)

 

Notice that a double arrow is used since this is a reversible reaction, or equilibrium, since it is a weak acid. A weak acid is one that doesn’t completely dissociate in solution.

In this experiment, you will use a pH meter to monitor pH as you titrate. The region of most rapid pH change will then be used to determine the equivalence point. The volume of NH3 titrant used at the equivalence point will be used to determine the molarity of the acid solution.

 

Calculating acid concentration

 

Molarity is defined as the moles of solute divided by the liters of solution. Therefore, if you know the molarity of NH3 and the volume used, you can determine the total moles.

 

First, convert the volume of NH3 added to liters (L). Next calculate the moles of NH3 using the volume and the molarity of the NH3, found on the bottle.

 

moles NH3 at equiv. = vol. (in L) NH3 at equiv. x [molarity of NH3]

 

OR moles NH3 = V x M

 

Next, determine the moles of acid in the vinegar at equivalence by using the balanced chemical reaction and the mole ratio:

 

1 mole HC2H3O2 = 1 mole NH3

 

Finally, calculate the molarity (M) of the acetic acid in the vinegar sample

 

Molarity of =

 

 

 

Procedure

 

1. Wear goggles and gloves for this experiment.

2. Obtain the buret (looks like a plastic syringe) from your kit. Remove the plunger from the syringe; it is not needed here.

3. Screw the stopcock onto the threaded tip of the syringe. You can rotate the valve on the stopcock to a position that is either closed (perpendicular to the syringe), open (parallel to the syringe), or anywhere in between.

4. Rinse the syringe with distilled water. Check that the stopcock at the bottom is working properly and is not leaking.

5. Using a very long (> 6 inches) piece of clear tape, secure the syringe to the wall or another sturdy support. Your set up should look very similar to the one shown in the picture below.

 

 

 

 

 

 

Stopcock

 

 

 

 

 

Figure 1. Titrations at home.

 

6. Use the funnel to fill the syringe with NH3 solution. Put an empty beaker under the stopcock and open it briefly to allow NH3 to fill the tip. Add more NH3 if needed, but don’t fill beyond the 35 mL mark. Record the Molarity of the NH3 in Table 1. (It is written on the bottle)

CAUTION: NH3 has a strong odor, so make sure you don’t inhale it directly. Use only in a well-ventilated area.

 

7. Use the 10-mL graduated cylinder to transfer 10.0 mL of vinegar into the 250 mL beaker. Add about 20 mL of distilled water with a graduated cylinder. Record the precise volume of vinegar used in Table 1 on p. 6.

8. Prepare the pH meter for measurements. Fill the 150 mL beaker with about 75 mL of distilled water. Remove the protective cap and put the probe into the distilled water. Turn the meter on by pressing the “ON/OFF” key.

 

 

 

 

9. Calibrate the pH meter. This should only need to be done once. Take the meter out of the distilled water, shake it off and dab dry with a paper towel.

 

a. Place the probe into the buffer pH = 6.86 and hold the “CAL” button for 5 seconds. When you release the button (keep probe in solution), the value “6.86” flashes three times. This point is now done. Rinse the probe in the distilled water, then shake off and dab dry.

b. Place the probe into the buffer pH = 4.00 and hold the “CAL” button for 5 seconds. When you release the button (keep probe in solution), the value “4.00” flashes three times. This point is now done. Rinse the probe in the distilled water, then shake off and dab dry.

c. Place the probe into the buffer pH = 9.18 and hold the “CAL” button for 5 seconds. When you release the button (keep probe in solution), the value “9.18” flashes three times. This point is now done. Place the probe back in the distilled water. You are now ready to measure.

 

10. Place the beaker with vinegar under the syringe. Place the pH meter in the solution. Record the starting pH and the initial volume of NH3 in the syringe to the nearest 0.1 mL on Table 2.

11. Take a picture  of your syringe set-up and include it with your lab report.

12. Add approximately 1 mL of NH3 to the vinegar solution. When the pH stabilizes, record the pH and the reading on the buret.

13. Continue adding NH3 solution in 1 mL increments and enter the buret reading after each increment. When a pH value of approximately 4.5 is reached (or you start to see pH increase more quickly), change to a 0.5 mL increment. Record the pH and buret reading after each addition of NH3.

14. You will eventually observe an abrupt increase in pH. After a pH value of approximately 9 is reached, add about 4 more 1 mL increments, recording the pH and volume in the buret after each addition.

15. Now that you have finished measuring the titration curve, you’ll need to neutralize the sample before flushing it down the drain. Add a couple milliliters of vinegar until the pH is between 6 and 8, and then you can dispose of it in the sink. Flush with lots of water.

16. Rinse the pH Sensor with distilled water in preparation for the second titration. Rinse the 250 mL beaker between samples.

17. Repeat Steps #6 – #13 for a second 10.0 mL sample of vinegar.

 

18. Construct a graph of the titration curve from each of your trials using Excel or another graphing program. Include these graphs in your final report.

d. Open a ‘blank workbook’ in Excel

e. Enter data in the spreadsheet, putting data for the x-axis (volume of NH3 added in mL) in Column A and data for the y-axis (pH) in Column B.

f. Select (i.e. highlight) all the data points in both columns

g. Click on the ‘Insert’ tab, find ‘Charts’ toward the middle, hover over ‘Insert Scatter (X, Y) and choose ‘Scatter with Smooth Lines & Markers’.

h. If you click on the Chart a “+” button will appear. When you click on the “+” it will give you the option to add axis titles and minor gridlines.

i. Label the axes and the title of the Chart.

 

19. Analyze your graph to find the midpoint of the steep vertical increase. This is the equivalence point, which is the largest increase in pH upon the addition of a very small amount of NH3 solution. This won’t always be at an exact data point, so it is often useful to overlay a line over this part of the graph. **If you’re having trouble with this step, please contact me!** Record the volume of NH3 added in mL and the pH at the equivalence point in Table 3.

 

 

 

Name __________________

 

Pre-laboratory Assignment

 

 

1. Provide definitions for the following terms:

a. Acid-Base titration curve

 

 

 

b. Equivalence point

 

 

 

c. Weak acid

 

 

 

2. What is the name and formula for the acid titrated in this experiment? What is the Ka value for the weak acid from Appendix H in your textbook?

 

 

 

 

 

3. Use the calculation instructions in the introduction to solve the following titration example: A 5.0 mL sample of aqueous HNO3 requires 18.5 mL of 0.11 M NaOH to reach the equivalence point. What is the molar concentration of HNO3? The equation for the reaction is:

 

 

 

 

 

 

 

4. What safety rules must be observed during this experiment?

 

Results

 

Table 1

HC2H3O2 Trial	Volume HC2H3O2 (mL)	[NH3] (M) (on bottle)
1
2

 

Table 2

Trial 1	Trial 2
pH	Reading on buret (mL)	Total Volume NH3 added	pH	Reading on buret (mL)	Total Volume NH3 added
		= (initial buret reading) – (current buret reading)

Table 3

 

HC2H3O2 Trial	Volume of NH3 at equivalence point (mL)	pH at equivalence point
1
2

 

 

Calculating Acid Concentration (see Introduction)

Trial 1		Trial 2
Moles of NH3 used at equivalence point Moles = M x V		Moles of NH3 used at equivalence point Moles = M x V
Moles of HC2H3O2		Moles of HC2H3O2
Molarity of HC2H3O2 M = moles/L of vinegar		Molarity of HC2H3O2 M = moles/L of vinegar

 

Average Molarity: _________

 

* SHOW ALL CALCULATIONS

 

 

Questions

 

1. The vinegar bottle normally has a 5% concentration, which is the equivalent of 0.83 M. Find your percent error using the average experimental value that you calculated for vinegar. Comment on your results. (Less than 10% error is acceptable, more than that requires additional discussion.)

 

 

 

 

 

 

 

 

 

 

 

2. Using the Ka value for HC2H3O2 from the pre-lab, find the Kb value for this anion.

 

 

 

 

3. If you reverse the equilibrium reaction described in the pre-lab, you can set up an ICE table and solve for x = [OH-], and use that to determine the pH. Using the concentration you found in this lab as the initial concentration of C2H3O2-, calculate the hydroxide concentration and the pH at the equivalence point.

 

 

 

 

 

 

 

 

 

4. How does this value compare to the pH value on your graph at the equivalence point?