Acids and Bases

Introduction

The course description for the College Board AP^® Chemistry lists the following topics for acids and bases: concepts of Arrhenius, Brønsted-Lowry and Lewis, coordination complex, amphoterism, pK_a and pH, buffers and hydrolysis. Use this set of three acid–base neutralization demonstrations to engage your students in a hands-on review and test their understanding of this key topic in chemistry.

The set of three interactive demonstrations includes:

• Weak Base–Weak Acid Titration—Start the review with a titration of a weak acid with a weak base. What is the pH at the end-point? Students graph the pH versus base added. Based on the graph, the students determine the relative strengths of the acid and base.
• Titrating an Acid Mixture—A mixture of hydrochloric and maleic acid is titrated with sodium hydroxide. Students graph the curve and use the data to determine the concentration of each acid and the K_a2 of maleic acid.
• Acid Strength and Electron Withdrawing Groups—Show the trend for K_a of acetic acid and its mono- and trichlo-derivatives. Challenge students to explain the trend based on these values and their knowledge of electron affinity.

The series of demonstrations may be presented in a variety of ways. Each demonstration may be used to review a specific AP test topic, or all the demonstrations can be performed together as a review student understanding and grasp of stoichiometry concepts normally covered in the AP exam. A student worksheet is included as an optional assessment tool for the instructor.

Concepts

• Weak acid–conjugate base
• Weak base–conjugate acid
• Equilibrium constant
• K_a and K_b
• pH
• Inductive effect

Experiment Overview

Activity 1. Weak Acid–Weak Base Titration
If a weak acid is titrated with a strong base, the equivalence point occurs at a pH value greater than 7. If a weak base is titrated with a strong acid, the equivalence point occurs at a pH value less than 7. Where does the equivalence point occur when a weak acid is titrated with a weak base?

Activity 2. Titrating an Acid Mixture
What happens if you mix a weak and strong acid and titrate this mixture with a strong base? Can you determine the concentration of each acid? The K_a of the weak acid? Put your students to the test with this thought-provoking demonstration.

Activity 3. Acid Strength and Electron Withdrawing Groups
Acids vary greatly in their strength—their ability to ionize or produce ions when dissolved in water. What factors determine the strength of an acid? In this demonstration, the effect of electron withdrawing groups on the strength of acids will be shown by preparing half-neutralized solutions of a series of chlorinated acetic acids and comparing their pH values using indicators.

Materials

Safety Precautions

Phenolphthalein is moderately toxic by ingestion. Universal indicator solution is slightly toxic by ingestion. All the carboxylic acid solutions are corrosive liquids and body tissue irritants. The sodium hydroxide solution is corrosive to skin and eyes and is also a body tissue irritant. Avoid contact with eyes and skin and clean up all spills immediately. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please consult current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. Remaining solutions may be neutralized and flushed down the drain with excess water.

Prelab Preparation

Activity 2. Titrating an Acid Mixture

0.1 M Maleic Acid Solution: Add 50 mL of distilled or deionized water to a 100-mL volumetric flask. Mass 1.12 g of maleic acid and add this to the flask. Fill to the 100-mL mark with distilled or deionized water. Cap and mix thoroughly..

Activity 3. Acid Strength and Electron Withdrawing Groups
Three acids will be compared: acetic acid, chloroacetic acid and trichloracetic acid. Prepare a half-neutralized buffer solution for each acid as follows:

1. Using a 50-mL graduated cylinder, precisely measure 20.0 mL of the appropriate 0.1 M acid solution into a clean 100-mL beaker.
2. Add 3 drops of phenolphthalein indicator to the acid solution in the beaker.
3. Using a Beral-type pipet, add 1.0 M sodium hydroxide solution dropwise to the beaker. Gently swirl the beaker while adding the sodium hydroxide.
4. Continue adding the sodium hydroxide dropwise until a faint pink color persists throughout the solution for at least five seconds. Note: A pink color develops immediately when the base is added, but fades quickly when the solution is swirled. When nearing the endpoint, the pink color begins to fade more slowly. Proceed cautiously when nearing the endpoint, so as not to “overshoot” it.
5. Using a 50-mL graduated cylinder, precisely measure 20.0 mL of the acid solution and pour this into the beaker containing the neutralized solution. Note: The pink color of the indicator will disappear.

Procedure

Activity 1. Weak Acid–Weak Base Titration

1. Set up a pH meter and electrode. Calibrate the pH meter using a buffer solution of pH 7.00. Rinse the electrode well with distilled water.
2. Using a graduated cylinder, add 20 mL of the 0.1 M chloroacetic acid solution to a 100-mL beaker.
3. Fill the buret with about 25 mL of the 0.1 M ammonia solution. Have students record the initial volume as the “initial buret reading ” in the Titration Data Table.
4. Set the beaker containing the weak acid solution on a magnetic stirrer. Clamp the pH electrode so it is submerged in the acid solution (see Figure 1). Be sure the stir bar does not hit the electrode. Set the stir bar gently spinning.
   {12161_Procedure_Figure_1_Setup}
5. When the pH reading has stabilized, have students record the initial pH of the solution in the Titration Data Table.
6. Add about 1 mL of ammonia solution to the beaker. Have students record the exact buret reading and the pH of the solution in the Titration Data Table.
7. Add another 1-mL increment of of ammonia solution. Have students record Figure 1 both the buret reading and the pH in the Titration Data Table.
8. Continue adding ammonia solution in 1-mL portions. Record both the buret reading and the pH after each addition.
9. When the pH begins to increase by more than 0.3 pH units after an addition, decrease the portions of ammonia solution added to about 0.2 mL.
10. Continue adding ammonia solution in about 0.2 mL increments. Record both the buret reading and the pH after each addition.
11. When the pH change is again about 0.3 pH units, resume adding the ammonia solution in 1-mL increments. Continue to record both the buret reading and the pH after each addition.
12. Stop the titration when the pH change is less than 0.1 pH units. Record the final volume of solution in the buret and the final pH.
13. Have students graph the data, with pH on the vertical axis and volume NH₃ on the horizontal axis. Identify the pH at the equivalence point.

Activity 2. Titrating an Acid Mixture

1. Set up a pH meter and electrode. Calibrate the pH meter using a buffer solution of pH 7.00. Rinse the electrode well with distilled water.
2. Fill the buret with about 25 mL of the 0.1 M sodium hydroxide solution. Have students record the initital volume as the “initial buret reading ” in the Titration Data Table.
3. Add 4 mL of hydrochloric acid solution to a 100-mL beaker, then 6 mL of maleic acid solution to the same 100-mL beaker.
4. Set the beaker containing the mixed acid solution on a magnetic stirrer. Clamp the pH electrode so it is submerged in the acid solution (see Figure 1). Be sure the stir bar does not hit the electrode. Set the stir bar gently spinning.
5. When the pH reading has stabilized, have students record the initial pH of the solution in the Titration Data Table.
6. Add about 1 mL of hydroxide solution to the beaker. Have students record the exact buret reading and the pH of the solution in the Titration Data Table.
7. Add another 1-mL increment of of hydroxide solution. Have students record Figure 1 both the buret reading and the pH in the Titration Data Table.
8. Continue adding hydroxide solution in 1-mL portions. Record both the buret reading and the pH after each addition.
9. When the pH begins to increase by more than 0.3 pH units after an addition, decrease the portions of hydroxide solution added to about 0.2 mL.
10. Continue adding hydroxide solution in about 0.2 mL increments. Record both the buret reading and the pH after each addition.
11. When the pH begins to again increase by more than 0.3 pH units after an addition, decrease the portions of hydroxide solution added to about 0.2 mL.
12. When the pH change is again about 0.3 pH units, resume adding the hydroxide solution in 1-mL increments. Continue to record both the buret reading and the pH after each addition.
13. Stop the titration when the pH change is less than 0.1 pH units. Record the final volume of solution in the buret and the final pH.
14. Have students graph the data and determine the pH of the neutralization reaction at the equivalence point, the concentration of each acid in the mixture and the K_a2 of the weak acid.

Activity 3. Acid Strength and Electron Withdrawing Groups

1. Assemble the tops and bottoms of the disposable Petri dishes into a 3 x 4 grid pattern on the overhead projector stage. (There will be 12 “wells.”) See Figure 2.
   {12161_Procedure_Figure_2}
2. Using a clean Beral-type pipet, add approximately 8–10 mL of the trichloroacetic acid “half neutralized” solution to each of thefour Petri dishes in the first column. To mimic the data table, make sure this column is on the left as you face the projected image.
3. Using a separate clean Beral-type pipet, add 8–10mL of the chloroacetic acid “half-neutral¬ized” solution to each of the four Petri dishes in the middle column.
4. Repeat step 3 adding acetic acid “half-neutralized” solution to each of the four Petri dishes in the third column.
5. Add 1–2 mL of methyl red indicator solution each Petri dish in the first row. These dishes should contain, in order, trichloroacetic acid, chloroacetic acid, and acetic acid.
6. Swirl to mix. Have the students record the colors of each solution in the worksheet handout. Estimate the pH range for each acid in the “Methyl Red” row.
7. Repeat steps 5 and 6 using the bromphenol blue indicator solution, then the orange IV indicator solution, and, finally, the “rainbow acid” universal indicator solution in rows two through four, respectively.

Student Worksheet PDF

12161_Student1.pdf

Teacher Tips

• In the Weak Acid–Weak Base Titration activity, if a trial run is done first, the increments of buret additions can be increased at the beginning and end to shorten the titration time.
• The use of computer-based technology for data collection and analysis is tailor-made for acid–base titrations. The graph showing pH versus volume titrant can be drawn using a graphing calculator or a graphical analysis program on a computer.
• In Acid Strength and Electron Withdrawing Groups activity, the K_a for trichloroacetic acid is 0.20. When calculating the pH of the half-neutralized 0.1 M solution, the equation for K_a is:
  {12161_Tips_Equation_1}
  Since K_a is greater than 1 x 10^–2, the x terms in the numerator and denominator cannot be neglected. The value of x must be solved quadradically. This yields a pH value of 1.45 at half-neutralization.
• The indicators were chosen to bracket the pH values of the half-neutralized weak acid solutions and give vivid colors. Pass around the color chart for the “rainbow acid” universal indicator to determine pH values for this indicator. Other indicators may also be used.

Sample Data

Activity 3. Acid Strength

Indicator Chart

Answers to Questions

Activity 1. Weak Acid–Weak Base Titration

1. Record the titration data, then graph volume of NaOH added versus pH. 
   {12161_Answers_Figure_3}
2. What is the pH of the solution at the neutralization point?

   Approximately 6.1

3. Based on the pH at the neutralization point and the fact that the concentration of the acid and base solutions are equal, which is larger, K_b for ammonia, the weak base, or K_a for chloroacetic acid, the weak acid?

   Since the pH at the equivalence point is < 7.00, the K_a of the conjugate acid of ammonia is larger than the K_b of the conjugate base of chloroacetic acid. This means that the K_a of chloracetic acid is larger than the K_b of ammonia.
   K_a of chloracetic acid = 1.3 x 10^–3
   K_b of ammonia = 1.8 x 10^–5

Activity 2. Titrating an Acid Mixture
Record the titration data, then graph volume of NaOH added versus pH.

1. The acid solution is a mixture of hydrochloric acid, HCl, and the diprotic weak acid, maleic acid, (CH₂)₂C₂O₄H₂. The concentration of the sodium hydroxide solution is 0.10 molar. Recalling that HCl is a strong acid, that is it completely dissociates in solution and maleic acid is a weak diprotic acid, use the titration curve data and the sodium hydroxide concentration to determine the initial concentration of each acid in the original mixture.

   The first endpoint corresponds to the consumption of both HCl and the first ionizable hydrogen of maleic acid. The second endpoint is solely due to the second ionizable hydrogen of maleic acid. The moles of HA^– titrated to this second endpoint are equal to the initial moles of H₂A, maleic acid, in the original solution.
   0.006 L x 0.10 moles/L = 6.0 x 10^–4 mole maleic acid
   6.0 x 10^–4 mole maleic acid/ 0.010 L solution = 0.06 M maleic acid
   The first endpoint corresponds to the consumption of both HCl and the first ionizable hydrogen of maleic acid. The concentration of the first ionizable hydrogen of maleic acid is the same as the concentration of its second ionizable hydrogen. Since 6 mL of base were required to neutralize the second hydrogen, it must also take 6 mL of base to neutralize maleic acid’s first hydrogen. That leaves 4 mL to neutralize all of the HCl in solution.
   0.004 L x 0.10 moles/L = 4.0 x 10^–4 moles hydrochloric acid
   4.0 x 10^–4 mole hydrochloric acid/0.010 L solution = 0.04 M hydrochloric acid

2. The pK₁ and pK₂ for a diprotic acid H₂A are given by the equations
   {12161_Answers_Equation_3}
   Use the titration curve to determine pK_2.
   {12161_Answers_Figure_5}
   pK₂ = 6.1

Activity 3. Acid Strength

1. Based on your observations, what is the pH range for the half-neutralized acetic acid solution? What is the range for the half-neutralized chloroacetic acid solution? For the half-neutralized trichloroacetic acid solution?

   Acetic acid: pH is >4.6 and <4.8
   Chloroacetic acid: pH is >2 and <2.8
   Trichloroacetic acid: pH is >1 and <1.4

2. For a weak acid (HA), K_a, the dissociation constant, is equal to:
   {12161_Answers_Equation_4}
   The pH of a weak acid solution can be expressed as the Henderson-Hasselbach equation:
   {12161_Answers_Equation_2}
   For weak acids with K_a values of 1 x 10^–2 or less, at half-neutralization the conjugate base concentration, [A^–], is essentially equal to the weak acid concentration, [HA]. Equation 2 becomes

   pH = pKa + log(1) or
   pH = pKa
   The pK_a for the three acids are:

   {12161_Answers_Table_3}
   Do your pH range estimations agree with these values? If not, what are some possible explanations?
   The pH ranges for the acetic acid and the chloroacetic acid agree with the pK_a values. The K_a for trichloracetic acid is greater than 1 x 10^–2. Because of this, the concentration of the conjugate base is larger than the undissociated acid at equilibrium.
   pH = pK_a + log(>1)
   Since the log of any positive number greater than 1 is positive, the pH of the solution is greater than the pK_a or the pH is greater than 0.7.

Discussion

Activity 1. Weak Acid–Weak Base Titration
In acid–base titrations, the plot of pH versus volume of titrant results in an S-shaped curve (see Figure 6)

The steepness of the curve and the pH value at the equivalence point depend on the strength of both the acid and the base. If both the acid and base are strong, the curve is very steep and the equivalence point pH value is 7. If a weak acid is titrated by a strong base, the titration curve is less steep and the equivalence point pH value is >7. At the equivalence point, assuming the acid is monobasic (HA). or The overall neutralization reaction is, At this point, the initial moles of the weak acid (HA) have been completely converted to its conjugate base (A^–). This conjugate base is a weak base and equilibrates with water to form a basic solution. K_b for this reaction is, The story for a weak base is similar. The conjugate acid of a weak base produces an acidic solution at the equivalence point when titrated with a strong acid. What happens at the equivalence point when a weak base is titrated with a weak acid? For a weak acid with a dissociation constant K_a' the value of K_b' for its conjugate base is: For a weak base with a dissociation constant of K_b, the value of K_a for its conjugate acid is: Combining Equations 12 and 13 yields: If K_a' of the weak acid is greater than K_b of the weak base, then K_a is greater than K_b'. The pH of the solution at the equivalence point is acidic. If K_b of the weak base is greater than K_a' of the weak acid, the pH of the solution at the equivalence point is basic.

Activity 2. Titrating an Acid Mixture
Dealing with a mixture of acids seems tricky at first glance. If, for example, you were to titrate a solution of only HCl or diprotic weak acid H₂A, it would be a simple matter to determine the concentration of the acid and/or the dissociation constant, K_a, of the weak acid. However, when you titrate an aqueous solution containing both HCl and H₂A the task of determining much of anything about the mixture seems quite complex—until you think about equilibrium. Let’s explore the weak diprotic acid in more detail. Now, when you mix some HCl with a solution of H₂A, you are adding a huge amount of H₃O⁺ ions because HCl is a strong acid (dissociates completely). Recall what Le Chatelier had to say about equilibria and you will see that the 1st dissociation equation contains a whole lot of H₂A and H₃O⁺ and practically no HA^–. This is what your acid mixture looks like before you start the titration.

As you titrate the mixture of HCl and H₂A with sodium hydroxide, the OH^– ions react with H₃O⁺ ions from the HCl and the H₂A molecules to reach the 1st equivalence point. On the titration plot, however, there is no way to identify the source of the H₃O⁺ ions up to the 1st equivalence point. But, we’ll be able to figure it out.

As the titration continues past the 1st equivalence point, you are now neutralizing the HA^– ions. The number of moles of NaOH you add from the 1st eq. pt. to the 2nd eq. pt. is equal to the number of moles of HA^– ion. Also, the original amount of H₂A in the mixture is equal to the moles of HA^– ions (see the 1st dissociation equation). Thus, to calculate the number of moles of HCl in the mixture you subtract the moles of HA^– from the moles of NaOH used to reach the 1st eq. pt. From here, calculating the molar concentration of HCl and H₂A is simple. That piece of the puzzle is solved.

How do we calculate the K_a2 (or the pK_a2) of the diprotic acid? The equilibrium expression for the 2nd dissociation is shown: If we take the logarithm of both sides of this equation, and do a bit of rearranging, we get the Henderson–Hasselbalch equation: In this version of the Henderson-Hasselbalch equation, if the molar concentration of the two ionic species is equal then the pK_a2 and the pH are equal. This event occurs exactly halfway between the 1st eq. pt. and the 2nd eq. pt. Thus, you can read the pK_a2 of a diprotic acid directly from the titration plot.

Here’s an example. The following graph shows the titration of 10.0 mL of a mixture of HCl and maleic acid with 0.10 M NaOH solution. (0.006 L) x (0.10 mol/L) = 6.0 x 10^–4 mol of OH^– used between 1st and 2nd eq. pts.
∴6.0 x 10^–4 mol of H₂A and 6.0 x 10^–4 mol of HA^– were in the mixture
[H₂A] = 6.0 x 10^–4 mol 0.010 L of total mixture = 0.06 M maleic acid

(0.010 L) x (0.10 mol/L) = 0.0010 mol of OH^– used to reach the 1st eq. pt.
∴0.0010 mol of H₃O⁺ ions came from HCl + H₂A
...and, we know that there were 6.0 x 10^–4 mol of H₂A
∴mol HCl = 0.0010 mol – 6.0 x 10^–4 mol = 4.0 x 10^–4 mol
[HCl] = 4.0 x 10^–4 mol 0.010 L of total mixture = 0.04 M HCl
Examine another region of the titration plot to determine the pK_a2 of maleic acid.

Halfway between the 1st equivalence point and the 2nd equivalence point is the point at which 13.00 mL of NaOH were added. This also corresponds to [A^2–] = [HA^–]. The pH at this point is 6.07, which is the pK_a2 of the diprotic acid, maleic acid.

Activity 3. Acid Strength and Electron Withdrawing Groups
When an acid dissolves in water, it donates hydrogen ions to water molecules to form H₃O⁺ ions. The general form of this reaction, called an ionization reaction, is shown in Equation 16, where HA is the acid and A^– its conjugate base after loss of a hydrogen ion. The double arrows represent a reversible reaction. The equilibrium constant expression (K_a) for the reversible ionization of an acid is given in Equation 17. Not all acids, of course, are created equal. The strength of an acid depends on the value of its equilibrium constant K_a for Equation 17. Weak acids ionize only partially in aqueous solution. The value of K_a for a weak acid is much less than one, so that Equation 17 is reversible—all species (HA, A^– and H₃O⁺) are present at equilibrium. The three weak acids used in this demonstration are acetic acid, chloroacetic acid and trichloroacetic acid. Any factor that stabilizes the conjugate base anion more that the weak acid should increase the strength of the acid and its K_a value. Electron-withdrawing groups, such as chlorine atoms, disperse the negative charge on the conjugate base anion and thus stabilize the anion relative to the acid (see Figure 9). In fact, chloroacetic acid is on the order of 100 times stronger than acetic acid, and trichloroacetic acid is more than 30,000 times stronger! The arrows in Figure 9 are meant to indicate the electron withdrawing nature of the chlorine atoms. The ionization constant of a weak acid can be determined experimentally by measuring the H₃O⁺ concentration in a dilute aqueous solution of the weak acid. This procedure is most accurate when the solution contains equal molar amounts of the weak acid and its conjugate base. If the concentrations of the weak acid [HA] and the conjugate base [A^–] are equal, then these two terms cancel out in the equilibrium constant expression, and Equation 18 reduces to Equation 19. The half-neutralized pH values for the three weak acids are:

Acetic acid pH = 4.75
Chloroacetic acid pH = 2.85
Trichloroacetic acid pH = 1.45*

*See Tips section When methyl red is added to the half-neutralized acid solutions, all produce a red color, indicating all have a pH value less than 4.8. When bromphenol blue is added, the acetic acid solution turns purple and the other two are yellow. This shows that the acetic acid solution pH is greater than 4.6 and the other two have pH values that are less than 3.0. Orange IV is now added to separate the chlorinated weak acids. At pH values of 1.4 or less, solutions are red, at pH values of 2.8 or greater, solutions are yellow, and at values between 1.4 and 2.8, the solutions are shades of orange. The red color of the trichloroacetic acid solution shows its pH is lower than that of the orange chloroacetic acid solution. Finally, “rainbow acid” universal indicator is added to show that the pH of the half-neutralized trichloroacetic acid solution falls between 1 and 1.4, the chloroacetic acid solution pH is between 2 and 2.8, and the acetic acid pH is between 4.6 and 4.8.