Acid Strength and Electron Withdrawing Groups

Introduction

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.

Concepts

• Weak acid
• Conjugate base
• Equilibrium constant
• Inductive effect

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. All the weak acid solutions and waste solutions can be neutralized and rinsed down the drain with excess water according to Flinn Suggested Disposal Method #24a.

Prelab Preparation

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 25-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 25-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

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 1.
2. Using a clean Beral-type pipet, add approximately 8–10 mL of the trichloroacetic acid “half-neutralized” solution to each of the four 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 8–10 mL of the chloroacetic acid “half-neutralized” 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

12676_Student1.pdf

Teacher Tips

• This kit contains enough chemicals to perform the demonstration seven times: 280 mL of 0.1 M acetic acid solution, 280 mL of 0.1 M chloroacetic acid solution, 280 mL of 0.1 M trichloroacetic acid solution, 50 mL of methyl red indicator solution, 50 mL of bromphenol blue indicator solution, 50 mL of orange IV indicator solution, 50 mL of universal indicator, rainbow acid indicator solution, 50 mL of phenolphthalein indicator solution, 100 mL of 1.0 M sodium hydroxide solution, 7 reusable plastic Petri dishes and 50 Beral-type pipets.

• 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:
  {12676_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 allow students to determine pH value for this indicator. Other indicators may also be used. A comprehensive list of acid–base indicators can be found in Section 8–19 of the CRC Handbook of Chemistry and Physics.

Sample Data

Record the color of each solution then refer to the indicator chart to determine the pH range for each of the added indicators.

Data Table

Indicator Chart

Answers to Questions

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 2 and 1 and 1)
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:

Acetic acid: 4.75
Chloroacetic acid: 2.85
Trichloroacetic acid: 0.70

Discussion

The modern Brønsted definition of an acid relies on the ability of the compound to donate hydrogen ions to other substances. 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 1, 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 2. The square brackets refer to the molar concentrations of the reactants and products.

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 1. 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 1 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 3). 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 3 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 2 reduces to Equation 3.

And therefore

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.

References

Special thanks to Lee Marek, retired, Naperville North H.S., Naperville, IL, for providing the idea and the instructions for this activity to Flinn Scientific.