Acid Strength and Electron Withdrawing Groups
Publication No. 12676
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.
(for each demonstration)
Acetic acid solution, CH3COOH, 0.1 M, 40 mL*
Bromphenol blue indicator solution, 6 mL*
Chloroacetic acid, ClCH2COOH, 0.1 M, 40 mL*
Methyl red indicator solution, 6 mL*
Orange IV indicator solution, 6 mL*
Phenolphthalein indicator solution, 1%, 2 mL*
Sodium hydroxide, NaOH, 1.0 M, 10 mL*
Trichloroacetic acid, Cl3CCOOH, 0.1 M, 40 mL*
Universal indicator solution, rainbow acid, 6 mL*
Water, distilled or deionized
Beakers, 100-mL, 3
Graduated cylinder, 10-mL
Graduated cylinder, 25-mL
Light box or overhead projector
Petri dishes, disposable, small, 7*
Pipets, Beral-type, disposable, 7*
*Materials included in kit.
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.
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.
Three acids will be compared: acetic acid, chloroacetic acid and trichloracetic acid. Prepare a half-neutralized buffer solution for each acid as follows:
Student Worksheet PDF
Correlation to Next Generation Science Standards (NGSS)†
Science & Engineering PracticesAsking questions and defining problems
Developing and using models
Planning and carrying out investigations
Analyzing and interpreting data
Using mathematics and computational thinking
Disciplinary Core IdeasHS-PS1.B: Chemical Reactions
Crosscutting ConceptsCause and effect
Systems and system models
Scale, proportion, and quantity
HS-PS1-2: Construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.
Record the color of each solution then refer to the indicator chart to determine the pH range for each of the added indicators.
Answers to Questions
Acetic acid: pH is >4.6 and 2 and 1 and 1)
The pH of a weak acid solution can be expressed as the Henderson-Hasselbach equation:
For weak acids with Ka 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 pKa for the 3 weak acids are:
Acetic acid: 4.75
The pH ranges for the acetic acid and the chloroacetic acid agree with the pKa values. The Ka 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 = pKa + log(>1)
Since the log of any positive number greater than 1 is positive, the pH of the solution is greater than the pKa or the pH is greater than 0.7.
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 (Ka) 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 Ka for Equation 1. Weak acids ionize only partially in aqueous solution. The value of Ka 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 Ka 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.
The half-neutralized pH values for the three weak acids are:
Acetic acid pH = 4.75
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.
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.