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Chapter 16: Problem 66

Explain why trichloroacetic acid, \(\mathrm{CCl}_{3} \mathrm{COOH},\) is a stronger acid than acetic acid, \(\mathrm{CH}_{3} \mathrm{COOH}\).

Short Answer

Expert verified
Trichloroacetic acid (CCl3COOH) is a stronger acid than acetic acid (CH3COOH) due to the effect of the three chlorine atoms. The electronegativity of the chlorine atoms pulls electron density away from the -COOH group, making it easier for the proton (H+) to be released, thus increasing the acidity.

Step by step solution

01

Understand the Concept of Acidity

The acidity of a molecule is determined by how easily it can lose a proton (H+). In this comparison, both acetic acid and trichloroacetic acid can lose a proton from the -COOH group, becoming -COO-.
02

Practicing Inductive Effect on Acetic Acid

In acetic acid, the CH3 group is attached to the -COOH acid group. Carbon (C) and Hydrogen (H) exhibit similar electronegativities, meaning there is no significant \'pull\' or shift of electron density from the -COOH group. Hence, the ability of acetic acid to lose a H+ from the -COOH group is not significantly affected by the CH3 group.
03

Practicing Inductive Effect on Trichloroacetic Acid

In contrast, trichloroacetic acid (CCl3COOH) has three Chlorine (Cl) atoms attached to the -COOH group. Chlorine is highly electronegative and therefore, each of them will pull electron density towards itself. This makes the H+ of the -COOH group more positively charged and therefore, easier to be released as a proton.
04

Comparing the Acidity

The more easily a molecule can lose a H+ proton, the stronger an acid it is. Therefore, trichloroacetic acid is a stronger acid than acetic acid because the CCl3 group in trichloroacetic acid pulls electron density away from the -COOH group, thereby reducing the proton’s electron density and making it easier for the proton to be released.

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Most popular questions from this chapter

Suppose you wanted to produce an aqueous solution of \(\mathrm{pH}=8.65\) by dissolving one of the following salts in water. Which salt would you use, and at what molarity? (a) \(\mathrm{NH}_{4} \mathrm{Cl} ;\) (b) \(\mathrm{KHSO}_{4} ;\) (c) \(\mathrm{KNO}_{2}\); (d) \(\mathrm{NaNO}_{3}\).

Explain why \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) in a strong acid solution doubles as the total acid concentration doubles, whereas in a weak acid solution, \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) increases only by about a factor of \(\sqrt{2}\)

A saturated aqueous solution of \(\mathrm{Ca}(\mathrm{OH})_{2}\) has a pH of 12.35. What is the solubility of \(\mathrm{Ca}(\mathrm{OH})_{2}\), expressed in milligrams per 100 mL of solution?

What is the pH of each of the following solutions? (a) \(0.0045 \mathrm{M} \mathrm{HCl} ;\) (b) \(6.14 \times 10^{-4} \mathrm{M} \mathrm{HNO}_{3} ;\) (c) 0.00683 \(\mathrm{M} \mathrm{NaOH} ;\) (d) \(4.8 \times 10^{-3} \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\)

Explain why \(\left[\mathrm{PO}_{4}^{3-}\right]\) in \(1.00 \mathrm{M} \mathrm{H}_{3} \mathrm{PO}_{4}\) is not simply \(\frac{1}{3}\left[\mathrm{H}_{3} \mathrm{O}^{+}\right],\) but much, much less than \(\frac{1}{3}\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\)
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