In this chapter, \(\mathrm{HCl}, \mathrm{HBr},\) and \(\mathrm{HI}\) are all listed as strong acids because they are assumed to be ionized completely in water. If, however, we choose a solvent such as acetic acid that is a weaker Bronsted base than water, it is possible to rank the acids in increasing strength as \(\mathrm{HCl}<\mathrm{HBr}<\mathrm{HI}\). (a) Write equations showing proton transfer between the acids and \(\mathrm{CH}_{3} \mathrm{COOH}\). Describe how you would compare the strength of the acids in this solvent experimentally. (b) Draw a Lewis structure of the conjugate \(\operatorname{acid} \mathrm{CH}_{3} \mathrm{COOH}_{2}^{+}\).

In the realm of chemistry, the Bronsted-Lowry theory offers a straightforward way of identifying acids and bases. Simply put, a Bronsted-Lowry acid is a substance that can donate a proton (which is essentially a hydrogen ion, H+), while a Bronsted-Lowry base is one that can accept a proton.
When an acid donates a proton to a base, the acid becomes a conjugate base, and conversely, the base becomes a conjugate acid. This concept is fundamental to understanding reactions such as the one presented in the exercise, where HCl, HBr, and HI are treated as Bronsted-Lowry acids that donate protons to acetic acid, a Bronsted-Lowry base.
By observing these proton transfer reactions, we can determine the strength of an acid based on its tendency to donate a proton. Stronger acids will readily give up their protons, whereas weaker acids will not.

Proton transfer reactions are at the heart of Bronsted-Lowry acid-base chemistry. These reactions are characterized by the movement of protons from acids to bases. For example, when acetic acid (CH3COOH) acts as a base in a proton transfer reaction with strong acids like HCl, HBr, or HI, it gains a proton becoming CH3COOH2+.
This reaction is pivotal in determining the relative strength of these strong acids in a given solvent. In the type of solvent mentioned in the exercise, acetic acid, the extent to which each acid transfers its proton can be measured; hence, indicating the varying strengths of the acids in this particular environment. The stronger the acid, the more complete the transfer of the proton will be.

In chemical reactions, the equilibrium concentration of substances is used to describe the final concentrations of reactants and products when the reaction has balanced out and gets to a point where it no longer changes with time. This is known as dynamic equilibrium.
Understanding equilibrium concentrations is crucial when comparing acid strengths, as shown in the exercise. When acids such as HCl, HBr, and HI are dissolved in a solvent like acetic acid, they reach an equilibrium state where both the acids and their conjugate anions (Cl-, Br-, I-) coexist. By measuring the concentration of these anions at equilibrium, we can infer the strength of the original acids; higher anion concentration indicates a stronger acid, which is because a strong acid dissociates more fully, leaving more of its anion in the solution.

Lewis structures are a visual representation of molecules showing all the atoms, bonds, and lone pairs of electrons. When drawing Lewis structures, certain rules must be followed, such as determining the total count of valence electrons to be used and arranging atoms with lower electronegativity in the center.
Drawing Lewis structures helps chemists visualize the changes that occur during chemical reactions. As part of the exercise, the Lewis structure of the conjugate acid CH3COOH2+ is described, which illustrates the molecule after acetic acid has accepted a proton. This visualization is essential not only to comprehend the structure of the molecule but also provides insight into its reactivity, polarity, and the capacity to form further chemical bonds.