Question

Students are often surprised to learn that organic acids, such as acetic acid, contain \(-\) OH groups. Actually, all oxyacids contain hydroxyl groups. Sulfuric acid, usually written as \(\mathrm{H}_{2} \mathrm{SO}_{4}\), has the structural formula \(\mathrm{SO}_{2}(\mathrm{OH})_{2}\), where \(\mathrm{S}\) is the central atom. Identify the acids whose structural formulas are shown below. Why do they behave as acids, while \(\mathrm{NaOH}\) and \(\mathrm{KOH}\) are bases? a. \(\mathrm{SO}(\mathrm{OH})_{2}\) b. \(\mathrm{ClO}_{2}(\mathrm{OH})\) c. \(\mathrm{HPO}(\mathrm{OH})_{2}\)

Short answer

The given acids are: a. \(H_2SO_3\) (sulfurous acid) b. \(HClO_2\) (chlorous acid) c. \(H_3PO_3\) (phosphorous acid) These acids behave as acids because the hydrogen atoms bonded to highly electronegative oxygen atoms can easily be released as protons (H⁺ ions), increasing the hydrogen ion concentration in the solution. In contrast, bases like \(\mathrm{NaOH}\) and \(\mathrm{KOH}\) contain hydroxide (OH⁻) ions that accept protons to form water.

Step-by-step solution

Identify the acids based on their structural formulas

The given structural formulas can be written in chemical formula notation as follows: a. \(\mathrm{SO}(\mathrm{OH})_{2} \Rightarrow \mathrm{H}_2\mathrm{SO}_3\) (sulfurous acid) b. \(\mathrm{ClO}_{2}(\mathrm{OH}) \Rightarrow \mathrm{HClO}_2\) (chlorous acid) c. \(\mathrm{HPO}(\mathrm{OH})_{2} \Rightarrow \mathrm{H}_3\mathrm{PO}_3\) (phosphorous acid)

Explain why the given acids are acids and not bases

The main reason these compounds behave as acids instead of bases is their ability to donate protons (H⁺ ions) to other compounds. In their structural formulas, the hydrogen atoms are bonded to highly electronegative oxygen atoms, making the O-H bond highly polar. Consequently, the hydrogen can be easily released as a proton (H⁺ ion), which is the characteristic of an acidic compound. When these acids dissolve in water, they release H⁺ ions, increasing the hydrogen ion concentration in the solution.

Differentiate between the given acids and bases like \(\mathrm{NaOH}\) and \(\mathrm{KOH}\)

Bases like \(\mathrm{NaOH}\) and \(\mathrm{KOH}\), although containing a hydroxyl group, can accept protons (H⁺ ions) from other compounds. In \(\mathrm{NaOH}\) and \(\mathrm{KOH}\), the OH⁻ ion (hydroxide) is released when dissolved in water. The hydroxide is a highly reactive ion that readily accepts a proton to form water. In summary, the given acids are classified as such due to their ability to donate protons, while bases like \(\mathrm{NaOH}\) and \(\mathrm{KOH}\) are classified as bases because they can accept protons.

Key concepts

Properties of Acids and Bases

Understanding the properties of acids and bases is fundamental in chemistry. Acids, such as those in the given exercise, are characterized by their ability to donate protons (hydrogen ions, H⁺). This is due to the presence of polar covalent bonds between hydrogen and more electronegative elements, such as oxygen in the hydroxyl group (-OH). The more electronegative oxygen atom attracts the electron density from the bond, making the hydrogen more likely to be released as a proton.

On the other hand, bases are substances that can accept protons. This often involves a hydroxide ion (OH⁻), as seen in substances like NaOH and KOH. In water, these bases dissociate to release hydroxide ions, which are eager to bond with protons, thereby decreasing the hydrogen ion concentration in the solution.

This fundamental difference in behavior between acids and bases is expressed by the Bronsted-Lowry acid-base theory, which defines acids as proton donors and bases as proton acceptors. The tendency of an acid or base to donate or accept protons is influenced by the structure and stability of the resulting products after the donation or acceptance of the proton.

Acid-Base Reactions

The behavior of acids and bases in chemical reactions is governed by their ability to exchange protons. In an acid-base reaction, an acid donates a proton to a base. This kind of reactance is fundamental in chemistry and is known as a neutralization reaction when an acid reacts with a base to produce a salt and water.

For instance, when the acids from our exercise react with a base, they will lose a hydrogen ion, forming its conjugate base. For example, H₂SO₃ will become HSO₃⁻ after donating a proton. Similarly, a base like KOH will accept a proton to form water, H₂O, and its conjugate acid, K⁺.

These reactions often involve a transfer of electrons and can be represented in equation form to show the movement of protons from acids to bases. Understanding these reactions is crucial because they are pervasive in both natural processes, like metabolic pathways in the body, and industrial applications, such as the making of fertilizers and pH regulation in swimming pools.

Hydroxyl Groups in Oxyacids

The presence of the hydroxyl group (-OH) is a constant feature in all oxyacids, which makes them distinct from other acids. A hydroxyl group is composed of an oxygen atom bonded to a hydrogen atom. In oxyacids, this group is directly attached to a central atom, often a nonmetal, like the sulfur in H₂SO₃ or the phosphorous in H₃PO₃.

Oxyacids can have one or more hydroxyl groups attached to the central atom, and the acidity of these compounds is directly related to the number of these groups. For example, H₂SO₄ is a stronger acid than H₂SO₃ because it has one more hydroxyl group that can donate a proton.

The strength of an acid can also depend on other factors like the electronegativity of the central atom and the stability of the conjugate base formed after proton donation. By understanding the role of hydroxyl groups in the structure of oxyacids, students can begin to predict the properties and reactivity of these acids in various chemical contexts.