Question

Explain why ions such as \(\mathrm{Cl}^{-}, \mathrm{NO}_{3}^{-}\), and \(\mathrm{I}^{-}\)do not act as bases in water.

Short answer

The ions \(\mathrm{Cl}^{-}, \(\mathrm{NO}_{3}^{-}\) and \(\mathrm{I}^{-}\) do not act as bases in water because they are the anions of strong acids and thus are fully ionized, with no protons to donate. Even though these ions can theoretically accept protons, in the presence of water, the water molecules (being amphiprotic) are more likely to accept protons.

Step-by-step solution

Understanding Acids and Bases

Bases are substances that can donate electron pairs, or accept protons. They have the ability to increase the hydroxide ion concentration in water, and react with acids to form salts.

Familiarity with the Given Ions

\(\mathrm{Cl}^{-}, \(\mathrm{NO}_{3}^{-}\) and \(\mathrm{I}^{-}\) are anions of strong acids (HCl, HNO_3 and HI respectively). In water, they are fully ionized and thus do not have protons to donate.

The reason why these ions do not act as bases

Although these anions can accept protons in principle, they do not do so in water because they are the conjugate bases of strong acids. That means in water, which can function as a base or an acid (amphiprotic behaviour), the water molecules are more likely to accept protons. Therefore, in water they will not act as bases.

Key concepts

Hydroxide Ion

The hydroxide ion, denoted as \(OH^-\), plays a critical role in the realm of acid-base chemistry. Understanding the behavior of hydroxide ions is crucial for anyone delving into chemistry because:

• Hydroxide ions are key components in defining the pH of a solution, as they increase the alkalinity when their concentration rises.
• They form when a base accepts a proton (\(H^+\)) or when water dissociates into \(OH^-\) and \(H^+\) ions.
• Due to their negative charge, hydroxide ions are efficient at attracting protons, leading to neutralization reactions in which they react with acids to form water molecules.

Therefore, hydroxide ions are essential identifiers of basicity in aqueous solutions, directly influencing chemical reactions by interacting with acids to either donate electrons or accept protons. This behavior underpins the balancing act of acidity and basicity in a wide array of chemical processes.

Conjugate Base

In acid-base chemistry, conjugate bases are fascinating participants in chemical reactions. A conjugate base is formed when an acid donates a proton to a solvent. This concept elucidates the behavior observed in weak and strong acids.

• When an acid loses a \(H^+\), it transforms into its conjugate base; for example, hydrochloric acid (\(HCl\)) becomes chloride ion (\(Cl^-\)).
• The strength of a conjugate base is inversely related to the strength of its parent acid. Strong acids like \(HCl\) have very weak conjugate bases that scarcely interact with protons in solution.
• In the context of ions such as \(Cl^-\), \(NO_3^-\), and \(I^-\), their presence as conjugate bases of strong acids explains their lack of base-like behavior in water. This is because the equilibrium strongly favors retaining the acid in its protonated form.

Understanding conjugate bases is central to grasping why some ions do not act as typical bases in water. Their parent acids' potency and the resulting non-reactivity highlight the dynamics of chemical equilibriums.

Amphiprotic Behavior

Amphiprotic molecules have the unique ability to react both as an acid and a base. Water is the classic example of an amphiprotic substance. This versatile behavior is pivotal to understanding why certain ions like \(Cl^-\), \(NO_3^-\), and \(I^-\) do not act as bases in water. Here's why:

• An amphiprotic substance can either donate a \(H^+\) or accept one, depending on the substances it interacts with.
• In the presence of such ions, water's amphiprotic nature predominantly has it behave as a base, using its oxygen to form stronger bonds by accepting \(H^+\) ions.
• This means that in aqueous solutions, these ions are less likely to behave as bases because they do not compete effectively with water's amphiprotic capability to gain protons.

Hence, water's dual behavior heavily influences the reactions and equilibriums of other substances in solution, often overshadowing the basic potential of these anions in aqueous environments.