Problem 110 - Chemistry: The Central Science in SI Units, Global Edition

Question

Arrange the following \(0.10 \mathrm{M}\) solutions in order of increasing acidity: (i) \(\mathrm{HCOONH}_{4}\), (ii) \(\mathrm{NH}_{4} \mathrm{Br}\), (iii) \(\mathrm{NaNO}_{3}\), (iv) \(\mathrm{HCOOK},(\mathrm{v}) \mathrm{KF} .\)

Step-by-Step Solution

Answer

The order of increasing acidity for the given solutions is (iii) \(\mathrm{NaNO}_{3}\) < (iv) \(\mathrm{HCOOK}\) < (v) \(\mathrm{KF}\) < (ii) \(\mathrm{NH}_{4} \mathrm{Br}\) < (i) \(\mathrm{HCOONH}_{4}\).

Step 1: Identify acidic or basic properties of cations and anions for each solution

We will analyze each of the given solutions to identify the acidic or basic properties of their cations and anions. (i) \(\mathrm{HCOONH}_{4}\) - The cation is \(\mathrm{NH}_{4}^{+}\) (ammonium ion) which is acidic as it donates proton and the anion is \(\mathrm{HCOO}^{-}\) (formate ion) which is basic as it can accept a proton. (ii) \(\mathrm{NH}_{4} \mathrm{Br}\) - The cation is \(\mathrm{NH}_{4}^{+}\) (ammonium ion) which is acidic and the anion is \(\mathrm{Br}^{-}\) (bromide ion), which is neutral. (iii) \(\mathrm{NaNO}_{3}\) - The cation is \(\mathrm{Na}^{+}\) (sodium ion) which is neutral and the anion is \(\mathrm{NO}_{3}^{-}\) (nitrate ion) which is also neutral. (iv) \(\mathrm{HCOOK}\) - The cation is \(\mathrm{K}^{+}\) (potassium ion) which is neutral and the anion is \(\mathrm{HCOO}^{-}\) (formate ion) which is basic. (v) \(\mathrm{KF}\) - The cation is \(\mathrm{K}^{+}\) (potassium ion) which is neutral and the anion is \(\mathrm{F}^{-}\) (fluoride ion) which is basic.

Step 2: Analyze the acidic and basic properties of given solutions and arrange them in order

Below are the acidic and basic properties of the given solutions: (i) \(\mathrm{HCOONH}_{4}\) - Acidic cation and basic anion. (ii) \(\mathrm{NH}_{4} \mathrm{Br}\) - Acidic cation and neutral anion. (iii) \(\mathrm{NaNO}_{3}\) - Neutral cation and neutral anion. (iv) \(\mathrm{HCOOK}\) - Neutral cation and basic anion. (v) \(\mathrm{KF}\) - Neutral cation and basic anion. Now, keeping in mind that acidity is increased by having more acidic cations and decreased by having more basic anions, we can arrange these in order of increasing acidity: Neutral cation and neutral anion: \(\mathrm{NaNO}_{3}\) (iii) Neutral cation and less basic anion: \(\mathrm{HCOOK}\) (iv) Neutral cation and more basic anion: \(\mathrm{KF}\) (v) Acidic cation and less basic anion: \(\mathrm{NH}_{4} \mathrm{Br}\) (ii) Acidic cation and more basic anion: \(\mathrm{HCOONH}_{4}\) (i) So, the order of increasing acidity is (iii) \(\mathrm{NaNO}_{3}\) < (iv) \(\mathrm{HCOOK}\) < (v) \(\mathrm{KF}\) < (ii) \(\mathrm{NH}_{4} \mathrm{Br}\) < (i) \(\mathrm{HCOONH}_{4}\).

Key Concepts

Buffer Solutions

Understanding buffer solutions is essential when studying the acidity and basicity of solutions. A buffer solution is a special type of solution that resists changes in pH when small amounts of acid or base are added. In essence, it "buffers" the solution against dramatic pH changes.

Buffer solutions are typically made from a weak acid and its conjugate base or a weak base and its conjugate acid.

• For example, a common buffer system includes acetic acid ( CH₃COOH) and its conjugate base, acetate ( CH₃COO⁻).
• Similarly, ammonia ( NH₃) and ammonium ( NH₄⁺) form another classic buffer pair.

The presence of both an acidic and basic component allows the solution to neutralize added acids or bases.

This property is crucial in biological and chemical applications where maintaining a stable pH is necessary for proper functioning.

Buffer solutions are widely used in chemical laboratories, medical formulations, and various industrial applications.

Acid-Base Equilibria

Acid-base equilibria describe how acids and bases react with each other in water and influence the acidity or basicity of solutions. When discussing acid-base reactions, it's important to consider whether the acid or base fully dissociates in water.

There are two categories:

• **Strong acids** and **bases**, which dissociate completely in aqueous solutions. Examples include hydrochloric acid ( HCl) and sodium hydroxide ( NaOH).
• **Weak acids** and **bases**, which only partially dissociate in solutions. Examples are acetic acid ( CH₃COOH) and ammonia ( NH₃).

Acid-base equilibria explain how ions like H⁺ and OH⁻ affect pH levels. The equilibrium constant ( K_a for acids and K_b for bases) is used to predict the extent of dissociation and subsequently, the pH of a solution.

Understanding these interactions is critical for predicting the impact of substances in solution, particularly in reactions involving ionic compounds or buffer systems.

Ionic Compounds

Ionic compounds play a significant role in determining the acidity or basicity of a solution. These compounds consist of a cation (positively charged ion) and an anion (negatively charged ion). When dissolved in water, ionic compounds disassociate into their constituent ions. This disassociation can affect the pH of the solution, depending on the properties of the ions.

For example:

• **Sodium chloride ( NaCl)**, a neutral salt, completely dissociates into the neutral Na⁺ and Cl⁻ ions, having no effect on the solution's acidity.
• **Ammonium chloride ( NH₄Cl)**, on the other hand, dissociates into the acidic NH₄⁺ and neutral Cl⁻ ions, making the solution slightly acidic.

These characteristics of ionic compounds, such as ability of cations to donate protons or of anions to accept them, determine if the resulting solution will be acidic, basic, or neutral.

Understanding how these ions interact when dissolved is critical for predicting the outcomes in various chemical reactions and in formulating buffer solutions.