Question

Arrange the following \(0.10 \mathrm{M}\) solutions in order of increasing acidity (decreasing pH): (i) \(\mathrm{NH}_{4} \mathrm{NO}_{3}\) (ii) \(\mathrm{NaNO}_{3}\), (iii) $$ \mathrm{CH}_{3} \mathrm{COONH}_{4} \text { , (iv) } \mathrm{NaF} \text { , (v) } \mathrm{CH}_{3} \mathrm{COONa} \text { . } $$

Short answer

The order of increasing acidity (decreasing pH) for the given solutions is: (v) \(\mathrm{CH}_{3} \mathrm{COONa}\), (iv) \(\mathrm{NaF}\), (ii) \(\mathrm{NaNO}_{3}\), (iii) \(\mathrm{CH}_{3} \mathrm{COONH}_{4}\), and (i) \(\mathrm{NH}_{4} \mathrm{NO}_{3}\).

Step-by-step solution

Identify the ions present in each solution

In each of the given solutions, we have different ions present. Let's write them down: (i) \(\mathrm{NH}_{4} \mathrm{NO}_{3}\): \(\mathrm{NH}_{4}^{+}\) and \(\mathrm{NO}_{3}^{-}\) (ii) \(\mathrm{NaNO}_{3}\): \(\mathrm{Na}^{+}\) and \(\mathrm{NO}_{3}^{-}\) (iii) \(\mathrm{CH}_{3} \mathrm{COONH}_{4}\): \(\mathrm{CH}_{3} \mathrm{COO}^{-}\) and \(\mathrm{NH}_{4}^{+}\) (iv) \(\mathrm{NaF}\): \(\mathrm{Na}^{+}\) and \(F^{-}\) (v) \(\mathrm{CH}_{3} \mathrm{COONa}\): \(\mathrm{CH}_{3} \mathrm{COO}^{-}\) and \(\mathrm{Na}^{+}\)

Identify the acidic or basic properties of each ion

Now, let's identify the acidic or basic properties of the ions in each solution: - \(\mathrm{NH}_{4}^{+}\): Acidic (ammonium ion) - \(\mathrm{NO}_{3}^{-}\): Neutral (nitrate ion) - \(\mathrm{Na}^{+}\): Neutral (sodium ion) - \(\mathrm{CH}_{3} \mathrm{COO}^{-}\): Basic (acetate ion) - \(F^{-}\): Basic (fluoride ion)

Determine the acidity/basicity of solutions

Using the information from steps 1 and 2, we can now determine the acidity or basicity of each solution: (i) \(\mathrm{NH}_{4} \mathrm{NO}_{3}\): Acidic due to \(\mathrm{NH}_{4}^{+}\) ion (other ion is neutral) (ii) \(\mathrm{NaNO}_{3}\): Neutral, both ions are neutral (iii) \(\mathrm{CH}_{3} \mathrm{COONH}_{4}\): Mixed acidity and basicity (acetate ion is basic, ammonium ion is acidic) (iv) \(\mathrm{NaF}\): Basic due to \(F^{-}\) ion (other ion is neutral) (v) \(\mathrm{CH}_{3} \mathrm{COONa}\): Basic due to \(\mathrm{CH}_{3} \mathrm{COO}^{-}\) ion (other ion is neutral)

Arrange the solutions in order of increasing acidity

Now that we have established the acidity or basicity of each solution, we can arrange them in order of increasing acidity (decreasing pH): 1. \(\mathrm{CH}_{3} \mathrm{COONa}\) (basic solution) 2. \(\mathrm{NaF}\) (basic solution) 3. \(\mathrm{NaNO}_{3}\) (neutral solution) 4. \(\mathrm{CH}_{3} \mathrm{COONH}_{4}\) (mixed acidity and basicity) 5. \(\mathrm{NH}_{4} \mathrm{NO}_{3}\) (acidic solution) In conclusion, the order of increasing acidity (decreasing pH) for the given solutions is: (v), (iv), (ii), (iii), (i).

Key concepts

Acid-Base Equilibrium

Understanding acid-base equilibrium is crucial for analyzing the acidity or basicity of solutions. At the heart of this concept lies the dissociation of water into hydrogen ions (\(H^+\)) and hydroxide ions (\(OH^-\)), which is a reversible reaction. When an acid is dissolved in water, it increases the concentration of hydrogen ions, whereas a base increases the concentration of hydroxide ions. The various ions contribute differently to the acidity or basicity of the solution. For example, ammonium ion (\(NH_4^+\)) is an acidic ion because it can donate a hydrogen ion to the water, while acetate ion (\(CH_3COO^-\)) is a basic ion as it can accept hydrogen ions.

The exercise provided illuminates the importance of recognizing whether an ion will act as an acid or a base in water. The balance between the hydrogen ions and hydroxide ions provided by the dissolved substances dictates the pH of the solution, representing its acid-base equilibrium. As seen in the solutions, the presence of different ions dictates the nature of the solutions' acidity or basicity, and this principle is also crucial when interpreting reactions involving acid-base pairs in more complex chemical systems.

pH Scale

The pH scale is a measure of the acidity or basicity of an aqueous solution. It is a logarithmic scale based on the molarity of hydrogen ions (\(H^+\)) in solution, where a lower pH corresponds to a higher concentration of hydrogen ions and a higher acidity. The scale typically ranges from 0 to 14, with 7 being neutral, values less than 7 representing acidic solutions, and values greater than 7 indicating basic or alkaline solutions.

Understanding the pH scale is essential when comparing the acidity of the solutions as in the exercise. For instance, a solution of ammonium nitrate (\(NH_4NO_3\)) will have a lower pH (higher acidity) compared to neutral sodium nitrate (\(NaNO_3\)) because of the presence of the acidic ammonium ion. Therefore, the pH scale not only provides an immediate sense of the solution's acidity but also aids in arranging various solutions in order of increasing or decreasing acidity. This ease of comparison is crucial as it helps visualize the relative strength of acidic or basic solutions in a straightforward and quantitative manner.

Ionic Composition of Solutions

Analyzing the ionic composition of solutions is key to understanding their chemical behavior. The ions present in a solution can significantly affect its physical and chemical properties, including conductivity, reactivity, and pH. The ions originate from dissociation of soluble ionic compounds in water. For example, upon dissolution, sodium fluoride (\(NaF\)) will separate into sodium (\(Na^+\)) and fluoride (\(F^-\)) ions.

The exercise demonstrates that each ion's unique properties contribute to the overall nature of a solution. While neutral ions like nitrate (\(NO_3^-\)) have no significant effect on the solution's pH, the acidic ammonium ion increases the solution's acidity. Conversely, basic ions like fluoride and acetate will increase its basicity. When tackling chemistry problems involving ionic solutions, it is crucial to consider the behavior of individual ions as it helps predict the outcome of chemical reactions and provides insights into the solution's characteristics.