Question

The following boxes represent aqueous solutions containing a weak acid, HA and its conjugate base, \(A^{-}\). Water molecules, hydronium ions, and cations are not shown. Which solution has the highest pH? Explain. [Section 17.1]

Short answer

The solution with the highest concentration of \(A^{-}\) relative to HA has the highest pH.

Step-by-step solution

Understand the Components

In this exercise, we're focusing on weak acid HA and its conjugate base, \(A^{-}\). The pH of a solution depends on the concentration of these components.

Identify pH Relationships

pH is an indicator of acidity, with lower pH values being more acidic and higher values being more basic. In a solution containing a weak acid and its conjugate base, the pH is determined by the ratio of \([A^{-}]/[HA]\). A higher ratio implies a higher pH.

Evaluate Solutions

View each solution and note the relative concentrations of HA and \(A^{-}\). The solution where \([A^{-}]\) is greater than \([HA]\) will have the highest pH, since it indicates a greater presence of the conjugate base which shifts the solution towards basicity.

Determine the Highest pH

Compare the boxes visually for each solution. Identify the box with the largest number of \(A^{-}\) compared to \(HA\). This solution will exhibit the highest pH as it is less acidic due to the larger concentration of the conjugate base.

Key concepts

Understanding Weak Acids

A weak acid is a type of acid that only partially dissociates in solution. This means when it is dissolved in water, only a small portion of the acid molecules donate hydrogen ions ( H⁺ ) to the solution, forming hydronium ions ( H₃O⁺ ). Weak acids are important in chemistry because they do not release all their hydrogen ions at once, which makes them less aggressive in changing the pH of a solution compared to strong acids.

Here's how you can identify a weak acid:

• It doesn't completely ionize in water.
• The equilibrium constant for its dissociation, known as the acid dissociation constant ( K_a ), is relatively small.
• In the molecular equation, it often remains mostly in its undissociated form.

Examples of weak acids include acetic acid ( CH₃COOH ), citric acid, and formic acid. Because of their partial ionization, weak acids create a balance between the undissociated acid and the ions formed in the solution. As a result, they're often used in buffer solutions where maintaining a stable pH is important.

The Role of Conjugate Bases

In the chemical world, every acid has a secretly dependent partner called a conjugate base. Simply put, a conjugate base is what remains of an acid molecule after it has given up a hydrogen ion ( H⁺ ). In the case of weak acids, this conjugate base plays a crucial part in the acid-base equilibrium.

The conjugate base is essential because:

• It can take up hydrogen ions to reform the acid, demonstrating the reversible nature of acid-base reactions.
• Its presence helps in maintaining equilibrium by reacting with hydronium ions, thereby neutralizing some of the acidity.
• It is instrumental in buffer solutions, which help resist drastic changes in pH when small amounts of acid or base are added.

For example, in an acetic acid solution ( CH₃COOH ), the acetate ion ( CH₃COO⁻ ) is the conjugate base. It can "soak up" protons (H⁺ ions) to help maintain the solution's pH by counteracting added acids or bases. The interplay between weak acids and their conjugate bases is a beautiful dance that maintains the delicate balance in acid-base equilibrium.

Exploring Acid-Base Equilibrium

Acid-base equilibrium is a fundamental concept that describes the balance between acids and bases in a solution. It's all about the continuous interplay between the dissociation of acids to release hydrogen ions ( H⁺ ) and the recombination of those ions with conjugate bases to form the original acid. In solutions with weak acids, this equilibrium is especially important.

Several things happen in acid-base equilibrium:

• Weak acids only partially dissociate, which means not all molecules break apart into ions, maintaining a dynamic equilibrium between the intact acid and its dissociated ions.
• The equilibrium can shift towards more molecules dissociating or recombining based on changes to the system, such as changes in concentration or temperature.
• This balance is crucial for buffer solutions that help resist changes in pH.

The equilibrium position is determined by the acid dissociation constant ( K_a ), which indicates how strongly an acid donates protons to the solution. Larger K_a values suggest stronger acids (more dissociation), while smaller values indicate weaker acids (less dissociation). Understanding this equilibrium helps us make predictions about how an increase in acid or base concentration will affect the pH, as well as how solutions will react to external stresses like dilution.