Weak Acid

Definition

A weak acid is a type of acid that only partially dissociates into hydrogen ions (H+) and the conjugate base of the acid when dissolved in water. This partial dissociation results in a solution with a pH that is higher than that of a strong acid at the same concentration.

5 Must Know Facts For Your Next Test

1. Weak acids are commonly found in biological systems, such as carbonic acid (H2CO3) in the blood and acetic acid (CH3COOH) in the digestive system.
2. The degree of dissociation of a weak acid is determined by its acid dissociation constant (Ka), with a higher Ka value indicating a stronger acid.
3. The Henderson-Hasselbalch equation is used to relate the pH of a weak acid solution to its Ka and the concentrations of the acid and its conjugate base.
4. Biological amines, such as histamine and serotonin, can act as weak acids or bases depending on the pH of their environment.
5. The pH of a weak acid solution can be adjusted by the addition of a base, which shifts the equilibrium and changes the concentrations of the acid and its conjugate base.

Review Questions

• Explain how the partial dissociation of a weak acid affects the pH of the solution compared to a strong acid.
  • Unlike strong acids, which completely dissociate in water, weak acids only partially dissociate, resulting in a solution with a higher pH. This is because the equilibrium between the undissociated acid and the dissociated hydrogen ions and conjugate base is reached at a higher pH value. The degree of dissociation is determined by the acid dissociation constant (Ka) of the weak acid, with a higher Ka indicating a stronger acid and a lower pH in the solution.
• Describe the role of weak acids in biological systems and how the Henderson-Hasselbalch equation is used to understand their behavior.
  • Weak acids are prevalent in biological systems, such as carbonic acid (H2CO3) in the blood and acetic acid (CH3COOH) in the digestive system. The Henderson-Hasselbalch equation, $\text{pH} = \text{pK}_\text{a} + \log\left(\frac{\left[\text{A}^-\right]}{\left[\text{HA}\right]}\right)$, is used to describe the relationship between the pH of a weak acid solution and the concentrations of the acid and its conjugate base. This equation allows for the prediction and manipulation of the pH in biological systems, which is crucial for maintaining homeostasis and proper physiological function.
• Analyze how the pH of a weak acid solution can be adjusted by the addition of a base, and explain the implications of this on the equilibrium and the concentrations of the acid and its conjugate base.
  • The pH of a weak acid solution can be adjusted by the addition of a base, which shifts the equilibrium between the undissociated acid and the dissociated hydrogen ions and conjugate base. When a base is added, it reacts with the hydrogen ions (H+) produced by the weak acid, reducing their concentration and increasing the concentration of the conjugate base. This shift in equilibrium is described by the Henderson-Hasselbalch equation, where the pH of the solution increases as the ratio of the conjugate base to the undissociated acid increases. This ability to manipulate the pH of weak acid solutions is crucial in biological systems, where maintaining appropriate pH levels is essential for proper physiological function.

Related terms

Strong Acid: A strong acid is an acid that completely dissociates into hydrogen ions (H+) and the conjugate base of the acid when dissolved in water, resulting in a solution with a low pH.

pH: pH is a measure of the acidity or basicity of a solution, with a range from 0 to 14. A pH less than 7 is acidic, 7 is neutral, and greater than 7 is basic.

Conjugate Base: The conjugate base of an acid is the species formed when the acid loses a proton (H+) and is left with a negative charge.