HA

Definition

HA, or Brønsted-Lowry acid, is a chemical species that can donate a proton (H+) to another substance, thereby acting as an acid in a chemical reaction. This term is particularly relevant in the context of understanding Brønsted-Lowry acid-base theory and the relative strengths of acids and bases.

5 Must Know Facts For Your Next Test

1. The strength of a Brønsted-Lowry acid, HA, is determined by its ability to donate a proton (H+) to a base in an aqueous solution.
2. The conjugate base of a Brønsted-Lowry acid, HA, is the species that remains after the acid has donated a proton (H+).
3. The equilibrium constant, $K_a$, is a measure of the strength of a Brønsted-Lowry acid, HA, with larger $K_a$ values indicating a stronger acid.
4. The pH of a solution containing a Brønsted-Lowry acid, HA, is determined by the concentration of H+ ions in the solution, which is related to the acid's $K_a$ value.
5. The relative strengths of Brønsted-Lowry acids can be compared using the concept of acid dissociation constant, $K_a$, with stronger acids having larger $K_a$ values.

Review Questions

• Explain how the strength of a Brønsted-Lowry acid, HA, is related to its ability to donate a proton (H+) in an aqueous solution.
  • The strength of a Brønsted-Lowry acid, HA, is directly related to its ability to donate a proton (H+) in an aqueous solution. Stronger acids have a greater tendency to donate protons, which is quantified by the acid dissociation constant, $K_a$. Acids with larger $K_a$ values are considered stronger because they are more likely to dissociate and release H+ ions in water, leading to a lower pH in the solution. The strength of an acid, HA, is a key factor in determining the equilibrium concentrations of the acid and its conjugate base in an aqueous solution.
• Describe the relationship between the strength of a Brønsted-Lowry acid, HA, and the strength of its conjugate base.
  • The strength of a Brønsted-Lowry acid, HA, is inversely related to the strength of its conjugate base. Stronger acids have weaker conjugate bases, and vice versa. This relationship is governed by the acid dissociation constant, $K_a$, and the corresponding conjugate base dissociation constant, $K_b$. Specifically, the product of $K_a$ and $K_b$ for a conjugate acid-base pair is equal to the ion product of water, $K_w$. As a result, if an acid is a strong donor of protons (H+), its conjugate base will be a weak acceptor of protons, and therefore a weak base.
• Analyze how the pH of a solution containing a Brønsted-Lowry acid, HA, is influenced by the acid's strength and concentration.
  • The pH of a solution containing a Brønsted-Lowry acid, HA, is directly influenced by the acid's strength and concentration. Stronger acids, with larger acid dissociation constants ($K_a$), will have a greater tendency to donate protons (H+) in an aqueous solution, leading to a lower pH. Conversely, weaker acids will have a smaller $K_a$ value and will produce a higher pH in the solution. Additionally, the concentration of the acid, HA, will also affect the pH, with higher concentrations of a strong acid resulting in a lower pH. By considering both the acid strength, as measured by $K_a$, and the acid concentration, one can predict and understand the pH of a solution containing a Brønsted-Lowry acid.

Related terms

Brønsted-Lowry Acid: A Brønsted-Lowry acid is a substance that can donate a proton (H+) to another substance, thereby acting as an acid in a chemical reaction.

Conjugate Base:

The conjugate base of a Brønsted-Lowry acid is the species that remains after the acid has donated a proton (H+) to another substance.

Acid Strength: The strength of an acid refers to its ability to donate protons (H+) in an aqueous solution, with stronger acids being more likely to donate protons.