The dissociation constant ($K_a$) of acetic acid is 1.5 × 10⁻⁵ at 25°C. This value quantifies the extent to which acetic acid dissociates into its ions in an aqueous solution.
Understanding the Dissociation Constant ($K_a$)
The dissociation constant, often denoted as $K_a$, is a quantitative measure of the strength of an acid in solution. It represents the equilibrium constant for the dissociation of a weak acid into its constituent ions. A larger $K_a$ value indicates a stronger acid, meaning it dissociates more extensively in water, while a smaller $K_a$ value indicates a weaker acid that dissociates only slightly.
For a generic weak acid, HA, the dissociation reaction in water can be represented as:
HA(aq) + H₂O(l) ⇌ H₃O⁺(aq) + A⁻(aq)
The acid dissociation constant ($K_a$) is expressed as:
$K_a = \frac{[H_3O^+][A^-]}{[HA]}$
Where:
- $[H_3O^+]$ is the concentration of hydronium ions.
- $[A^-]$ is the concentration of the conjugate base.
- $[HA]$ is the concentration of the undissociated acid.
The Acetic Acid Dissociation Constant Value
Acetic acid (CH₃COOH) is a classic example of a weak organic acid. Its dissociation in water can be written as:
CH₃COOH(aq) + H₂O(l) ⇌ H₃O⁺(aq) + CH₃COO⁻(aq)
At a standard temperature of 25°C, the acid dissociation constant ($K_a$) for acetic acid is 1.5 × 10⁻⁵. This relatively small value confirms that acetic acid is a weak acid; it does not fully dissociate in water, maintaining a significant concentration of undissociated CH₃COOH molecules at equilibrium.
Factors Affecting $K_a$
While the $K_a$ value is generally considered constant for a given acid, it is specifically defined for a particular temperature. The standard reference temperature for most $K_a$ values is 25°C. Changes in temperature can slightly affect the equilibrium position and thus the $K_a$ value. Other factors like the solvent can also influence dissociation, but $K_a$ values are typically reported for aqueous solutions.
Practical Implications of Acetic Acid's $K_a$
The $K_a$ of acetic acid has several important practical implications:
- Weak Acid Nature: Its low $K_a$ value explains why acetic acid is considered a weak acid. Unlike strong acids (e.g., HCl), which dissociate almost completely, acetic acid dissociates only partially, leading to a lower concentration of H₃O⁺ ions and thus a higher pH for solutions of equivalent concentration.
- Buffer Systems: Acetic acid and its conjugate base, the acetate ion (CH₃COO⁻), form an effective buffer system. This acetic acid/acetate buffer is crucial in maintaining stable pH levels in chemical and biological systems, resisting significant pH changes upon the addition of small amounts of strong acid or base.
- Common Applications: Acetic acid is the main component of vinegar (typically 4-7% acetic acid by volume). Its mild acidity, governed by its $K_a$, makes it safe for use in food preservation, cooking, and cleaning agents. Its weak nature contributes to its relatively non-corrosive properties compared to strong acids.
Comparison of Common Weak Acids
To put acetic acid's $K_a$ into perspective, consider its value relative to other weak acids.
| Acid | Chemical Formula | Dissociation Constant ($K_a$) at 25°C | Relative Strength |
|---|---|---|---|
| Acetic Acid | CH₃COOH | 1.5 × 10⁻⁵ | Moderate Weak Acid |
| Formic Acid | HCOOH | 1.8 × 10⁻⁴ | Stronger Weak Acid |
| Carbonic Acid | H₂CO₃ | 4.3 × 10⁻⁷ | Weaker Weak Acid |
| Hydrocyanic Acid | HCN | 4.5 × 10⁻¹⁰ | Very Weak Acid |
As seen in the table, acetic acid is a moderately weak acid. For instance, formic acid is a stronger weak acid with a larger $K_a$, while hydrocyanic acid (HCN) is a much weaker acid with a significantly smaller $K_a$ value of 4.5 × 10⁻¹⁰, indicating it dissociates to a far lesser extent than acetic acid.
Understanding the dissociation constant of acetic acid is fundamental to comprehending its chemical behavior and its widespread applications in various fields.
