The equilibrium constant for acetic acid, often referred to as its acid dissociation constant (Ka), is 1.8 × 10⁻⁵.
Understanding the Equilibrium Constant (Ka)
The equilibrium constant, specifically the acid dissociation constant (Ka), quantifies the strength of an acid in solution. For a weak acid like acetic acid (CH₃COOH), it represents the extent to which the acid dissociates into its conjugate base and a hydrogen ion (proton) when dissolved in water. A smaller Ka value indicates a weaker acid, meaning it dissociates less, while a larger Ka value indicates a stronger acid.
Acetic acid dissociates in water according to the following reversible reaction:
CH₃COOH(aq) + H₂O(l) ⇌ CH₃COO⁻(aq) + H₃O⁺(aq)
The Ka expression for acetic acid is:
Ka = $\frac{[CH_3COO^-][H_3O^+]}{[CH_3COOH]}$
Where:
- $[CH_3COO^-]$ is the equilibrium concentration of the acetate ion.
- $[H_3O^+]$ is the equilibrium concentration of the hydronium ion (often simplified as $[H^+]$).
- $[CH_3COOH]$ is the equilibrium concentration of undissociated acetic acid.
The Specific Value for Acetic Acid
The acid dissociation constant (Ka) for acetic acid is 1.8 × 10⁻⁵. This value is a fundamental property of acetic acid and is widely used in chemical calculations and analyses. While this constant is typically provided in academic settings, it can also be determined experimentally by measuring the equilibrium concentrations of the species involved in the dissociation reaction.
- Ka Value: 1.8 × 10⁻⁵
Here's a comparison of acetic acid's Ka value with a few other common weak acids:
| Acid | Chemical Formula | Ka Value |
|---|---|---|
| Acetic Acid | CH₃COOH | 1.8 × 10⁻⁵ |
| Formic Acid | HCOOH | 1.8 × 10⁻⁴ |
| Hydrofluoric Acid | HF | 6.6 × 10⁻⁴ |
| Carbonic Acid | H₂CO₃ | 4.3 × 10⁻⁷ (Ka1) |
(Note: These are approximate values for comparison and can vary slightly depending on temperature and ionic strength.)
Significance and Implications of Ka
The relatively small Ka value of 1.8 × 10⁻⁵ for acetic acid indicates that it is a weak acid. This means that when acetic acid is dissolved in water, only a small fraction of its molecules dissociate into ions; the majority remain as undissociated CH₃COOH molecules.
- Weak Acidity: The low Ka signifies limited proton donation, resulting in a less acidic solution compared to strong acids (e.g., hydrochloric acid, sulfuric acid) which dissociate almost completely.
- Equilibrium Position: The equilibrium lies far to the left, favoring the undissociated acetic acid molecules.
- Buffering Capacity: Weak acids like acetic acid and their conjugate bases (acetate ion) are crucial components of buffer solutions, which resist changes in pH. Vinegar, a dilute solution of acetic acid, is a common household example demonstrating this property.
Practical Applications and Calculations
Understanding the Ka of acetic acid is vital in various chemical contexts:
- pH Calculation: Knowing Ka allows for the calculation of the pH of acetic acid solutions, which is essential in analytical chemistry and industrial processes. For example, to find the pH of a 0.1 M acetic acid solution, one would typically use an ICE table (Initial, Change, Equilibrium) in conjunction with the Ka value.
- Buffer Preparation: Acetic acid and its salt, sodium acetate, form an effective buffer system. The Ka value is directly used in the Henderson-Hasselbalch equation to prepare buffers of specific pH levels, crucial in biological and pharmaceutical applications.
- Titrations: The Ka value helps in predicting the equivalence point and selecting appropriate indicators for the titration of acetic acid with a strong base.
- Biological Systems: Acetic acid is a metabolic intermediate in many biological processes. Its acidic properties, governed by its Ka, are relevant to its role in cellular biochemistry.
For further reading on acid dissociation constants and weak acids, you can refer to resources like Chemistry LibreTexts on Acid Dissociation Constant or the IUPAC Gold Book definition of Acid Dissociation Constant.
