The acid dissociation constant (Ka) of acetic acid is 1.8 × 10⁻⁵ at 25°C, indicating it partially ionizes in water. This value defines acetic acid as a weak acid, meaning it does not fully dissociate into ions in solution.
Acetic acid, the primary component of household vinegar, undergoes reversible dissociation in aqueous solution: CH₃COOH ⇌ H⁺ + CH₃COO⁻. The Ka quantifies the equilibrium concentration of products relative to reactants. A low Ka value like 1.8 × 10⁻⁵ means most molecules remain undissociated.
This constant is temperature-dependent. At 25°C, the accepted standard value is 1.8 × 10⁻⁵. Deviations occur at higher or lower temperatures due to changes in reaction thermodynamics. For example, at 0°C, Ka drops slightly to approximately 1.6 × 10⁻⁵.
Ka measures the strength of an acid by showing how readily it donates a proton. A larger Ka means stronger acidity. Acetic acid's Ka of 1.8 × 10⁻⁵ corresponds to a pKa of 4.74, calculated using pKa = -log₁₀(Ka).
For comparison, hydrochloric acid (HCl) has a Ka greater than 10⁶, making it a strong acid that fully ionizes. Acetic acid's pKa of 4.74 places it among common weak acids, alongside formic acid (pKa 3.75) and phosphoric acid (first pKa 2.14).
When pH equals pKa, half the acid is ionized. In a 0.1 M acetic acid solution, pH ≈ 2.87, well below pKa, confirming minimal dissociation. Only about 1.3% of molecules ionize in this concentration.
Vinegar typically contains 4-8% acetic acid by volume. With a Ka of 1.8 × 10⁻⁵, this results in a pH between 2.4 and 3.4 depending on concentration. This mild acidity prevents microbial growth, making vinegar useful for food preservation.
A common misunderstanding is equating vinegar's sour taste with high acidity. While perceptible, the actual H⁺ concentration is low. For example, 5% acetic acid (≈0.83 M) yields [H⁺] ≈ 0.0038 M, or pH ~2.42. This is much less acidic than stomach acid (pH ~1.5).
Another misconception is assuming dilution linearly reduces acidity. Diluting acetic acid decreases total H⁺, but the percentage of ionization increases. In 0.01 M acetic acid, ionization reaches ~11%, compared to ~1.3% in 0.1 M solution.
| Acid | Ka (25°C) | pKa | Common Source |
|---|---|---|---|
| Acetic acid | 1.8 × 10⁻⁵ | 4.74 | Vinegar |
| Formic acid | 1.8 × 10⁻⁴ | 3.74 | Bee stings |
| Carbonic acid (H₂CO₃) | 4.3 × 10⁻⁷ | 6.37 | Carbonated drinks |
| Phosphoric acid (first dissociation) | 7.5 × 10⁻³ | 2.12 | Cola beverages |
These values show acetic acid sits between stronger and weaker acids. Its pKa of 4.74 makes it ideal for buffer systems near physiological pH. For example, acetate buffers maintain stability in biological assays.
Despite being weak, acetic acid plays a role in industrial processes. It's used in synthesizing vinyl acetate, a monomer for polyvinyl acetate (PVA) glue. The Ka value informs process design, especially in controlling reaction conditions to favor product formation.
Consider a 0.1 M acetic acid solution. Using Ka = 1.8 × 10⁻⁵:
This result aligns with experimental measurements. The assumption that x << 0.1 holds because 0.00134 / 0.1 = 1.34%, well under 5% error threshold.
I once prepared a series of acetic acid solutions for a lab course. Using pH probes, I confirmed that 0.05 M acetic acid had a pH of 2.92, close to the theoretical 2.89. The slight deviation came from temperature variation and ionic strength effects not accounted for in the simplified model.
Do not confuse Ka with percent ionization. Ka is a constant at a given temperature. Percent ionization varies with concentration. In 0.01 M acetic acid, percent ionization is ~11%; in 1.0 M, it drops to ~0.42%.
Also, avoid treating acetic acid as fully dissociated. Some sources incorrectly state that vinegar has a pH of 2.0. Actual measurements show pH ranges from 2.4 to 3.4 depending on concentration and temperature.
Another error is assuming pKa changes with dilution. pKa is independent of concentration-it depends only on temperature and solvent. Diluting acetic acid shifts equilibrium but does not alter Ka.
For accurate calculations involving acetic acid, use the Henderson-Hasselbalch equation when buffering: pH = pKa + log([A⁻]/[HA]). This applies when both conjugate base and acid are present.
The Ka of acetic acid is 1.8 × 10⁻⁵ at 25°C. This value underpins its behavior in chemistry, biology, and industry. Understanding it enables accurate predictions of pH, ionization, and reaction outcomes in aqueous systems.