Ka (Acid Dissociation Constant) Calculators

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Ka — the acid dissociation constant — is the equilibrium constant that describes how completely a weak acid dissociates in water. A large Ka means strong dissociation and a stronger acid; a small Ka means weak dissociation and a weaker acid. Ka and its logarithmic form pKa are central to acid-base chemistry, buffer calculations, titration analysis, and biochemistry. Every pH calculation involving a weak acid depends on Ka, making it one of the most important constants in analytical and biological chemistry.

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What Is Ka?

For a weak acid HA dissociating in water:

HA ⇌ H⁺ + A⁻

The acid dissociation constant is:

Ka = [H⁺][A⁻] / [HA]

Where square brackets denote molar concentrations at equilibrium. Ka is dimensionless (or expressed in mol/L depending on convention) and is a fixed value at a given temperature.

pKa: The Logarithmic Form

Because Ka values range over many orders of magnitude, the logarithmic form pKa is more convenient:

pKa = −log₁₀(Ka)

A lower pKa means a stronger acid (larger Ka, more dissociation). Examples:

  • Hydrofluoric acid: Ka = 6.3 × 10⁻⁴, pKa = 3.17
  • Acetic acid: Ka = 1.8 × 10⁻⁵, pKa = 4.76
  • Ammonium ion: Ka = 5.6 × 10⁻¹⁰, pKa = 9.25
  • Water: Ka = 1.8 × 10⁻¹⁶, pKa = 15.74

Calculating pH from Ka

For a weak acid at concentration C with Ka ≪ C (the standard approximation):

[H⁺] ≈ √(Ka × C)
pH ≈ ½(pKa − log C)

Example: For 0.1 M acetic acid (Ka = 1.8 × 10⁻⁵):
[H⁺] = √(1.8 × 10⁻⁵ × 0.1) = √(1.8 × 10⁻⁶) = 1.34 × 10⁻³ M → pH ≈ 2.87

Ka and Buffer Chemistry

pKa is central to the Henderson-Hasselbalch equation for buffers: pH = pKa + log([A⁻]/[HA]). A buffer works best when pH is within ±1 unit of pKa — at pH = pKa, the buffer has maximum capacity. This is why knowing pKa values is essential for selecting the right buffer for any experiment.

Ka Relationships

For a conjugate acid-base pair, Ka and Kb are related by the water autoionization constant Kw:

Ka × Kb = Kw = 1.0 × 10⁻¹⁴ (at 25°C)
pKa + pKb = 14

Glossary

Ka (Acid Dissociation Constant)
The equilibrium constant for the dissociation of a weak acid in water: Ka = [H⁺][A⁻]/[HA]. Larger Ka means stronger dissociation and a stronger acid.
pKa
The negative base-10 logarithm of Ka (pKa = −log Ka). Lower pKa indicates a stronger acid. The pKa equals the pH at which the acid is 50% dissociated — equal concentrations of HA and A⁻.
Conjugate Acid-Base Pair
An acid and the base formed when it loses a proton, or a base and the acid formed when it gains a proton. Ka × Kb = Kw for any conjugate pair at 25°C.

Frequently Asked Questions

A large Ka indicates that the acid dissociates extensively in water — the equilibrium strongly favors products (H⁺ and A⁻). This means the acid is relatively strong. Conversely, a small Ka means the acid dissociates very little and is weak. Strong acids like HCl and H₂SO₄ have Ka values much greater than 1; typical weak acids have Ka values between 10⁻² and 10⁻¹².

pKa = −log₁₀(Ka). For Ka = 1.8 × 10⁻⁵: pKa = −log(1.8 × 10⁻⁵) = −(−4.74) = 4.74. To convert back: Ka = 10^(−pKa). Lower pKa values correspond to stronger acids; higher pKa values indicate weaker acids.

In the Henderson-Hasselbalch equation — pH = pKa + log([A⁻]/[HA]) — Ka determines the center of the buffer's working range. To buffer at a specific pH, choose a weak acid whose pKa is within ±1 unit of the target. The ratio of conjugate base to acid needed can be calculated from: [A⁻]/[HA] = 10^(pH − pKa).

For a conjugate acid-base pair, Ka (of the acid) × Kb (of the conjugate base) = Kw = 1.0 × 10⁻¹⁴ at 25°C. Equivalently, pKa + pKb = 14. A strong acid (large Ka, small pKa) has a very weak conjugate base (small Kb, large pKb), and vice versa.