4.76
0.000017378
9.24
0.0000000006
4.76
0.000017378
9.24
0.0000000006
The pKa Calculator converts between the acid dissociation constant (Ka) and its negative logarithm pKa, two of the most important quantities in acid-base chemistry. The pKa value characterizes the strength of an acid — a lower pKa indicates a stronger acid that dissociates more readily in water. The relationship is simple: pKa = -log₁₀(Ka) and Ka = 10⁻ᵖᴷᵃ. This calculator also computes the corresponding pKb and Kb for the conjugate base using the relationship pKa + pKb = 14 (at 25°C). Understanding pKa is essential in organic chemistry, biochemistry, pharmacology, and environmental chemistry, where it predicts protonation states, buffer capacity, drug absorption, and the environmental fate of chemicals.
The acid dissociation reaction is: HA ⇌ H⁺ + A⁻
The equilibrium constant for this reaction is:
Ka = [H⁺][A⁻] / [HA]
Taking the negative logarithm gives: pKa = -log₁₀(Ka)
Strong acids like HCl have very large Ka values (Ka >> 1) and negative or near-zero pKa values. Weak acids like acetic acid have small Ka values (Ka = 1.75 × 10⁻⁵) and moderate pKa values (pKa = 4.76). Very weak acids like water have extremely small Ka values and high pKa values.
The conjugate base relationship follows from the water autoionization equilibrium: Ka × Kb = Kw = 10⁻¹⁴ at 25°C. Therefore pKa + pKb = 14. A strong acid (low pKa) has a weak conjugate base (high pKb), and vice versa.
In biochemistry, pKa values determine which form of an amino acid, drug, or metabolite predominates at a given pH. At pH = pKa, exactly 50% of the molecule is protonated and 50% is deprotonated. This is the basis of the Henderson-Hasselbalch equation and is central to understanding buffer systems, protein structure, and drug bioavailability.
Lower pKa means a stronger acid that donates protons more readily. Negative pKa values indicate very strong acids. A pKa of 4-5 is typical for weak organic acids (carboxylic acids). A pKa of 9-10 is typical for ammonium groups and phenols. The pKb of the conjugate base tells you about base strength — lower pKb means a stronger base. At any pH, you can determine the predominant form: if pH < pKa, the protonated form (HA) predominates; if pH > pKa, the deprotonated form (A⁻) predominates.
Inputs
Results
Acetic acid has Ka = 1.75 × 10⁻⁵, giving pKa ≈ 4.76. Its conjugate base (acetate) has pKb = 9.24, confirming it is a weak base.
Inputs
Results
NH₄⁺ has pKa = 9.25, so Ka = 5.62 × 10⁻¹⁰. Its conjugate base NH₃ has pKb = 4.75, making ammonia a moderately strong base.
pKa is inversely related to acid strength. Lower pKa = stronger acid. HCl (pKa ≈ -7) is a much stronger acid than acetic acid (pKa = 4.76). The pKa represents the pH at which the acid is exactly half-dissociated.
Yes. Strong acids like HCl (pKa ≈ -7), H₂SO₄ (pKa₁ ≈ -3), and HNO₃ (pKa ≈ -1.4) have negative pKa values. This means Ka > 1, indicating the equilibrium strongly favors dissociation.
For a conjugate acid-base pair: pKa + pKb = 14 at 25°C (since Ka × Kb = Kw = 10⁻¹⁴). A strong acid (low pKa) always has a weak conjugate base (high pKb), and vice versa.
A drug's pKa determines its ionization state at physiological pH (7.4). Only the un-ionized form can cross lipid membranes efficiently. Medicinal chemists tune pKa values to optimize oral absorption, blood-brain barrier penetration, and renal elimination of drug molecules.
Ka and pKa are temperature-dependent. For most acids, Ka increases slightly with temperature (pKa decreases). The change is usually small over normal laboratory temperature ranges but can be significant for precise biochemical work. Additionally, Kw changes with temperature, so the pKa + pKb = 14 relationship shifts.
Polyprotic acids (like H₂SO₄, H₃PO₄) have multiple dissociation steps, each with its own Ka. Ka₁ > Ka₂ > Ka₃ always, because removing each successive proton becomes harder. For phosphoric acid: pKa₁ = 2.15, pKa₂ = 7.20, pKa₃ = 12.35.
Roboculator Team
The Roboculator Team explains calculations, planning tools, and practical formulas in clear language for real-life situations.
How helpful was this calculator?
Be the first to rate!
