Ka & pKa Converter

⚗️ Ka & pKa Converter – Acid Dissociation Made Simple

The Ka & pKa Converter is an essential chemistry tool for students, biochemists, pharmacists, and researchers who routinely work with acid-base equilibria. It converts between the acid dissociation constant (Ka) and its logarithmic form (pKa), calculates degree of dissociation, estimates equilibrium pH, and enables side-by-side comparison of multiple acids — all in one interface.

What Are Ka and pKa?

When a weak acid HA dissolves in water, it partially dissociates according to the equilibrium: HA ⇌ H⁺ + A⁻. The acid dissociation constant Ka quantifies the extent of this equilibrium:

Ka = [H⁺][A⁻] / [HA]

Because Ka values span many orders of magnitude (from 10 for strong acids to 10⁻¹⁵ for very weak ones), chemists use the pKa scale — the negative base-10 logarithm of Ka:

pKa = −log₁₀(Ka)     Ka = 10^(−pKa)

A lower pKa means a stronger acid. Acetic acid (pKa = 4.76) is much stronger than ammonium (pKa = 9.25). Strong mineral acids like HCl have negative pKa values (pKa ≈ −7).

Acid Strength Classification

Degree of Dissociation and Equilibrium pH

For a weak acid at initial concentration C, the exact equilibrium hydrogen ion concentration is found by solving the quadratic:

x² + Ka·x − Ka·C = 0 x = (−Ka + √(Ka² + 4·Ka·C)) / 2

where x = [H⁺] = [A⁻]. The degree of dissociationα = x / C expresses the fraction of the acid that ionised. The equilibrium pH is then pH = −log₁₀(x). This quadratic method is exact — it does not rely on the simplifying approximation (valid only when α < 5%) that C − x ≈ C.

Henderson–Hasselbalch and Buffer Design

The Henderson–Hasselbalch equation connects pKa to solution pH for a conjugate acid/base pair:

pH = pKa + log₁₀([A⁻] / [HA])

At the half-equivalence point of a titration, exactly half the acid has been neutralised, so [A⁻] = [HA] and log(1) = 0; therefore pH = pKa. This is why knowing the pKa is the first step in designing a buffer: choose a weak acid whose pKa is within one unit of your target pH for maximum buffering capacity.

Multi-Acid Comparison Mode

When comparing several acids — for instance, in pharmaceutical research to predict drug ionisation at physiological pH (7.4), or in food chemistry to balance flavour profiles — the multi-acid table ranks them by pKa and shows the degree of dissociation at a shared concentration. This makes relative acid strength immediately visible.

Polyprotic Acids

Diprotic acids such as carbonic acid (H₂CO₃), oxalic acid, and sulfurous acid undergo two successive dissociation steps, each with its own Ka and pKa. The converter handles both steps:

H₂A  ⇌  H⁺ + HA⁻     Ka1, pKa1 HA⁻  ⇌  H⁺ + A²⁻     Ka2, pKa2

Ka2 is always smaller than Ka1 (pKa2 > pKa1) because removing a proton from an already-negative ion is harder. At any given pH, you can determine the dominant species from the two pKa values. For example, at blood pH 7.4, bicarbonate (HCO₃⁻) dominates because pKa1 = 6.37 and pKa2 = 10.33.

Practical Applications

Ka and pKa values are fundamental across many disciplines:

• Pharmaceutical sciences — predicting drug absorption (most drugs must be unionised to cross cell membranes; pKa determines the fraction ionised at gut pH)
• Biochemistry — understanding enzyme active sites (histidine pKa ≈ 6 makes it a key proton shuttle at physiological pH)
• Environmental chemistry — modelling pesticide or pollutant speciation in natural waters
• Analytical chemistry — selecting indicators for titrations and designing HPLC mobile phases
• Food science — controlling preservation, flavour balance, and fermentation using organic acid pKa values

Scientific Notation Support

Ka values are often inconveniently small numbers. This tool accepts all common input formats: plain decimal (0.000018), engineering notation (1.8e-5or 1.8E-5), and Unicode superscript style (1.8×10⁻⁵). All three are normalised internally to a standard floating-point number before calculation.