Sinhyu/iStock/GettyImages
When you think of acids—vinegar, bleach, or even the tartness of citric acid—you’re engaging with the pH scale, the tool that quantifies acidity in aqueous solutions. Understanding how to translate the intrinsic acidity of a substance (its acid dissociation constant (Ka)) into a measurable pH value is essential for chemists, educators, and anyone working with acids.
In water, an acid donates a proton (H⁺) and becomes an anion. The freed proton associates with a water molecule, forming the hydronium ion (H₃O⁺). The original acid becomes its conjugate base. For example, carbonic acid (H₂CO₃) dissociates to H₃O⁺ and bicarbonate (HCO₃⁻).
Strong acids such as hydrochloric acid (HCl) release protons readily even in already acidic media, whereas weak acids only ionize appreciably when the surrounding proton concentration is low (i.e., at higher pH).
The pH scale is logarithmic, ranging from about 1 (very acidic) to 14 (very basic). It is defined by:
\(\text{pH} = -\log_{10}[\text{H}^+]\)
where \([\text{H}^+]\) is the molar concentration of free protons. Each tenfold increase in proton concentration lowers the pH by one unit.
Example: A 0.025 M proton solution has
\(\text{pH} = -\log_{10}(0.025) = 1.602\)
Ka quantifies an acid’s tendency to dissociate:
\(K_a = \dfrac{[A^-][\text{H}_3\text{O}^+]}{[HA]}\)
Higher Ka values indicate stronger acids, meaning more complete dissociation at equilibrium.
By taking the negative logarithm of Ka, we obtain the acid’s pKa:
\(\text{pKa} = -\log_{10}K_a\)
The Henderson–Hasselbalch equation links pH, pKa, and the ratio of conjugate base to acid:
\(\text{pH} = \text{pKa} + \log_{10}\dfrac{[A^-]}{[HA]}\)
This relationship is especially useful for buffer solutions, where both the acid and its conjugate base are present.
Example: Acetic acid (CH₃COOH) has \(K_a = 1.77 \times 10^{-5}\). If only 10 % of the acid is dissociated, then \([A^-]/[HA] = 0.1\). First, compute the pKa:
\(\text{pKa} = -\log_{10}(1.77 \times 10^{-5}) = 4.75\)
Then apply Henderson–Hasselbalch:
\(\text{pH} = 4.75 + \log_{10}(0.1) = 4.75 - 1 = 3.75\)
At a pH equal to the pKa, the concentrations of acid and conjugate base are equal, meaning 50 % of the acid is dissociated.
These equations provide a straightforward method to predict the pH of any solution where the acid’s Ka and concentration are known.
• The pH scale measures proton concentration on a logarithmic scale.
• Ka expresses an acid’s dissociation propensity; pKa is its logarithmic counterpart.
• The Henderson–Hasselbalch equation bridges pH, pKa, and the acid/base ratio, enabling accurate pH calculations for buffers and weak acids.
