degree of dissociation
Imagine you drop 100 sugar-cube-sized particles into water and ask: how many of them came apart? If 30 split into pieces and 70 stayed whole, the fraction that came apart is 0.30. The degree of dissociation is exactly that fraction for a chemical species — the share of the original molecules that have broken into smaller parts (ions or fragments) at equilibrium.
Written as the symbol α (alpha), the degree of dissociation runs from 0 (nothing has split) to 1 (everything has split). It is often given as a percentage. For a weak acid or a partly dissociating gas, α tells you how complete the breakup is, and it ties directly to the equilibrium constant: knowing α and the starting amount lets you compute K, and knowing K and the conditions lets you predict α.
Its great practical value is that α responds to conditions in revealing ways. Diluting a weak electrolyte usually increases its degree of dissociation; raising temperature often does too. The honest caveat is that α is not a fixed property of a substance — it depends on concentration, temperature, and what else is present (the common-ion effect can suppress it). So always quote α together with the conditions under which it was measured.
In a 0.1 mol/L solution of acetic acid, only about 1.3% of the molecules ionise into ions, so α ≈ 0.013. Dilute the same acid tenfold and the degree of dissociation rises noticeably — a hallmark of weak electrolytes that Ostwald's dilution law captures.
α is the fraction that has broken apart — it grows on dilution.
Degree of dissociation is not the same as the equilibrium constant. α changes with concentration and temperature; the constant changes only with temperature. The two are linked, but they answer different questions.