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Boiling is vaporization through the formation and growth of vapor bubbles within a liquid. It differs from surface evaporation, which can occur without bubbles in the bulk. At equilibrium, the boiling point of a pure liquid is set by equality of its saturation vapor pressure and the external pressure.

A real bubble must also overcome local hydrostatic pressure and the capillary pressure of its curved surface. Creating the interface costs energy, giving a nucleation barrier; a smooth clean vessel can therefore permit superheating. Surface cavities and boiling chips supply sites where bubbles form more readily. The observed onset depends on these sites and heating conditions, while the equilibrium boiling point remains a property specified at the chosen pressure.

McMaster University's 1997/98 laboratory manual compares simple distillation and fractional distillation using charges of $25\mathrm{mL}$ each of cyclohexane and p-xylene. Part B uses a $100\mathrm{mL}$ round-bottom flask and a fractionating column. The protocol specifies a boiling chip, a heating mantle on a ring support, and the top of the thermometer bulb midway in the opening toward the laboratory condenser. Heating is adjusted for roughly one drop per second. Students record head temperature every $2\mathrm{mL}$ collected, up to $45\mathrm{mL}$, and plot both runs against collected volume on the same axes.^[ref]

Charge, atmospheric operation, and collection rate provide comparison conditions; the fractionating column is changed. A sharper separation of temperature regions would support improved fractionation. It would not directly measure product mole fractions or prove purity. The manual distinguishes collected volume from composition and states that this pair forms no azeotrope. No student dataset is supplied here, and the constant-relative volatility example is not assigned to these liquids.^[ref]

Superheating is persistence of a phase above the equilibrium temperature at which a higher-temperature phase becomes favorable. The familiar case is a liquid above its boiling point at the imposed pressure, although delayed melting can also be discussed as superheating.

The liquid can remain in a metastable state because producing a vapor bubble costs interfacial energy. A sufficiently large bubble overcomes the nucleation barrier and grows, converting the liquid and consuming latent heat. Smooth clean containers can delay nucleation, while cavities and boiling chips provide heterogeneous nucleation sites. The observed temperature of sudden bubble formation therefore need not equal the equilibrium boiling point. The amount of superheating is a kinetic and preparation-dependent observation, rather than a new equilibrium phase boundary.