The determination of a substance’s boiling point provides chemists with a fundamental physical property, serving both to identify an unknown compound and to assess its purity. The boiling point is a characteristic constant for a pure substance under a specific external pressure. Understanding how this temperature is determined requires looking at the underlying physics of phase change and the practical methods used in the laboratory. The measurement process is a direct application of the principle of pressure equilibrium, which dictates when a liquid will transition into a gas.
Defining Boiling Point Through Vapor Pressure
Boiling is a phase transition where a liquid rapidly changes into a gas, forming vapor bubbles throughout the bulk of the liquid. This process is governed by internal vapor pressure, which is the pressure exerted by gaseous molecules above the liquid surface. As a liquid is heated, the kinetic energy of its molecules increases, raising the vapor pressure.
The liquid will only boil when its vapor pressure becomes equal to or slightly exceeds the external pressure, which is typically the atmospheric pressure. Once this equilibrium is reached, bubbles can form and expand anywhere within the liquid, allowing the substance to boil. At this specific temperature, the energy input is used to overcome the intermolecular forces and convert the liquid into a gas, rather than raising the liquid’s temperature.
The temperature at which this pressure balance occurs is the boiling point, and it remains constant throughout the boiling process for a pure substance. Any heat added after this point simply increases the rate of vaporization. This physical principle means that the boiling point is directly tied to the surrounding atmospheric pressure, as the internal vapor pressure must match the external force.
Standard Laboratory Measurement Techniques
In the laboratory, the boiling point of a substance is typically measured using one of two primary methods, depending on the available sample size. For larger quantities, a simple distillation setup is used, where the liquid is heated in a flask and the resulting vapor is condensed back into a liquid. The thermometer bulb must be precisely placed to measure the temperature of the vapor as it condenses, not the boiling liquid itself.
The temperature reading is taken once the vapor ring reaches the thermometer bulb and the temperature stabilizes, indicating that the pure substance’s vapor is in equilibrium with the external air pressure. When only a small amount of liquid is available, a micro-scale method, often involving a Thiele tube or similar heating apparatus, is preferred. This technique requires less than a milliliter of sample.
This micro-method involves placing a tiny inverted capillary tube into a small sample of the liquid, which is then heated slowly. As the temperature rises, the liquid’s vapor pressure eventually exceeds the atmospheric pressure, causing a stream of bubbles to escape from the capillary tube. The heat is then removed, and the boiling point is recorded at the moment the vapor pressure drops just below the external pressure, causing the liquid to be sucked back into the capillary tube.
The Impact of External Pressure and Purity
The measured boiling point is sensitive to the surrounding external pressure. Because boiling occurs when vapor pressure matches external pressure, a decrease in atmospheric pressure, such as at a higher altitude, lowers the temperature required for boiling. Chemists define the “normal boiling point” as the temperature measured specifically at a standard pressure of one atmosphere (101.3 kilopascals or 760 Torr).
When a boiling point is measured at a non-standard pressure, a correction must be applied to compare the result to literature values. A general rule is that for pressures near one atmosphere, a drop of \(10 \text{ Torr}\) in pressure results in approximately a \(0.5^\circ \text{C}\) decrease in the boiling point. This correction ensures the recorded value reflects the substance’s inherent properties.
The presence of non-volatile impurities, such as dissolved salts or sugars, significantly affects the measured temperature, leading to boiling point elevation. These impurities reduce the solvent’s vapor pressure, meaning a higher temperature is required for the solution’s vapor pressure to reach the external pressure. Consequently, an impure substance will boil at a higher temperature and over a range of temperatures, rather than at a single, sharp temperature.