Final answer:
When a liquid's vapor pressure equals the surrounding atmospheric pressure, it reaches its boiling point and begins to boil, undergoing a phase change from liquid to gas.
Step-by-step explanation:
Understanding Vapor Pressure and Boiling Point
When a liquid's vapor pressure becomes equal to the external atmospheric pressure, the liquid begins to boil. This factor is evident when observing how the temperature affects the vapor pressure of a liquid, as demonstrated in various figures showing the correlation. The boiling point of a substance is particularly important because it signifies the temperature where the liquid's equilibrium vapor pressure matches the pressure exerted by its surrounding atmosphere, usually measured in one atmosphere (1 atm) for the normal boiling point.
Within a closed system, a dynamic equilibrium is established between the rate of vaporization (liquid turning into gas) and condensation (gas turning back into liquid). This occurs as molecules possess varying kinetic energies, allowing them to enter and exit the vapor phase. The Clausius-Clapeyron equation further explains this nonlinear relationship between vapor pressure and temperature, allowing calculations for the enthalpy of vaporization.
When discussing liquids in an open container, boiling occurs when the vapor pressure inside the liquid forms bubbles that are capable of overcoming the external pressure, causing the bubbles to rise and release the vapor into the atmosphere. The point at which this process takes place defines the boiling point of the liquid under the specific atmospheric conditions. Therefore, it's crucial to understand that when a liquid reaches its boiling point, it undergoes a phase change from liquid to gas.