At what temperature does the vapor pressure of a liquid equal the external pressure?

The temperature at which the vapor pressure of a liquid equals the external pressure is known as the boiling temperature. At this point, the molecules in the liquid have enough energy to overcome the intermolecular forces holding them together and convert to the gaseous phase.

As temperature increases, the kinetic energy of the molecules also increases, which raises the vapor pressure of the liquid. When this vapor pressure reaches the level of the external pressure (such as atmospheric pressure), the liquid will start to boil, allowing bubbles of vapor to form within the liquid itself.

For example, under standard atmospheric pressure (1 atm), water boils at 100°C. If the external pressure were higher (like underwater or in a pressure cooker), the boiling temperature would also be higher, demonstrating the dependence of the boiling point on external pressure.

The other options refer to different physical states or transitions: the melting temperature is concerned with phase change from solid to liquid, the critical temperature is the maximum temperature at which a substance can exist as a liquid regardless of pressure, and the triple point temperature is the specific condition where all three phases (solid, liquid, and gas) coexist in equilibrium. None of these denote the specific condition where vapor pressure equals the external pressure leading to a boiling state.