How is it possible to have boiling water at a temperature other than \(100^{\circ} \mathrm{C} ?\)

The boiling point of a substance is typically thought of as a fixed property—water boils at 100°C, right? Not quite. The boiling point is the specific temperature at which a liquid converts into vapor. At this temperature, the vapor pressure of the liquid equals the pressure exerted on the liquid by its surrounding environment, allowing bubbles of vapor to form within the liquid.

However, boiling points are not static; they change in response to the environment. This flexible characteristic of boiling points is integral to understanding how we can manipulate them, which is vital in a variety of applications ranging from culinary to chemical processes. For example, when cooking at high altitudes, adjustments must be made due to the variation in boiling points.

Pressure is a game-changer when it comes to boiling points. When we talk about boiling water on a mountain versus sea level, we're looking at the influence of atmospheric pressure. The lower atmospheric pressure at higher altitudes means that water can transition to vapor at temperatures lower than the standard boiling point.

For instance, at the top of Mount Everest, water boils at about 69°C! Conversely, in a pressure cooker, the pressure is higher, making it harder for water molecules to escape as vapor. This increased pressure leads to a higher boiling point. This principle is also what enables autoclaves to sterilize equipment, by using pressure to reach higher temperatures that can kill bacteria and spores effectively.

Colligative properties are fascinating because they depend on the number of particles in a solution and not on what those particles are. Boiling point elevation, freezing point depression, vapor pressure lowering, and osmotic pressure are the four main colligative properties.

Boiling point elevation, for example, occurs when a non-volatile solute is dissolved in a solvent, like salt in water, and raises the boiling point of the solution. The solute particles interfere with the molecules of the solvent trying to evaporate, requiring more heat and therefore a higher temperature to reach the boiling point. This principle is not limited to cooking—it's also applied in systems like automotive engines, where antifreeze is used to modify the boiling point of the coolant.

Vapor pressure is all about escape. It's the measure of how much the molecules of a liquid are trying to spread out and form vapor. Vapor pressure increases with temperature, as molecules move more energetically and are thus more likely to enter the gas phase.

A liquid boils when its vapor pressure equals the external pressure. That's why changes in atmospheric conditions or the addition of a solute can alter a liquid's boiling point. If you reduce the vapor pressure—by introducing a solute—the system requires a higher temperature to get the vapor pressure back to the point where it equals the atmospheric pressure, which is the boiling point. Understanding vapor pressure is vital in fields from meteorology to cooking, and even in the design of pressure-relieving systems.