Boiling and External Pressure

Introduction

A liquid can evaporate from its surface at many temperatures, but boiling involves vapour bubbles forming and surviving within the liquid. The boiling temperature is reached when the liquid's equilibrium vapour pressure equals the pressure imposed by its surroundings. Because that surrounding pressure can change, boiling point is not one immutable number independent of location or apparatus.

Core explanation

As a liquid warms, its vapour pressure rises. At a temperature where that vapour pressure matches external pressure, vapour bubbles can form in the bulk and avoid being crushed by the surroundings; the liquid boils. At one atmosphere external pressure, this temperature is called the normal boiling point. For pure water it is near 100 °C at 1 atm, but the exact boiling temperature depends on pressure and purity.

Lower external pressure requires a lower vapour pressure for boiling, so the liquid reaches its boiling condition at a lower temperature. This is why water boils at a lower temperature at high altitude, where atmospheric pressure is lower. Higher external pressure requires a higher vapour pressure, so boiling temperature rises. A pressure cooker uses this principle to allow liquid water and steam to reach temperatures above the ordinary open-pot boiling temperature, subject to its design and controls.

The condition is about equilibrium vapour pressure and surrounding pressure, not about temperature alone. A table value of “100 °C” is shorthand for a particular pressure. An exam problem might supply vapour-pressure data: find where the liquid's vapour-pressure curve crosses a horizontal line for the external pressure. If the intersection lies between tabulated points, interpolate or use an appropriate model as directed, while recognising that vapour pressure usually changes nonlinearly with temperature.

Evaporation occurs below boiling from the liquid surface. Some surface molecules have sufficient energy to escape even while vapour pressure is below external pressure and bulk bubbles cannot persist. A puddle drying at room temperature is therefore not “boiling very slowly.” Boiling is a bulk phase-change phenomenon under a pressure condition; evaporation can be a surface process over a much wider temperature range.

During boiling of a pure liquid at constant external pressure, adding heat can drive vaporisation while temperature remains nearly constant until the liquid is exhausted, assuming equilibrium and no significant superheating. Energy goes into phase change rather than raising temperature at that moment. Real systems can have temperature gradients, dissolved solutes or boiling delays, so the ideal plateau is an approximation.

Intermolecular forces affect boiling temperature through vapour pressure. Stronger attractions generally lower vapour pressure at a given T and require higher T to reach the same external pressure. But comparisons across substances also depend on molecular size and shape. A high boiling point does not imply a stronger covalent bond inside each molecule; the phase change mostly separates molecules from one another.

Pressure inside a bubble can include contributions from liquid depth and surface curvature, so the simple statement “vapour pressure equals atmospheric pressure” is an introductory bulk condition. For a deeper bubble, local liquid pressure can exceed atmospheric pressure. A precise fluid-mechanics treatment includes hydrostatic and surface-tension effects. In ordinary school problems, external pressure is the stated pressure applied to the liquid surface.

If a mixture is boiling, composition can change as more volatile components preferentially enter vapour. It may boil over a range rather than at one fixed temperature. The simple pure-liquid boiling point rule should therefore be applied to a pure substance or with mixture details supplied.

Step-by-step reasoning

1. Identify the pure liquid and the external pressure at its surface. 2. Find or compare its equilibrium vapour pressure as temperature changes. 3. Locate the temperature where vapour pressure equals external pressure. 4. Predict a lower boiling temperature for lower external pressure and vice versa. 5. Distinguish bulk boiling from surface evaporation and note mixture limitations.

Visual explanation

Plot vapour pressure on the vertical axis against temperature on the horizontal. Draw a rising curved vapour-pressure line and two horizontal external-pressure lines. The lower line intersects the curve at a lower temperature. Beside the graph draw a liquid with bubbles forming throughout at boiling, contrasting a few molecules escaping only from its surface during evaporation.

Real-world analogy

A bubble can expand only when its internal push is strong enough against the surrounding push. Lowering the surroundings' push lets expansion happen at a lower internal pressure. This analogy captures the pressure balance but not the full molecular phase-equilibrium physics.

Real-world example

Cooking water on a mountain can boil below its sea-level normal boiling temperature because lower atmospheric pressure needs a lower water vapour pressure for the boiling condition. Food preparation may take longer because the liquid water temperature is lower, though cooking time depends on the food and method as well.

Why?

Why does a pressure cooker raise boiling temperature? It raises the pressure surrounding the liquid. Water must then be heated until its vapour pressure rises to that higher pressure before sustained boiling occurs.

Common misconception

“A liquid cannot become vapour until its boiling point.” Evaporation can occur below boiling at the surface. The boiling point marks the condition for sustained vapour-bubble formation in the liquid under the surrounding pressure.

Worked example

A hypothetical liquid has vapour pressure 40 kPa at 60 °C, 60 kPa at 70 °C and 100 kPa at 90 °C. At 60 kPa external pressure, its boiling temperature from the supplied data is about 70 °C. At 100 kPa, it is about 90 °C. Lower external pressure gives lower boiling temperature, consistent with the graph. No calculation with PV = nRT for a fixed gas amount is needed; the relevant information is the liquid's equilibrium vapour-pressure curve.

Quick check

1. What happens to a pure liquid's boiling temperature when external pressure decreases? Answer: It decreases because the vapour-pressure curve meets the lower surrounding pressure at a lower temperature.

Exam focus

Write “vapour pressure equals external pressure” and include the pressure when quoting a boiling point. Distinguish boiling from evaporation and avoid claiming a fixed boiling temperature for every pressure or mixture.

Advanced insight

Boiling-point elevation by a nonvolatile solute arises because the solution has lower solvent vapour pressure than pure solvent at the same temperature. It must be heated more to reach the same external pressure. This connects liquid-state behavior with solution colligative properties, though concentration limits and nonideality matter for quantitative formulas.

Summary

Boiling occurs when a liquid's equilibrium vapour pressure reaches surrounding pressure. Lower external pressure lowers boiling temperature and higher pressure raises it. Evaporation can occur below boiling, and mixture or depth effects require additional detail beyond the simple pure-liquid rule.

Practice questions

1. Why does water boil at a lower temperature on a high mountain? Answer: Atmospheric pressure is lower, so water's vapour pressure reaches it at a lower temperature. 2. Can a puddle evaporate at 20 °C without boiling? Answer: Yes. Surface molecules can escape below the boiling point. 3. A pure liquid's vapour pressure reaches 80 kPa at 75 °C. What is its approximate boiling point under 80 kPa external pressure? Answer: About 75 °C, assuming the supplied equilibrium value is correct.