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# 13.2: Liquids

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## Lesson Objectives

• Describe a liquid according to the kinetic-molecular theory.
• Describe how a liquid exhibits surface tension.
• Describe the evaporation of a liquid and its relationship to the kinetic energy of the evaporating particles.
• Define vapor pressure and understand its relationship to intermolecular forces and to the temperature of the liquid.
• Describe the process of boiling and differentiate between boiling point and normal boiling point.
• Use a vapor pressure curve to determine boiling points at different atmospheric pressures.

## Lesson Vocabulary

• boiling point
• condensation
• evaporation
• fluid
• normal boiling point
• surface tension
• vapor pressure
• vaporization

### Recalling Prior Knowledge

• What is the kinetic-molecular theory?
• What makes gases different from liquids and solids?

Gases are easy to study because of the extremely large distances between the gas particles. Because there is so much space between particles, intermolecular forces can largely be ignored, which vastly simplifies any analysis of the motion exhibited by individual particles. In this lesson, we begin to study the properties of liquids and discover the importance of a liquid’s intermolecular attractive forces.

## Properties of Liquids

The primary difference between liquids and gases is that the particles of a liquid are much closer together, and there is very little empty space between them. Liquids are essentially not compressible and are far denser than gases. According to the kinetic-molecular theory for gases, any attractive forces between the particles of a gas are so minor that they can be ignored in most cases. For liquids, the intermolecular attractive forces are the only thing that keeps the particles close together. Liquids and solids are referred to as the condensed states of matter. One way in which liquids and gases are similar is that they are both fluids. A fluid is a substance that is capable of flowing from one place to another and takes the shape of its container.

### Surface Tension

Molecules within a liquid are pulled equally in all directions by intermolecular forces. However, molecules at the surface are pulled downwards and sideways by other liquid molecules, but not upwards away from the surface. The overall effect is that the surface molecules are pulled into the liquid, creating a surface that is tightened like a film (see Figure below (A)). The surface tension of a liquid is a measure of the elastic force in the liquid’s surface. Liquids that have strong intermolecular forces, like the hydrogen bonding in water, exhibit the greatest surface tension. Surface tension allows objects that are denser than water, such as the paper clip shown in Figure below (B), to nonetheless float on its surface. It is also responsible for the beading up of water droplets on a freshly waxed car, because there are no attractions between the polar water molecules and the nonpolar wax.

(A) Molecules at the surface of a liquid are pulled downwards into the liquid, creating a tightened surface. (B) Surface tension allows a paper clip to float on water’s surface.

### Evaporation

A puddle of water left undisturbed eventually disappears. The liquid molecules escape into the gas phase, becoming water vapor. Vaporization is the process in which a liquid is converted to a gas. Evaporation is the conversion of a liquid to its vapor form below the boiling temperature of the liquid. If the water is instead kept in a closed container, the water vapor molecules do not have a chance to escape into the surroundings and so the water level does not change. As some water molecules become vapor, an equal number of water vapor molecules condense back into the liquid state. Condensation is the change of state from a gas to a liquid.

In order for a molecule to escape into the gas state, it must have enough kinetic energy to overcome the intermolecular attractive forces in the liquid. Recall that a given liquid sample will have molecules with a wide range of kinetic energies. Liquid molecules with a kinetic energy that is above a certain threshold are able to escape the surface and become vapor. Because only the highest energy molecules are leaving the liquid state, the collection of molecules that remain in the liquid now have a lower average kinetic energy. Thus, as evaporation occurs, the temperature of the remaining liquid decreases. You have observed the effects of evaporative cooling. On a hot day, the water molecules in your perspiration absorb body heat and evaporate from the surface of your skin. The evaporating process leaves the remaining perspiration cooler, which in turn absorbs more heat from your body.

A given liquid will evaporate more quickly when it is heated. This is because the heating process results in a greater fraction of the liquid’s molecules having the necessary kinetic energy to escape the surface of the liquid. Figure below shows the kinetic energy distribution of liquid molecules at two temperatures. The number of molecules that have the required kinetic energy to evaporate are shown in the shaded area under the curve at the right. The higher temperature liquid (T2) has more molecules that are capable of escaping into the vapor phase than the lower temperature liquid (T1).

Kinetic energy distribution curves for a liquid at two temperatures, T1 and T2. The shaded area represents the molecules with enough kinetic energy to escape the liquid and become vapor.

## Vapor Pressure

When a partially filled container of liquid is sealed with a stopper, some liquid molecules at the surface evaporate into the vapor phase. However, the vapor molecules cannot escape from the container. Over time, some of the molecules lose energy through collisions with other molecules or with the walls of the container. At this point, the less energetic vapor molecules are trapped by the attractive forces of the molecules in the liquid, and they begin to condense back into the liquid form. Eventually, the system reaches a point where the rate of evaporation is equal to the rate of condensation (Figure below). This is called a dynamic equilibrium between the liquid and vapor phases.

A dynamic equilibrium can be illustrated by an equation with a double arrow, meaning that the reaction is occurring in both directions and at the same rate.

$\mathrm{H_2O}(l) \rightleftharpoons \mathrm{H_2O}(g)$

The forward direction represents the evaporation process, while the reverse direction represents the condensation process.

Because they cannot escape the container, the vapor molecules above the surface of the liquid exert a pressure on the walls of the container. The vapor pressure is a measure of the pressure exerted by the vapor that forms above its liquid form in a sealed container. Vapor pressure is considered a property of the liquid and is constant for a given substance at a set temperature. The vapor pressure of a substance at a given temperature is based on the strength of its intermolecular forces. A liquid with weak intermolecular forces evaporates more easily and has a high vapor pressure. A liquid with stronger intermolecular forces does not evaporate easily and thus has a lower vapor pressure. For example, diethyl ether is a nonpolar liquid with weak dispersion forces. Its vapor pressure at 20°C is 58.96 kPa. Water is a polar liquid whose molecules are attracted to one another by relatively strong hydrogen bonding. The vapor pressure of water at 20°C is only 2.33 kPa, far less than that of diethyl ether.

Vapor pressure can be measured by the use of a manometer (Figure below).

In the top picture, the flask of liquid ethanol has just been sealed and no vapor has accumulated, so the pressure inside is equal to the external atmospheric pressure. This is seen by the equal levels of the mercury in the U-tube. In the bottom picture, the system has been allowed to reach a dynamic equilibrium and the ethanol vapor is exerting a pressure equal to its vapor pressure. The vapor pressure can be measured by the height difference between the levels of mercury on each side of the U-tube.

### Vapor Pressure and Temperature

Vapor pressure is dependent upon temperature. When the liquid in a closed container is heated, more molecules escape the liquid phase and evaporate. The greater number of vapor molecules strike the container walls more frequently, resulting in an increase in pressure. Table below shows the relationship between temperature and vapor pressure for three different liquids.

Vapor Pressure (in kPa) of Three Liquids at Different Temperatures
0°C 20°C 40°C 60°C 80°C 100°C
Water 0.61 2.33 7.37 19.92 47.34 101.33
Ethanol 1.63 5.85 18.04 47.02 108.34 225.75
Diethyl ether 24.70 58.96 122.80 230.65 399.11 647.87

Notice that the temperature dependence of the vapor pressure is not linear. From 0°C to 80°C, the vapor pressure of water increases by 46.73 kPa, but it increases by another 53.99 kPa in a span of only twenty degrees from 80°C to 100°C.

Watch a simulation of vapor pressure at http://www.dlt.ncssm.edu/core/Chapter10-Intermolecular_Forces/Chapter10-Animations/VaporPressure.html.

You can participate in an online lab, the Dumas Molar Mass Lab, that determines the molar mass of a volatile liquid at http://www.youtube.com/watch?v=0UJXa9Hd88I&feature=player_embedded. The document that accompanies this lab is found at http://www.dlt.ncssm.edu/core/Chapter7-Gas_Laws/Chapter7-Labs/Dumas_Molar_Mass_web_01-02.doc.

### Boiling Point

As a liquid is heated, the average kinetic energy of its particles increases. The rate of evaporation increases as more and more molecules are able to escape the liquid’s surface into the vapor phase. Eventually, a point is reached when the molecules all throughout the liquid have enough kinetic energy to vaporize. At this point, the liquid begins to boil. The boiling point is the temperature at which the vapor pressure of a liquid is equal to the external pressure. Figure below illustrates the boiling of a liquid.

In the picture on the left, the liquid is below its boiling point, yet some of the liquid evaporates. On the right, the temperature has been increased until bubbles begin to form in the body of the liquid. When the vapor pressure inside the bubble is equal to the external atmospheric pressure, the bubbles rise to the surface of liquid and burst. The temperature at which this process occurs is the boiling point of the liquid.

The normal boiling point of a liquid is the temperature at which the vapor pressure of the liquid is equal to the standard atmospheric pressure of 760 mmHg. Because atmospheric pressure can change based on location, the boiling point of a liquid changes with the external pressure. The normal boiling point is a constant because it is defined relative to a standard pressure (760 mmHg, or 1 atm, or 101.3 kPa). The boiling points of various liquids can be illustrated by vapor pressure curves (Figure below). A vapor pressure curve is a graph of vapor pressure as a function of temperature. To find the normal boiling point of a liquid, a horizontal line is drawn from the left at a pressure equal to the standard pressure. The temperature at which that line intersects with the vapor pressure curve of a liquid is the boiling point of that liquid.

The boiling point of a liquid also correlates to the strength of its intermolecular forces. Recall that diethyl ether has relatively weak intermolecular forces, so the liquid has a relatively high vapor pressure at a given temperature. The weak forces also mean that it does not require a large an input of energy to make diethyl ether boil, so it has a relatively low normal boiling point of 34.6°C. Water, with its much stronger hydrogen bonding, has a lower vapor pressure than diethyl ether at any given temperature, and it has a higher normal boiling point of 100°C.

As stated earlier, boiling points are affected by external pressure. At higher altitudes, the atmospheric pressure is lower. With less pressure pushing down on the surface of the liquid, it boils at a lower temperature. This can also be seen from the vapor pressure curves. If one draws a horizontal line at a lower vapor pressure, it intersects each curve at a lower temperature. The boiling point of water is 100°C at sea level, where the atmospheric pressure is close to the standard value. In Denver, CO, at 1600 m above sea level, the atmospheric pressure is only about 640 mmHg, so water boils at about 95°C. On the summit of Mt. Everest, the atmospheric pressure is about 255 mmHg, so water boils at only 70°C. On the other hand, water boils at temperatures greater than 100°C if the external pressure is higher than the standard value. Pressure cookers do not allow the water vapor to escape, so the total pressure inside the cooker increases. Since water now boils at a temperature above 100°C, the food cooks more quickly.

The effect of decreased air pressure can be demonstrated by placing a beaker of water in a vacuum chamber. At a low enough pressure, about 20 mmHg, water will boil at room temperature!

Watch a video demonstration showing the relationship between vapor pressure and concentration at http://www.youtube.com/watch?v=PXYorrMu0Mw.

Watch a video demonstration of the effect of changing pressure on the boiling point of water at http://www.youtube.com/watch?v=Cshd5MVGpfk (1:23).

## Lesson Summary

• Intermolecular forces keep the particles of liquids close together, but due to relatively loose associations between any two particles, liquids are still able to flow freely.
• Surface tension is a property of liquids that makes it slightly more favorable for a liquid to minimize its surface area. This can lead to interesting properties, such as the ability of certain objects to float on the surface of a liquid even if they are denser than that liquid.
• Molecules of a liquid evaporate even when the temperature of the liquid is below its boiling point. Molecules with high kinetic energies escape the surface of the liquid, leaving the remaining liquid cooler than it was before.
• The vapor pressure of a liquid is determined by the strength of the intermolecular forces between its particles. Vapor pressure increases as temperature increases due to an increased rate of evaporation. The relationship between vapor pressure and temperature is shown on a vapor pressure curve.
• The boiling of a liquid occurs when its vapor pressure is equal to the external pressure.

## Lesson Review Questions

### Reviewing Concepts

1. Explain why gases are compressible, while liquids are almost completely incompressible.
2. Explain how a molecule of a liquid evaporates in terms of its kinetic energy. Why does the rate of evaporation increase with temperature?
3. When a highly volatile (evaporates rapidly) liquid such as acetone is spilled on your skin, it feels cold even if the liquid is originally at room temperature. Explain why this is true.
4. A liquid in a closed container has a constant vapor pressure. What can be said about the rate of evaporation of the liquid compared to the rate of condensation of the vapor?
5. Explain how the boiling point of a liquid varies with the atmospheric pressure.
6. What is the difference between boiling point and normal boiling point?
7. Liquid A has weaker intermolecular forces than liquid B.
1. Which liquid has the higher vapor pressure?
2. Which liquid has a higher normal boiling point?
3. Which liquid would demonstrate greater surface tension?

### Problems

1. Use the vapor pressure curve of water to answer the following questions.
1. What is the vapor pressure of water at 60°C?
2. What is water’s temperature when its vapor pressure is 100 mmHg?
3. At what temperature would water boil if the external pressure were 500 mmHg?
4. The atmospheric pressure is 250 mmHg. Is water a liquid or a gas at a temperature of 80°C?
2. A mountain climber in a camp at 3000 m above sea level decides to make pasta for dinner. The package states that the pasta should be cooked in boiling water for 10 minutes. How should the climber adjust this cooking time, if at all? Explain.

## Points to Consider

Solids are the other condensed phase of matter, with particles held closely together by attractive forces.

• What are some of the ways in which the particles of a solid are arranged?
• What happens when the temperature of a solid reaches its melting point?

Aug 02, 2012

Sep 09, 2014