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C4 - Phase Change

I can explain what is happening at the molecular level when a substance changes from one state to another


Liquids and solids differ from gases in that the particles (atoms, molecules, or ions) are much closer together, so the total volume of a liquid or solid is much closer to the sum of the volumes of the particles. The volume of a gas, as you may recall from the chapter on “The Behavior of Gases,” is related to the volume of the spaces between the particles and not to the volume of the particles.

At all temperatures above absolute zero, atoms and molecules are in constant random motion. The particles travel in a straight line unless they collide with another particle. In the absence of any attractive forces, this molecular motion would cause all substances to be in gaseous form. The fact that solid and liquid states exist tells us that there are forces that hold molecules and atoms together even when they are not chemically bonded. The forces of attraction that hold atoms and molecules together in solid and liquid phases are called intermolecular forces of attraction. These forces are different from chemical bonds, which are called intramolecular forces.

The phase (solid, liquid, or gas) of a substance is the result of a competition between the molecular motion that pushes the molecules apart and the attractive forces that pull them together. If the molecular motion is much greater than the attractive forces, the substance will be gaseous. If the molecular motion is nearly the same in strength as the attractive forces, the substance will be liquid, and if the molecular motion is much less than the attractive forces, the substance will be solid. When the molecular motion is increased or decreased by changing the temperature, the relationship between the molecular motion and the attractive forces changes, and the substance may change its phase.
The Structure of Solids and Liquids

The diagram below illustrates the molecular arrangements in the three phases of matter.

In the gaseous phase, the molecular motion dominates. The molecules of the substance are completely separated and move about independently of each other. The spaces between the molecules are very large compared to the size of the particles, so the measured volume of a gas is actually a measurement of the spaces between the molecules. In comparison, the molecular structure of the liquid phase has some spaces between the particles that allow the particles to move past one another, but the attraction between the particles is strong enough to prevent them from moving very far apart. In the solid phase, the forces of attraction have completely overcome molecular motion, and the movement of the particles has been reduced to vibrating in place. The particles cannot move past one another and are held in a tightly-packed pattern, so there is very little space between the particles.

For a brief animation showing the differences in molecular motion and relative position for gases, liquids, and solids (2d), see http://www.youtube.com/watch?v=s-KvoVzukHo (0:52).

For a video demonstration of a physical model of molecular motion and states of matter (2d), see http://www.youtube.com/watch?v=ynUso6rJ0rE (3:02).

Properties of Solids

The intermolecular forces of attraction in solids hold the particles so tightly in place that they cannot pull away from each other to expand their volume, nor can they flow past one another to change shape. Therefore, solids hold their own shape and volume regardless of their container. There is very little empty space in the solid structure, so solids are virtually incompressible. Since molecules cannot pass each other in the structure, diffusion or mixing is essentially non-existent beyond the surface layer.
Properties of Liquids

Attractive forces between molecules are a major factor in the behavior of the liquids. Since the particles in a liquid remain in touch with each other, liquids maintain their volume, but since the particles can flow past each other, liquids take the shape of their container. of liquid will be in any container, but because the liquid molecules are not held in a tightly-packed pattern like solids are, the molecules can move past one another, allowing the liquid to fit the shape of the container.

In gases, the distance between the molecules is so great that the size of the molecules themselves become inconsequential. A gas, then, is considered to be essentially empty space. If we consider the volumes of molecules of oxygen gas, , and molecules of Freon gas, , at STP, we know that the volumes will be the same, . This is because the sizes of the molecules themselves are negligible compared to the empty space in a gas. The fact that the Freon molecules are several times larger than the oxygen molecules makes no difference. In a gas, you are measuring the volume of the empty space. If we consider the volume of molecules of liquid oxygen and molecules of liquid Freon both at STP, we find the volume of the oxygen is about and the volume of the Freon is about . In the case of liquids, the volumes of the molecules themselves make a difference. In liquids, the volume of a group of molecules is related to the volume of the individual molecules, so equal moles of liquids do not occupy equal volumes under the same conditions. Since liquids have many more molecules in a smaller volume, they will have much greater densities than gases.

When a substance is compressed, it is not the molecules themselves that are compressed; it is the space between the molecules that is compressed. Gases have a great deal of empty space and are easily compressed. A pressure of compresses a gas to one-third its volume at . Liquids have very little space between molecules and do not compress easily – a pressure of will have virtually no effect on the volume of a liquid. Liquids are used in hydraulic systems because of their ability to flow to fit their container and their incompressibility.

Diffusion in gases (mixing) is nearly instantaneous. If you release a colored gas in a container of non-colored gas, the color spreads evenly throughout the container in a second or two. In liquids, diffusion is a much slower process. The dye molecules require much more time to move from one side of a container to the other due to the smaller spaces between molecules and the almost constant collisions with other molecules.
Lesson Summary
The molecular arrangement in solids is a highly organized, tightly-packed pattern with small spaces and molecular motion reduced to vibration in place.
Molecules in a solid maintain both their own shape and their own volume.
Solids are virtually incompressible and have little diffusion beyond the surface layer.
Molecules in the liquid phase have some freedom of movement but their motion is much more restricted than that of gases.
Liquids maintain their volume but take the shape of their container.
Liquids are only slightly compressible.
Diffusion in liquids occurs more slowly than in gases.

Evaporation and Condensation

The temperature in a beaker of water is a measure of the average kinetic energy of the molecules in the beaker. This does not mean that all the molecules in the beaker have the same amount of kinetic energy. Most of the molecules will be within a few degrees of the average, but a few molecules may be considerably hotter or colder than the average. The kinetic energy of the molecules in the breaker will have a distribution curve similar to a standard distribution curve for most naturally occurring phenomena. For most naturally occurring phenomena, most instances of the phenomena will occur near the average. Instances that occur further away from the average are increasingly rare, as seen in the image below.

In the case of a beaker of water, some of the molecules will have an average temperature below the boiling point, while some of the molecules will have a temperature above the boiling point (see figure below). The dashed yellow line is the average temperature of the molecules and would be the temperature shown on a thermometer inserted into the liquid. The red line represents the boiling point of water ( at pressure), and the area under the curve to the right of the red line represents the number of molecules that are above the boiling point. In order for a molecule above the boiling temperature to escape from the liquid, it must either be on the surface, or it must be adjacent to many other molecules that are above the boiling point so that the molecules can form a bubble and rise to the surface.

Water boils only when a sufficient number of adjacent molecules are above the boiling point and can form bubbles of gaseous water, as seen below. The process of molecules escaping from the surface of a liquid when the average temperature of the liquid is below the boiling point is called evaporation.

The phase change process is a little more complicated than just having the molecules reach the boiling point. Gaseous molecules have a force of attraction between them due to the separation between the molecules. Recall two oppositely charged objects that are separated have potential energy, and the amount of potential energy can be calculated by multiplying the force of attraction times the distance of separation. At the same temperature, the same molecules in the liquid state and the gaseous state do not have the same total energy. If they are at the same temperature, they have the same kinetic energy, but the gaseous molecules have additional potential energy that the liquid molecules do not have. As a result, molecules in the liquid state hot enough to exist in the gaseous state must absorb energy from their surroundings to gain the potential energy needed to change phase. This potential energy is called the heat of vaporization.

If a saucer of water is sitting out on the countertop, the water will slowly disappear – yet, at no time is the temperature of the water ever at the boiling point. When molecules of a liquid are evaporating, it is clear that it is the hottest molecules that are evaporating. It might seem that once the hottest molecules are gone, evaporation would no longer continue. This is not true, as the water in an open container continues to evaporate until it is all in the vapor state. When a substance is a vapor, the substance is in the gaseous phase even though the substance is at a temperature below its boiling point. Note that a substance in the gaseous phase at temperatures above the boiling point of its liquid is called a gas, not vapor. Evaporation continues because the temperature of the liquid is the average temperature of all the molecules. When the hottest molecules evaporate, the average temperature of those molecules left behind is lower, so the molecules left behind also contribute to the heat of vaporization to the evaporating molecules. The process of evaporation causes the remaining liquid to cool significantly. Heat flows from warmer objects to colder objects, so when the liquid cools due to evaporation, the surroundings will give heat to the liquid. The temperature of the liquid is raised so that it matches the temperature of the surroundings, thus producing more hot molecules. This process can continue in an open container until the liquid is all evaporated.

The rate of evaporation is related to the strength of the intermolecular forces of attraction, to the surface area of the liquid, and to the temperature of the liquid. As the temperature of a liquid gets closer to the boiling point, more of the molecules will have temperatures above the boiling point, resulting in faster evaporation. Substances with weak intermolecular forces of attraction evaporate more quickly than those with strong intermolecular forces of attraction. Substances that evaporate readily are called volatile, while those that hardly evaporate at all are called non-volatile.

Liquids in an open container will usually evaporate completely. What happens, however, if the container is closed? When a lid is placed over the container, the molecules that have evaporated are now kept in the space above the liquid. This makes it possible for a gaseous molecule to condense back to a liquid after colliding with another molecule or a wall, as seen in the figure below. This process where a gas or vapor is changed into a liquid is called condensation. Molecules at the boiling point can exist in either the liquid phase or gaseous phase – the only difference between them is the amount of potential energy they hold. See below figure.

For a liquid molecule with adequate temperature to exist in the gaseous phase, it needs to gain the heat of vaporization. It does this by colliding with adjacent molecules. For a gaseous molecule to return to the liquid phase, it must give up the same amount of potential energy that it gained. This amount of potential energy is called the heat of vaporization when it is being gained and the heat of condensation when it is being lost, but the amount of energy gained or lost is the same amount.

As more and more molecules evaporate in a closed container, the partial pressure of the gas in the space above the liquid increases. The rate at which the gas condenses is determined by the partial pressure of the gas, the surface area, and the substance involved. Once these factors are established, the rate of condensation will only vary depending on the partial pressure of the gas. As the partial pressure of the gas in the space above the liquid increases, the rate of condensation will increase.

It was pointed out that as a liquid evaporates, the remaining liquid cools because the hottest molecules are leaving, so the average temperature decreases and the heat of vaporization is absorbed from the remaining molecules. For similar reasons, when a gas is undergoing condensation, the temperature of the remaining gas increases because the coolest molecules are condensing, thus raising the average of those left behind. The condensing molecules must then give up the heat of condensation.