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General Chemistry--Unit 2

04/25/09

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Phase Transitions


�Before You Begin:

To master this material, you must be familiar with kinetic molecular theory and the states of matter: solid, liquid, and gas.

 


Changes in state are called phase transitions. Each of the phase transitions has a technical name and many have common names. The change from solid to liquid is fusion (or melting). The change from liquid to solid is solidification (or freezing). The change from liquid to gas is vaporization (or boiling). The change from gas to liquid is condensation. The change from solid to gas is sublimation. The change from gas to solid is deposition. Temperature and pressure changes cause substances to change state. We can use kinetic molecular theory to understand what happens on a molecular level when substances undergo phase transitions.

 

Phase Diagrams

Phase transitions occur with changes in temperature and pressure. The set of temperature and pressure data points at which phase transitions occur can be plotted in a graph to give concise phase information about a particular substance. These graphs are called phase diagrams. A generic phase diagram will have pressure on the y axis and temperature on the x axis. At low pressure and high temperature, a substance will be a gas. At high pressure and low temperature a substance will be a solid. At moderate temperatures and pressures a substance will be a liquid. The temperatures/pressures at which two states are in equilibrium (vaporization occurring at the same rate as condensation, for example) usually form a smooth curve on a phase diagram. Substances have a temperature/pressure at which all three phases can co-exist. This temperature/pressure data point is the triple point.

aphase diagram showing the phase changes

Note that the curve is more nearly vertical for the phase transitions between the solid and liquid states than the curves for the phase transitions involving the gaseous state. This is because a gas is compressible and changes in pressure radically affect a gas’s properties including phase changes. Solids and liquids, on the other hand, are not very compressible. There is much less distance between the particles that make up a solid or liquid compared to a gas.

 

Liquid and Gas Phase Transitions

Vaporization

The phase change from a liquid to a gas is known formally as vaporization and commonly as boiling, when the change is due to an increase in temperature. We can use kinetic molecular theory to visualize how this occurs on a molecular level. A liquid is composed of molecules in constant motion, but the molecules have strong enough intermolecular forces to keep them relatively close together. As the temperature of the liquid increases, the average kinetic energy of the molecules increases. Eventually, the molecules gain enough kinetic energy to escape the attractions of their neighbors, and the liquid boils.

The temperature at which the average kinetic energy is greater than the intermolecular attractions is known as the boiling point. As a liquid is heated, the temperature gradually rises until it reaches the boiling point. From that point the temperature stays the same until all of the liquid has vaporized. Any extra thermal energy goes into providing the energy needed for vaporization and as kinetic energy for the resulting gas molecules. This is why boiling water is 100 șC no matter how long you heat it, but steam can be very much hotter. The energy needed for the phase change is called the enthalpy of vaporization or heat of vaporization. This energy has to be supplied from the surroundings, so the vaporization process is endothermic.

The boiling point temperature varies with air pressure: the lower the air pressure is the lower the boiling temperature. Once again, kinetic molecular theory can give us insight as to why this happens. Air pressure pushes against the surface of the liquid. This force makes it more difficult for the liquid molecules to escape the intermolecular attractions. If the force drops, the liquid molecules need less kinetic energy to vaporize. That is why recipes and cooking instructions have to be adjusted for higher elevations. Food with high water content reaches the boiling point for that elevation and then it stops getting hotter until all of the water is vaporized. This temperature may not be hot enough for the food to cook properly. The boiling point of water at South Lake Tahoe, California (elevation 6200 feet) is about 95 șC instead of 100 șC. The boiling point temperature at sea level (or 1 atm) is called the normal boiling point for clarity.

The normal boiling point of a liquid depends upon the strength of its intermolecular attractions. If the attractions are low, the normal boiling point will be low. The liquid molecules need less kinetic energy to escape from the intermolecular attractions of neighboring molecules.

 


4Concept Check: Rank the halogens in order of increasing normal boiling point.

Answer: The rank is F2 (-188 șC) < Cl2 (-34 șC) < Br2 (59 șC) < I2 (457 șC). All of these molecules are non-polar, so the molecules are capable of induced dipole force only. The strength of induced dipole force increases as the size of the molecule increases, so the normal boiling point for these elements increases as you go down the family.  


 

Evaporation

Vaporization also occurs at temperatures below the boiling point. This process is commonly known as evaporation. Evaporation happens because, at any given temperature, some of the molecules have higher than average kinetic energy. Recall the kinetic energy distribution from kinetic molecular theory. Some of the molecules at the surface of the liquid have enough kinetic energy to escape the attractions of their neighbors. They leave the liquid phase and enter the gas phase. If the liquid is in a container open to the atmosphere, the vapor diffuses. This process continues until all of the liquid has evaporated and the resulting vapor has spread throughout the air. This is why a water puddle will dry up even though the temperature is far below the normal boiling point for water.

The evaporation process is able to continue even though the most energetic molecules evaporate. The average kinetic energy stays the same even though these higher than average energy molecules have left the liquid phase, so the kinetic energy distribution adjusts itself. Some molecules gain kinetic energy while others lose kinetic energy as the liquid reverts to the original kinetic energy distribution. Evaporation like boiling is an endothermic process. The boost of energy needed for the vaporization process and to replenish the kinetic energy of most energetic molecules comes from the thermal energy of the surroundings. This is how sweat cools the body.

Two factors influence the rate at which a liquid evaporates: temperature and volatility. When the temperature is high, the liquid molecules have greater kinetic energy, so a greater number of molecules have higher kinetic energy than the force of their intermolecular attractions. Volatility is the ease of which a liquid evaporates. If the volatility of a particular liquid is high, it evaporates readily. Molecules with weak intermolecular attractions evaporate the most easily.

 


4Concept Check: Which is more volatile, water or ammonia?

Answer: Ammonia, NH3, is more volatile than water, H2O. Both compounds have about the same molar mass, so their induced dipole forces are about the same strength. Both are polar molecules; they have tetrahedral electron geometries with non-bonding pairs of electrons giving them bent and pyramidal molecular geometries, respectively. Both are able to hydrogen bond, so the intermolecular forces of both are very strong. However, the H-O bond is more polar than the H-N bond, and water has two non-bonding electron pairs instead of one. Ammonia is less polar than water, so it is more volatile than water.

 


Condensation

The phase change from vapor to liquid is called condensation. Condensation occurs when vapor is cooled or when the pressure increases. We can use kinetic molecular theory to visualize how this occurs on a molecular level. As long as the kinetic energy is higher than the attractions due to intermolecular forces, vapor molecules behave as an ideal gas. This means that they have elastic collisions with one another and any surface with which they come in contact. At any given temperature, the kinetic energy of the vapor molecules is a distribution with some molecule having much lower than average kinetic energy and some having much higher kinetic energy. As the low energy molecules strike one another or a surface, they may not have enough kinetic energy to overcome intermolecular forces, so they stick together. If the temperature decreases, the number of molecules with kinetic energy high enough to overcome intermolecular attractions decreases, so the process of condensation happens faster than the process of evaporation.

Condensation will also occur if the pressure of the gas is high. When pressure is high, distance between molecules is low and the intermolecular attractions between particles are harder to overcome.

 

Critical Point and Supercritical Fluids

At very high temperatures, gas particles have so much kinetic energy that they cannot be condensed into a liquid. The highest temperature at which a gas and liquid can be in equilibrium is known as the critical temperature. The pressure needed to condense a gas into a liquid at the critical temperature is known as the critical pressure. The temperature and pressure data point on a phase diagram is the critical point. At temperatures above the critical point, the substance is in a new state known as a supercritical fluid. The properties of a supercritical fluid are an interesting blend of gas and liquid. A supercritical fluid has no surface tension and very low viscosity yet it has a density similar to that of a liquid and makes a good solvent. Supercritical carbon dioxide is an important cleaning agent in the semiconductor industry.

phase diagram showing supercritical fluid phase

 

Vapor Pressure

If a volatile liquid is open to the atmosphere, it will vaporize until none of the liquid is left. If a volatile liquid is in a closed container, only some of the liquid will vaporize. The vapor molecules have a kinetic energy distribution with some of the molecules lower in energy than the average. In a closed container, vapor molecules striking the surface of the liquid can become trapped by intermolecular attractions, if their kinetic energy happens to be low enough. The rate at which these molecules condense increases as the amount of vapor increases. Eventually an equilibrium in which a molecule condenses for every molecule that evaporates. The air in the container is said to be saturated with vapor. The pressure this vapor exerts is called vapor pressure. Like volatility, vapor pressure increases as temperature increases and as strength of intermolecular attraction decreases.

 

Solid and Liquid Phase Transitions

Fusion

Fusion, or melting, is the phase change from solid to liquid. Melting occurs when the temperature increases. It is less sensitive to pressure change than vaporization and condensation because solids and liquids are less compressible than gases. We can use kinetic molecular theory to understand what happens on a molecular level when a solid melts. The molecules, ions, or atoms that make up a solid are held relatively close to one another (compared to a gas) by strong intermolecular attractions or bonds. Although the particles move due to their kinetic energies, the relative movement is quite small. This gives a solid the property of having a definite shape. When a solid is heated, the particles gain kinetic energy. Eventually, some of the most energetic particles have enough kinetic energy to leave the vicinity of their neighbors in the ordered local structure. When the particles gain enough kinetic energy to overcome the attractions that hold them in an ordered structure but not enough kinetic energy to spread throughout the container, the substance melts.

The amount of thermal energy needed to melt a substance depends on the strength of the forces giving that substance its shape. Molecular substances are held together by relatively weak intermolecular forces, and they melt at low temperatures. Metallic substances are held together by metallic bonds. Metallic bonds are much stronger than intermolecular forces and the atoms are much closer together than the molecules in a molecular solid. However, the fluid nature of the valence electrons in a metal allow the nuclei and core electrons to slip past one another without breaking the metallic bonds. This is why metals are malleable and ductile. Since the atoms can move without disrupting bonds, metals have a wide range of melting points. Mercury is a liquid at room temperature, and gallium melts at slightly below body temperature. In contrast, metals in the lower center of the transition element region have melting points over 3000 șC. Ionic compounds are held together in crystalline arrays by ionic bonds. Ionic bonds are quite strong, and the proximity of opposite charges in a three dimensional lattice make it difficult for ions to move out of position. Ionic solids have high melting points. Network covalent solids are held in place by covalent bonds rather than intermolecular forces. These are quite strong, so melting points are very high for most network covalent solids. Carbon melts at 3527 șC, for example.

As with vaporization, melting is an endothermic process (requires an input of thermal energy or heat). When a solid is heated, the temperature increases until it reaches the melting point. The temperature remains at the melting point until the entire sample has liquefied. Any additional thermal energy supplies the energy needed for the phase change. This thermal energy is called the enthalpy of fusion or the heat of fusion. For a given substance, the heat of fusion is generally less than heat of vaporization. This is because the thermal energy is used to allow the molecules to move with respect to each other rather than to separate them completely.

Changes in pressure at constant temperature can cause a substance to either melt or freeze. This is very different from the behavior of a gas, since an increase in pressure will causes a gas to condense but never cause a liquid to boil.  An increase in pressure will cause a less dense phase to change into a denser phase. For most substances, the solid is denser than the liquid, and a pressure increase will cause a liquid to freeze. In most substances, the solid state is denser than the liquid state because the lower kinetic energy prevents the particles from escaping the intermolecular attractions or bonds. The molecules, ions, or atoms are packed tighter together with more mass per unit volume. Increasing the pressure on a liquid will force particles closer to one another, making the attractions stronger and favoring the solid state. An example of a substance that has a less dense liquid phase is candle wax. Substances of this type shrink when they freeze, and the liquid will float on top of the solid phase.   

 

phase diagram for a substance with a more dense solid phase than liquid phase (like wax)

 

Water is a very unusual substance. The solid phase is less dense than the liquid phase, so ice floats in liquid water and water expands when it freezes. Water molecules have strong hydrogen bonds and a three dimensional geometry that allows them to form extended networks that are especially stable. When water freezes, the molecules arrange themselves in a hexagonal lattice which maximizes this stability. This hexagonal structure is the reason that snowflakes have six sides. The lattice is also less dense than the random arrangement in liquid water. Increasing the pressure on ice causes it to melt.

In this discussion, we are talking about ‘normal’ ice. Water has 18 different solid phases! Most only occur at very high pressures. For detailed phase diagrams of water and molecular models of its ices, see Martin Chaplin’s web site about water. This site has lots of really interesting information that show how special (magical?) water really is.

 

phase diagram for a substance with a less dense solid phase than liquid phase (like water)

Solidification

Solidification, or freezing, is the phase change from liquid to solid. Freezing occurs when the temperature decreases. It is less influenced by pressure changes than vaporization or condensation because solids and liquids are less compressible than gases. We can use kinetic molecular theory to understand what happens on a molecular level when a liquid freezes. In the liquid state, particles are kept in close proximity by intermolecular forces, but they have enough kinetic energy to move with respect to one another. As the temperature drops, fewer particles have enough kinetic energy to overcome neighboring attractions, and the particles lock into place. As more and more of the particles settle into a relatively fixed position, the substance gains a shape of its own and becomes a solid.

As with condensation, freezing is an exothermic process (releases thermal energy or heat). As a liquid is cooled, it gradually decreases in temperature until it reaches the melting point/freezing point. The sample will stay at this temperature until all of it is a solid. If a liquid is very carefully and quickly cooled, the temperature drops more rapidly than the solidification process. This results in a supercooled liquid, one that is colder than its freezing point yet still in the liquid state. Because they are unstable, supercooled liquids will crystallize rapidly upon the slightest disturbance, such as stirring the liquid or adding a seed crystal.

 

cooling curve graph--temperature is on the y axis and time is on the x axis                       

Solid and Gas Phase Transitions

The phase change from solid into gas is known as sublimation. Substances undergo this phase transition at low temperatures and pressures. As the temperature increases, the kinetic energy of the particles increases enough to overcome intermolecular attractions. If the pressure is relatively low, molecules with enough kinetic energy to escape the rigid structure of the solid phase have enough kinetic energy to disperse as a gas rather than transform into the closely packed yet disorganized liquid phase. Like fusion, sublimation is an endothermic process (requires thermal energy input).

At atmospheric pressure, carbon dioxide solid (dry ice) sublimes rather than melts. The attractions that hold carbon dioxide solid together are induced dipole forces, which are quite weak. At room temperature and pressure, carbon dioxide is a gas (and a pretty darn ideal one, at that). In order to form dry ice, carbon dioxide is cooled and compressed until liquid. Then the pressure is suddenly decreased by allowing the liquid to expand in a spray. This has a cooling affect, just like an aerosol spray deodorant cools your armpit. The temperature drop causes some of the liquid to freeze in carbon dioxide snowflakes. These are scraped together and compacted to make blocks. For more info see the Dry Ice web site.

The phase change from gas into solid is known as deposition. Substances undergo this phase transition at low pressures. In a process called physical vapor deposition, thin films are grown on semiconductors and catalysts by bombarding a metal crystal surface under high vacuum with a stream of high energy particles.

 

 

 

 

 

 

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