3 5 PHASE CHANGES PHASE CHANGES Name(s) The activities presented here focus on the energy changes that occur in substances undergoing a phase change. The first activity will take the most time to complete. You may want to work on the other activities while you take data for the first one. Activity #1 Phase Changes in Water Fill a small beaker (about 250 ml) about half full with ice cubes. Place a thermometer in the beaker, and record the temperature of the ice in the table below. Next, place the beaker on a hot plate. Stir the contents of the beaker well every minute (Do Not use the thermometer as a stirring device). Every two minutes, remove the beaker from the hotplate, stir thoroughly, and record its temperature. (Be sure to stir well before taking each temperature measurement.) Record your results below, along with a verbal description of the water (responding with either ice cubes only ; ice cubes and liquid water ; liquid water only ; or boiling liquid water ). Continue taking data until half of the water is boiled away. While you take this data you may work on the activities that follow this one. Time Temperature Observations UNIT 3 ENERGY AND SYSTEMS III-35 CALVIN COLLEGE
PHASE CHANGES Time Temperature Observations Questions to be answered after completion of the observations for Activity #1: 1. The entire time that the beaker is on the hot plate, heat is being transferred to the material in the beaker. Did this energy always manifest itself as an increase in the temperature of the water? (If there were times when the energy was not causing the temperature of the water to rise, describe what was happening to the water, and explain what was happening to the energy being given to the water.) 2. What inferences can you draw about the melting and evaporation processes and whether they require or give off heat? Activity #2: Salt and Ice A) Fill a 500 ml beaker with ice, and place a thermometer into the beaker. Record the temperature. UNIT 3 ENERGY AND SYSTEMS III-36 CALVIN COLLEGE
3 5 PHASE CHANGES B) Pour a few teaspoons of salt onto the ice, and watch the ice and thermometer carefully. You can continue to add salt, and you can stir the ice cubes as you think appropriate. Record the lowest temperature that the ice and salt mixture reaches. Do not dispose of your beaker or its contents, as you will need them in Activity #4. 1. Why is salt put on icy roads? 2. Does melting absorb thermal energy (make things colder) or release thermal energy (make thinks warmer}? Why? Activity #3: Evaporation Place a little rubbing alcohol onto a piece of cotton, and then swab the back of your hand. Describe your observations. (Does your hand stay moist for very long?) 1. Why did the back of your hand get cool? 2. Describe how this system of alcohol on the warm hand is analogous to the beaker of water on the hot plate. UNIT 3 ENERGY AND SYSTEMS III-37 CALVIN COLLEGE
PHASE CHANGES Activity #4 Super-cooling a Liquid You might be accustomed to thinking that all water below 0 C (32 F) will be ice, but in fact, it is possible to super cool pure water and have it remain liquid even below its normal freezing temperature. A) Check the temperature of your beaker of salt water. Be sure its temperature is not above -7 C (20 F). B) Get a clean, large test tube. Put just enough cold tap water into it that you can rest the test tube in your beaker and have the same water level in both the test tube and in the beaker. C) Let the test tube set undisturbed in the beaker of salt and ice for 5 minutes. During this time you can find a clean ice cube, something to smash it with, and some tweezers. D) At the end of the 5 minutes, smash the ice cube and pick up a small crystal of ice with the tweezers. Gently lift the test tube from the beaker (hopefully the water in it has not frozen) and carefully observe what happens when you drop the small ice crystal into the test tube of water. Record your observations: 1. If you can, measure the temperature of the ice you have created in your test tube. Is it the same temperature as the surrounding salt-water bath? Why or why not? You may wish to repeat this activity several times. You might try it with a thermometer inside the test tube. (Does it still work?) E) Water is not the only liquid that can be super-cooled. The heat packs that you may have used when you began this unit on energy contain a super-cooled solution of sodium acetate. Activate another heat pack, if available. 2. Explain why the heat pack gets hot when the solution solidifies. drop the small ice crystal into the test tube of water. Record your observations: UNIT 3 ENERGY AND SYSTEMS III-38 CALVIN COLLEGE
Phase Changes Content Overview When evaporation or boiling takes place, a substance goes from the liquid phase to the gas phase. During this process, the liquid particles with the highest kinetic energies are the ones most likely to change into a gas, because they are the ones with enough energy to escape from the other particles. (It is the mutual attraction of the particles that causes the substance to be a liquid, and not a gas, in the first place.) Since the particles with the highest energies are escaping, the particles left behind have a lower average kinetic energy than the original liquid had. Therefore, the liquid left behind is cooler than the original sample. This explains why some people call evaporation a cooling process. Surely you have had the opportunity to feel the cooling associated with evaporation after you emerged from a shower or from a swimming pool. The particles that are escaping during evaporation also decrease in temperatures. Originally, they were the particles with the highest kinetic energy in the liquid. However, they lose much of this kinetic energy as they overcome the attractive forces from the other particles in the liquid that are trying to hold them down. In order to represent this process with a model, you might imagine a liquid as a collection of marbles rolling about in an indentation in the ground. (See the diagram below.) The sloped sides of the hole represent the attractive forces of the particles. This is the force that holds the liquid together. Only the highest-energy particles are able to roll all the way up the side of the hole and escape (evaporate). Once the particles have escaped, their speeds will be less than when they were in the hole, and thus their temperature will be lower than it was before they evaporated. As these particles evaporate, they become part of the gas making up the surrounding air. The new gas particles have lost much of their kinetic energy by escaping the liquid. Therefore, they are colder than the surrounding gas. As the original gas particles bump into the evaporated particles, the original gas decreases in temperature as it transfers some of its energy to the evaporated particles. The net effect is that the temperature of the air decreases. You may have experienced this effect while sitting under a tree on a hot, sunny day. Water in the leaves of the tree evaporates and moves into the surrounding air. Since the temperature of this water will be lower than that of the surrounding air, the total air temperature in the area will decrease. It is worth considering the evaporation process from the perspective of conservation of energy. Everything is cooling off during evaporation--both the liquid left behind and the gas surrounding the liquid. Does this mean that energy is lost? Certainly not. It is merely transformed from one form to another. Think again of the marbles in the hole. The marbles that escape lose kinetic energy (measured as temperature) by rolling up the side of the bowl, but once they have escaped, they are moving around on a higher surface, and thus they have more potential energy. Since temperature only measures average kinetic energy per particle, not total energy, we observe a decrease in temperature. However, if a marble were to roll back down the incline and return to the bowl, it would recover its kinetic energy, and the temperature would go up. This potential energy, which cannot be measured with a thermometer, is referred to as latent heat. UNIT 3 ENERGY AND SYSTEMS III-39 CALVIN COLLEGE
PHASE CHANGES In summary, then, evaporation is a cooling process in which energy is converted from the kinetic energy of the particles to stored potential energy. In the reverse process of condensation, energy is converted from stored energy to kinetic energy. The temperature of the liquid and the surrounding gas will tend to increase during condensation as the particles release their latent heat. Sometimes it becomes confusing to talk about evaporation as a cooling process. Evaporation of the alcohol from your hand clearly cooled your hand, but think about the beaker of water on the hot plate. Here we stimulated the evaporation by heating the water with the hot plate. (And, if you think about the alcohol on your hand, you stimulated it to evaporate by providing it with heat from your hand.) Since heating stimulates evaporation, we can also associate evaporation with heating. Perhaps the best way to describe this curious situation in which evaporation is stimulated by heating, but is a cooling process, is just to say that evaporation absorbs thermal energy, converting it into potential energy. Similarly, the reverse process of condensation releases thermal energy as potential energy is converted into thermal energy. The processes of melting and freezing are analogous to evaporation and condensation in terms of the temperature changes that result. When a solid melts, the particles must escape from the rigid attraction of neighboring particles. During this process, they lose some of their kinetic energy. The result is that the temperature of the solid and the liquid mixture will decrease as melting occurs. This phenomena can be observed when salt causes ice to melt. (Note: even though a chemical reaction occurs when calcium chloride mixes with water, causing the temperature of the mixture to increase [see the Temperature and Thermal Energy Content Overview] the net effect of adding calcium chloride to ice is that the temperature will go down, because the decrease in temperature due to melting more than offsets the temperature increase from the chemical reaction.) The opposite is true during freezing. The liquid particles lose their energy of freedom as they become bound to neighboring particles. The result is an increase in kinetic energy of the particles, with an associated increase in temperature. This phenomena can be observed when a super cooled solution solidifies. UNIT 3 ENERGY AND SYSTEMS III-40 CALVIN COLLEGE
GAS (Particles are free) boil or evaporate condense LIQUID (Particles are partially free) melt freeze SOLID (Particles are stuck) Consider the processes on the left hand side of the diagram. The particles need to find some energy in order to get free. If they use their own thermal energy in order to get free, they will become "colder" because they will not be moving as fast as they used to be. Consider the processes on the right hand side of the diagram. They are the opposite of the processes on the left. The particles will tend to gain thermal energy as they get trapped. Their "energy of freedom" is turned into thermal energy. UNIT 3 ENERGY AND SYSTEMS III-41 CALVIN COLLEGE
PHASE CHANGES UNIT 3 ENERGY AND SYSTEMS III-42 CALVIN COLLEGE