The nature of heat


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The nature of heat

  1. 1. THE NATURE of HEAT As we know, heat is a form of energy. In the form of infrared radiation, heat from the sun travels through space at the speed of 186,000 miles per second. Upon arriving on earth, much of the radiant heat is absorbed by different kinds of matter and is converted into heat that we can feel (sensible heat). When you sit in the sun for a period of time on a clear spring day, you may find that your clothing and other objects around you become warm. Similarly, when you walk barefoot across a beach on a summer day, you may find the sand so hot that it burns your feet. In both cases, radiant heat from the sun has been absorbed by matter, and has been converted into heat that you can feel. In this chapter, we will study the effect of heat upon matter. HEAT AND THE MOTION OF MOLECULES Have you ever tried to drill a hole through a piece of metal? Both the drill and the metal become very hot. Around 1800, an English scientist named Count Rumford noted that, when a drill was used to bore a cannon, the bit of the drill and the cannon both became very hot. To keep the metals cool, he placed a cylinder of water around the end of the cannon. As the boring continued, the water became warmer and eventually boiled. Since the bit of the drill and the cannon were cold at the start, Rumford concluded that the heat produced probably came from the friction created by the particles of the metal of the bit rubbing against the particles of the metal of the cannon.
  2. 2. Further, he theorized that the motions of the particles in the metals themselves (atoms or molecules) generated the heat. Recall that various forms of energy can be converted into other forms: When electrical energy passes through a thin wire, as in a toaster, the wire becomes hot (electrical energy to heat energy). When you rub your hands together, heat is produced from the friction between the rubbed surfaces (mechanical energy to heat energy). In general, when any form of energy is absorbed by matter, the energy is changed to heat. This may be explained by the kinetic- molecular theory: The energy excites the molecules in the matter, causing them to move faster and to collide more frequently. As more collisions take place, more heat is; produced. The effect of heat energy on the motion of molecules can demonstrated by using a sealed tube containing a little mercury with some glass beads floating on the surface of the mercury. At ordinary temperatures, the glass beads merely float on the surface of the mercury. However, when the tube of mercury is heated, the glass beads, bounce up and down in a violent but random fashion. As still more heat is supplied to the sealed tube, particles of mercury begin to move more swiftly. The glass beads are repeatedly struck by many mercury particles at the same time; consequently, the glass beads begin to move randomly themselves. Thus, Rumford's theory that heat is related to the motions of molecules appears to be correct. According to the kinetic-molecular theory, heat energy acquired by a body is transformed into increased kinetic
  3. 3. energy of the molecules of the body. We observe this increased kinetic energy whenever a solid, a liquid, or a gas expands on heating. A further increase in kinetic energy will eventually cause the particles of a solid or liquid to become a gas. Recall that when an ice cube (a solid) is heated, it melts and becomes liquid water. When the water is heated, it vaporizes and becomes gaseous water. According to the kinetic-molecular theory, as increasing amounts of heat are supplied to a piece of ice, the water molecules move more rapidly until they gain sufficient energy to overcome the attractive forces holding them together. This permits the ice to liquefy and become water. Similarly, as still more energy is received, the water molecules move at even greater speeds. The attractive forces in the liquid are weakened and the water is converted into gaseous water. EXPANSION OF SOLIDS Your laboratory experience with the ball and ring apparatus indicated the effect of heat on volume. The increase in size is not due to an increase in the size of the particles that make up the solid ball, but rather to an increase in the average distance between the particles. When an object is heated, its particles vibrate faster, collide more violently, and consequently move farther apart, thereby increasing the volume of the object. When the object is cooled, the opposite change occurs and the volume of the object decreases. This decrease in volume is calledcontraction.
  4. 4. The expansion of solids by heating may cause serious practical problems. For example, the expansion of railroad tracks, bridges, or the concrete in a roadbed can create dangerous situations. Thus, allowance for the expansion of solids daring hot weather must he made in the construction of rails, bridges, and roads. For example, when rails arc laid, gaps between the ends of the rails provide for expansion. If this were not done, consider what would happen to the railroad tracks on a very hot day. The metal would expand, making the tracks bend andbuckle, which might cause an oncoming train to be derailed. In bridge construction, expansion joints allow for changes in the length of the bridge. Concrete roadbeds are built with spaces between the sections of concrete to allow for expansion. The contraction of solids, by cooling, may also present problems. Thus telephone and electrical wires are strung loosely to prevent their snapping as contraction takes place during the colder times of the year. UNEQUAL EXPANSION OF SOLIDS Through extensive studies, scientists have found that different metals expand at different rates when they are heated. For example, when a piece of iron and a piece of aluminum of equal size are heated together, we find that the aluminum expands more than twice as much as the iron. When two strips of different metals are fastened together, they form a compound bar, or bimetallic strip, which is employed in, useful devices such as thermostats and metallic thermometers. In these devices the two different
  5. 5. metals, usually brass and steel, are welded together. When the bar is heated, it bends because the brass expands more than the iron and becomes longer than the iron. Thus, the brass strip will be on the outside of the bend. As it cools, the bar returns to its original shape. The thermostat is a device containing a compound liar that regulates the heating systems of our homes. When the temperature in the house falls below the setting on the thermostat, the compound bar, which contracts as it cools, closes the circuit, turning on the heat. As the room is warmed, the compound bar in the thermostat expands, bends, and thereby breaks the circuit, shutting off the heat. The bimetallic thermometer is often used as an oven thermometer to indicate the temperature within an oven, or within a piece of meat that is cooking. EXPANSION OF LIQUIDS Liquids, like solids, expand when heated. In the laboratory experience we demonstrated that when water is heated, it expands. When the same water is cooled to its original temperature, the water contracts to its original volume. Many other liquids, such as alcohol and mercury, behave in the same way. At lower temperatures, however, the behavior of water is an exception to this rule. As water is cooled from 100° C to 4° C, it contracts-like other liquids do. However, when water is cooled below 4° C, the water expands-unlike other liquids. Water continues to expand until it reaches 0° C, its freezing point. It has been found, as
  6. 6. shown in that the spaces between the water molecules in ice are larger than the spaces between the water molecules in liquid water. Ice is therefore said to have an open structure. Thus, as ice is formed, the need for increased space between the molecules causes the volume of the ice to be greater than that of the water from which it was formed. (This expansion in volume begins as liquid water is cooled below 4° C.) Since the volume of ice is greater than the volume of water from which the ice is formed, the density of ice is less than the density of water. (Recall that density equals weight divided by volume.) This is why ice floats on water. Like solids, different liquids expand at different rates. As we will see later, the expansion of liquids is used in alcohol and mercury thermometers. The expansion of liquids must be considered in certain heating systems. In a hot-water heating system, allowances must be made for the expansion of heated water. As the furnace heats the water in the heating system, the water expands. If the expansion continues, pressure would build up in the pipes and could damage the entire system. To avoid this difficulty, an expansion tank is provided. Excess heated water enters the expansion tank and thereby reduces the pressure in the system. EXPANSION OF GASES Cases, like solids and liquids, expand when heated. Our laboratory experience indicated that, as air is warmed, it expands. Scientists have made similar observations with
  7. 7. other gases which indicate that gases confined in an elastic container expand when they are heated and contract when they are cooled. The expansion of gases by heat must be considered by automobile tire. manufacturers, since tires may burst if allowed to remain in the sun indefinitely. A less serious hazard caused by expanding gases is that bottles of soda may crack or even explode if they are ex posed to heat for a considerable length of time. Products such as whipped cream, shaving cream, deodorants, and insect repellents are now supplied in aerosol cans. These cans contain the product itself and a gas that forces the product out of the can when the valve, is open. When the product is used up, the can still contains unused gas. If this can is thrown into incinerator, the gas becomes heated, expands, and may cause the can to explode. Such aerosol cans should be discarded in a manner that does not involve heating. Different solids and liquids expand at different rates when heated. Gases, however, generally expand at the same rate when heated to the same temperature, at a given pressure. TEMPERATURE Heat and temperature are two terms that are often confused. We know that the temperature of a small sample
  8. 8. of molten iron is considerably higher that the temperature of the water in the ocean. However, the total heat in a sample of molten iron is much less than the total heat of the water in the ocean. Scientist now accept Rumford’s theory that heat is related to the motions of particles in matter. Thus, heat depends on the total kinetic energy of the particles in a body. Recall the equation relating kinetic energy with mass and velocity: K.E. = ½ m v2. Thus, the total kinetic energy of the particles in a body depends on the number of particles (mass) and the velocity of these particles. Because the water in the ocean is colder than the sample of molten iron, the velocity of the particles in the water is less than the velocity of the particles in the molten iron. However, the much larger quantity (mass0 of water compensates for the smaller velocity of the particles and thus the particles of water in the ocean possess greater kinetic energy. This means that there is more heat in the water in the ocean than in a small sample of molten iron. But why is the temperature of molten iron higher? Temperature, unlike heat, depends on the average kinetic energy of the particles, that is, the kinetic energy per particle. To find this average, we divide the total kinetic energy by the number of particles. Thus, the large mass of ocean water has a smaller average kinetic energy per particle and consequently has a lower temperature than a small sample of molten iron.
  9. 9. MEASURING TEMPERATURE Instruments designed to measure temperature are called thermometers. Most thermometers are based on the principle that matter, on heating, expands and, on cooling, contracts. In general, matter expands and contacts regularly. This means that the amount of expansion or contraction in length are generally equal for the same increase or decrease in temperature. This regular expansion and contraction has made it possible to construct three different types of thermometers: gas (air), liquid (mercury and alcohol), and solid (bimetallic) thermometers. The gas (air) thermometer In this thermometer, the glass bulb contains air. When the bulb is warmed, the air in the bulb expands and forces some of the colored water out of the tube. This changes the level of the liquid in the tube. By placing a suitable scale alongside the tube, temperature changes can be measured. Air thermometers of this type, while interesting, are not very accurate because the volume of a gas is also influenced by the air pressure around it. (Note that the flask contains a tube open at both ends. Why?). LIQUID THERMOMETERS Thermometers containing liquids such as mercury and alcohol are useful and accurate because these liquids usually expand and contract uniformly (regularly).
  10. 10. Mercury thermometers are made by filling a thin glass tube with mercury at a temperature greater than the maximum to be measured. The tube is then cut and sealed at the top. When the mercury cools, it contracts, leaving a partial vacuum above the mercury. (Liquids expand and contract to a much greater extent than do solids. Thus, in the given temperature range, the glass tube is scarcely affected by the temperature change.) The partial vacuum eliminates the effect of air resistance to the expansion of the mercury. The scale of the. thermometer is generally fixed by locating the boiling and freezing points of water on the scale. The distance between the boiling and freezing points is then divided into units depending on the temperature scale used. This will be discussed in the next section. Since mercury freezes at -39° C, it cannot be used to measure very low temperatures. However, mercury boils at 357° C, which means that a mercury thermometer can be used to measure temperatures above the boiling point of water. On the other hand, alcohol freezes at -114° C. Accordingly, alcohol thermometers are used to measure low temperatures. However, since alcohol boils at 78° C, alcohol thermometers cannot be used to measure high temperatures. SOLID (BIMETALLIC) Recall that a bimetallic strip behaves as it does because different metals expand at different rates. Because most metals melt only at very high temperatures, a thermometer that uses a bimetallic strip can measure temperatures as high as 1000° C. The dial thermometer, used in most
  11. 11. homes as an oven thermometer, is an example of a bimetallic thermometer. A curved bimetallic strip, with the faster-expanding metal on the outside of the bend, is attached to a pointer. Upon heating, the bimetallic strip moves, causing the pointer to indicate the temperature on a circular scale. Other metallic thermometers, called resistance thermometers, use the principle that the resistance of a wire changes with temperature. Such thermometers also measure high temperatures. THE FAHRENHEIT AND CELSIUS TEMPERATURE SCALES Temperature markings on thermometers are indicated in Fahrenheit degrees or Celsius degrees. The Fahrenheit and Celsius scales are named after their originators, Gabriel Fahrenheit and Anders Celsius. (The Celsius scale is also called the centigrade scale.) Both Fahrenheit and Celsius scales are calibrated by using the boiling and freezing points of water. The Fahrenheit scale is used in the English system of measurement and the Celsius scale in the metric system. In the Fahrenheit scale, the freezing point of water is 32.° F, and the boiling point of water is 212° F. The remainder of the scale between these two points is marked off into 180 equal divisions (212 - 32 = 180) . In the Celsius scale, the freezing point of water is 0° C and the boiling point of water is 100° C. The remainder of the scale between these points is divided into 100 equal divisions (100 - 0 = 100). Note that there are 180 divisions between the freezing and boiling points of water in the Fahrenheit
  12. 12. scale and 100 divisions between these points in the Celsius scale. Thus, each Celsius division ( degree ) is 9/5 as large as Fahrenheit division. This relationship, together with the fact that there are 32 Fahrenheit divisions between 0 °F and 32° F, makes it possible to convert one scale into the other by using the following formulas: ° C = 5/9(° F – 32) ° F = 9/5 °C + 32 THE KELVIN SCALE Confined gases, like most solids and liquids, expand and contract uniformly. For this statement to be true, however, a gas must be heated or cooled in such a way that the pressure remains constant. ( Recall that the air thermometer is inaccurate because it is affected by surrounding air pressure.) If we start at 0° C, we find that, for every Celsius degree rise in temperature, the volume of a gas increases 273 of its original volume ( provided the pressure does not change). Similarly, if we again start from 0° C, we find that for every Celsius degree drop in temperature, the volume decreases 273 of its original volume. At -273° C, the volume of a gas would shrink to zero and all molecular motion would cease. This, in turn, means that the gas would contain no heat. (Actually, gases generally liquefy before this temperature is reached.) Scientists refer to -273°C as absolute zero, a temperature that has never been attained, although some scientists have come very close to this point.
  13. 13. Absolute zero, -273°C, is also called 0 Kelvin ( 0 K ). The Kelvin scale, named after its originator, Lord Kelvin, is based on absolute temperatures. Since the Kelvin scale begins with absolute zero (-273° C), we use the following formula to convert the Celsius scale to the Kelvin scale: degrees Kelvin = degrees Celsius + 273 degrees This formula may be written as K = ° C + 273 Let us find the freezing point of water (0° C) in the Kelvin scale: K=0 + 273=273 K Thus, 0° C is equivalent to 273 K. Now, let us find the boiling point of water: K =100 + 273 = 373 K Thus, 100° C is equivalent to 373 K. TRANSFER of HEAT When a metal spoon is placed in a bowl of hot soup, the entire spoon soon becomes hot because the heat travels from the soup to the bowl-shaped part of the spoon, and then to the handle. When ice is placed in warm water, the ice soon melts. Both of these examples show that heat travels from one body to another. Generally, when objects are at different temperatures, heat is transferred from the warmer object to the cooler object until both objects are at the same temperature. Heat transfer can occur through one of three methods: conduction, convection, or radiation. CONDUCTION
  14. 14. When one end of a metal rod is held in a flame, the entire rod will become hot enough to burn the hand. The heat from the flame reaches the hand by traveling through the rod. Substances that allow heat to travel through them are called conductors. In general, as we learned before, metals are good conductors. However, some metals conduct heat more readily than others. This can be demonstrated by inserting rods of aluminum, copper, iron, nickel, and brass into a brass sphere or disk and then attaching a small ball of wax to the end of each rod. When the center of the brass disk is heated, the wax at the tip of each metal melts in the order in which the different metals conduct heat. The wax at the tip of the copper melts first and the wax at the tip of the iron melts last. Conduction in most materials can be explained by the kinetic-molecular theory. When one end of a rod is heated, the molecules in that end of the rod vibrate faster and strike other nearby molecules, causing them to vibrate faster. In this manner, the increased molecular motion is transferred from one end of the rod to the other, permitting the heat to travel through the rod. Substances that do not readily allow heat to pass through them are called insulators. Gases and liquids are generally poor conductors of heat because their molecules are farther apart than are the molecules in solids. Therefore, neighboring molecules in a gas or in a liquid are less affected by the increased motions of heated molecules, and consequently heat is not conducted rapidly.
  15. 15. Substances like wood or plastic are poor conductors of heat, so they are used to make handles for metallic objects that are to be heated. The clothing we wear is also a poor conductor of heat, enabling us to retain body warmth. Porous material is generally non-conducting because it contains layers of trapped air which do not permit heat transfer. CONVECTION Although gases and liquids are poor conductors of heat, heat is transferred through them by the process of convection. Convection is the transfer of heat due to the motion of the liquid or gas itself. For example, when a beaker of water is heated the water layer closest to the heat source is warmed slowly by conduction. As the water becomes warmer, it expands, becomes less dense, and rises. This brings heat to the upper layer. At the same time, cooler water from the upper portion of the beaker moves down, takes the place of the rising water, and becomes heated itself. When warm enough, this water rises and carries heat upward. As these processes continue, heat that enters the bottom of the beaker is distributed throughout the beaker until all the water is at the same temperature. The moving water in such a case is said to have set up a convection current. Heat is also transferred through gases by convection. It is by this means, in part, that a stove or a radiator heats a room. Heat from the radiator warms the air above it, causing the air to expand, become less dense, and rise. The cooler air that moves in to take the place of the warmed air
  16. 16. is also soon warmed. As this air rises, a convection current is established. The convection current continues to distribute heat throughout the room until the entire room is warmed. The formation of a convection current in air is demonstrated with a convection box apparatus. First the candle is lighted, then smoking touch paper is placed over the chimney, opposite the candle. The smoke, coloring the air, can be seen to move down this chimney, across the box, and out through the other chimney. This occurs because the air over the candle is heated, becomes less dense, and rises, leaving a partial vacuum. Cooler, more dense air from the first chimney moves in to fill the partial vacuum. This cycle continues as long as heat is given off by the burning candle. RADIATION We know that light energy and heat energy travel from the sun to the earth through space, which is an almost perfect vacuum. These forms of energy, traveling without the aid of molecular collisions, are transferred from the sun to the earth by radiation, that is, by means of rays, or waves. You can understand this method of heat transfer by standing a short distance from an open fire or by placing your hand a little to one side of, but not touching, a hot radiator. Since neither source of heat is being touched, you cannot receive heat by conduction. Since warm air rises vertically from the heat source, the heat cannot reach you by convection. The heat that is transferred to you from the fire or radiator reaches you by radiation.
  17. 17. The heat radiated by one body ( the sun, for example) is most rapidly absorbed by other bodies that are black in color and rough in texture. In warm climates, white clothing which reflects the radiant heat of the sun is cooler than dark clothing which quickly absorbs the radiant heat. Similarly, bodies that are rough and dark tend to radiate heat better than shiny smooth bodies. This is why steam radiators are often dark and have a roughened surface. It is for the same reason that coal burning stoves are black. Bodies that are shiny and smooth do not absorb heat readily. Instead, these bodies reflect heat. Thus, aluminum used for roofing keeps homes cool in the summer and warm in the winter. This principle is utilized in the thermos (vacuum) bottle, which is so constructed as to permit liquids to retain their temperatures for a long time. A thermos bottle is double walled, with a partial vacuum between the walls to prevent heat transfer by conduction or convection. A cork stopper also prevents heat transfer by conduction. The inner glass walls are silvered to reflect radiant heat back into the liquid, thereby minimizing heat loss by radiation. Thus, a hot liquid remains hot because heat is lost very slowly. A cold liquid remains cold in thermos bottles because outside heat enters very slowly by conduction, convection, or radiation. MEASURING HEAT We learned that temperature is a measure of the average kinetic energy of the molecules of a substance. This is the same as saying that temperature represents the average intensity of the motion of the molecules, or the
  18. 18. degree of hotness of a substance. Average kinetic energy means the total kinetic energy divided by the total number of particles. Recall that the ocean contains much more heat than does a small amount of molten iron. However, since the ocean contains many more particles than the molten iron, the temperature of the ocean (that is, the total kinetic energy divided by the total number of particles) is lower than that of the molten iron. Temperature, therefore, does not tell us the quantity of heat present. The quantity of heat represents the total kinetic energy contained by allof the particles of the substance. In the metric system, we measure the quantity of heat by a unit called the calorie. The calorie is the amount of heat needed to raise the temperature of 1 gram of water 1 Celsius degree. In the English system, heat is measured by a unit called the British Thermal Unit (BTU). A BTU is the amount of heat needed to raise the temperature of 1 pound of water 1 Fahrenheit degree. The amount of heat energy present in a substance cannot be measured directly with simple measuring devices. Instead, it is measured by observing its effect on a given quantity of water in a device called a calorimeter. One type of calorimeter, consists of two polished metal cups surrounded by air, a poor conductor of heat. An insulating cover, holding a thermometer, makes up the top of the calorimeter. The polished cups reflect heat, thus maintaining the temperature of the liquid in the container. To determine the amount of heat energy absorbed (or lost) by a given quantity of water, we multiply the weight of the
  19. 19. water in grams by the change in temperature of the water in Celsius degrees. Thus: amount of heat = weight of water X change in temperature In a calorimeter, when 20 grams of water at 20°C are heated to a temperature of 30°C, how much heat is absorbed? The temperature change = the final temperature - the initial temperature = 30°C - 20°C = 10 Celsius degrees. Substituting, amount of heat = 20 grams X 10 C° = 200 calories We conclude that 200 calories have been ab sorbed. (We assume that no heat has escaped from the calorimeter.) HEAT EXCHANGE IN WATER In a calorimeter, when a quantity of water at a given temperature is mixed with a quantity of water at a different temperature, the amount of heat lost by the "hot" water is equal to the amount of heat gained by the "cold" water. Suppose we mix 100 grams of water at 90° C with 100 grams of water at 40° C and find that the final temperature of the mixture is 65° C. Let us calculate the number of calories lost. by the hot water and gained by the cold water. The temperature of the hot water dropped from 90° C to 65° C, a decrease of 25 Celsius degrees. Since we began with 100 grams of hot water that underwent a temperature change of 25° C, we determine the amount of heat lost:
  20. 20. amount of heat = weight of water X change in temperature (calories) (grams) (Celsius degrees) amount of heat = 100 grams X 25° C amount of heat = 2500 calories The minus sign in the answer indicates that 2500 calories of heat have been lost. The temperature of the cold water increased from 40° C to 65° C, an increase of 25 Celsius degrees. Since we began with 100 grams ofcold water that underwent a temperature change of 25° C, we determine the amount of heat gained: amount of heat = weight of water X change in temperature (calories) (grams) (Celsius degrees) amount of heat = 100 grams X 25° C amount of heat = 2500 calories Note that the amount of heat lost by the hot water (2500 calories) is the same as the amount of heat gained by the cold water (2500 calories). We assume that the heat exchange was "perfect" and that no heat escaped from the calorimeter. The quantity of heat needed to raise the temperature of 1 gram of a substance 1 Celsius degree is called the specific heat of the substance. For water, the specific heat is 1. This
  21. 21. means that 1 calorie of heat will raise the tem perature of 1 gram of water 1° C. Water is the only substance for which this is true. Other substances vary in the quantity of heat needed to raise 1 gram of the substance 1 Celsius degree. Consequently the formula amount of heat =weight of water X change in temperature applies only to water. CALORIES AND FOOD Your body requires energy in order to per form its daily tasks. Most of this energy comes from energy-rich foods such as carbohydrates and fats. This energy is released when the body utilizes these foods. Using special calo rimeters, scientists have measured the energy content, or the number of calories present, in fixed quantities of certain foods. For example, a slice of white bread contains about 60 000 calories; a typical chocolate bar may contain about 300 000 calories. Nutritionists use a special kind of notation when discussing the calorie content of foods. They define a food Calorie ( written with a capital letter) as 1000 calories. On a calorie table, therefore, we would read that a slice of white bread contains about 60 Calories and that a chocolate bar contains about 300 Calories. This information comes from "Physical Science Workbook", 1961
  22. 22. The History of Heat Aristotle and the Greeks had their idea of fire as one of the 4 Primal Elements. Even the ancients realized that heat and light were not alike as aspects of fire, though. After the fire had gone out and the light gone, the heat of the kettle and its contents remained. First modern chemist to study heat was Joseph Black (1728 - 1799). Black tried to explain heat in terms of a fluid. He explained how a kettle of water placed over a fire increased in temperature but a kettle filled with water and ice placed over a fire did not change in temperature till all the ice was melted. He said that until the ice was saturated with the heat-fluid and thus became melted could its temperature rise. Lavoisier accepted this theory and gave the name for this heat-fluid “caloric” from the Latin word for heat. Another idea competed with the caloric theory. Scientist knew that kinetic energy of motion plus the stored energy called potential energy was given the name mechanical energy and that friction was a part of the conservation of these energies. They knew friction could warm up an object so maybe the invisible motion of invisible particles was what we call heat. Summed up; friction was converting mechanical energy into heat. The
  23. 23. problem was this idea of really small particles of matter (i.e., atoms and molecules). Count Rumford (really Benjamin Thompson - a spy for the British authorities during the Revolutionary War) was to supervise the boring of cannon for Bavarian army. Using a horse to work a treadmill he realized that the solid block of brass grew hot as the borer cut its way in. Rumford calculated that if the caloric theory were correct the heat released during the boring would have melted the entire block of metal first. He pointed out that heat was produced without fire, without light, without chemical combustion. It came just out of motion. John Dalton comes along with his ideas of atoms and the kinetic energy idea of heat began to gain favor. An English physicist, James Prescott Joule (1818-1889) was attempting to find the mechanical equivalent of heat. In the end he found that a given amount of energy of whatever form always yielded that same amount of heat (at 4.18 joules per calorie). The relationship of the motion of atoms to temperature and heat was placed on firm theoretical basis about 1860 by the Scottish physicist James Clerk Maxwell. Specific Heat
  24. 24. The ability of water to stabilize temperature depends on its relatively high specific heat. The specific heat of a substance is defined at the amount of heat that must be absorbed or lost for 1 g of that substance to change its temperature by 1º C. The specific heat of water is 1.00 cal/g ºC. Compared with most other substances, water has an unusually high specific heat. For example, ethyl alcohol, the type in alcoholic beverages, has a specific heat of 0.6 cal/g ºC. Because of the high specific heat of water relative to other materials, water will change its temperature less when it absorbs or loses a given amount of heat. The reason you can burn your finger by touching the metal handle of a pot on the stove when the water in the pot is still lukewarm is that the specific heat of water is ten times greater than that of iron. In other words, it will take only 0.1 cal to raise the temperature of 1 g of iron 1ºC. Specific heat can be thought of as a measure of how well a substance resists changing its temperature when it absorbs or releases heat. Water resists changing its temperature; when it does change its temperature, it absorbs or loses a relatively large quantity of heat for each degree of change. We can trace water’s high specific heat, like many of its other properties, to hydrogen bonding. Heat must be absorbed in order to break hydrogen bonds, and heat is released when hydrogen bonds form. A calorie of heat causes a relatively small change in the temperature because must of the heat energy is used to disrupt hydrogen bonds before the water molecules can begin moving faster. And
  25. 25. when the temperature of water drops slightly, many additional hydrogen bonds form, releasing a considerable amount of energy in the form of heat. What is the relevance of water’s high specific heat to life on Earth? By warming up only a few degrees, a large body of water can absorb and store a huge amount of heat from the sun in the daytime and during summer. At night and during winter, the gradual cooling water can warm the air. This is the reason coastal areas generally have milder climates than inland regions. The high specific heat of water also makes ocean temperatures quite stable, creating a favorable environment for marine life. Thus, because of its high specific heat, the water that covers most of planet Earth keeps temperature fluctuations within limits that permit life. Also, because organisms are made primarily of water, they are more able to resist changes in their own temperatures than if they were made of a liquid with a lower specific heat. Water is one of the few substances that are less dense as a solid than as a liquid. While other materials contract when they solidify, water expands. The cause of this exotic behavior is, once again, hydrogen bonding. At temperatures above 4º C, water behaves like other liquids, expanding as it warms and contracting as it cools. Water begins to freeze when its molecules are no longer moving vigorously enough to break their hydrogen bonds. As the temperature reaches 0º C, the water becomes locked into a crystalline lattice, each water molecule bonded to the maximum of four partners. The hydrogen bonds keep the
  26. 26. molecules far enough apart to make ice about 10% less dense than liquid water at 4º C. When ice absorbs enough heat for its temperature to increase to above 0º C, hydrogen bonds between molecules are disrupted. As the crystal collapses, the ice melts, and molecules are free to slip closer together. Water reaches it greatest density at 4º C and then begins to expand as the molecules move faster. The ability of ice to float because of the expansion of water as it solidifies is an important factor in the fitness of the environment. If ice sank, then eventually all ponds, lakes, and even the oceans would freeze solid, making life as we know it impossible on Earth. During summer, only the upper few inches of the ocean would thaw. Instead, when a deep body of water cools, the floating ice insulates the liquid water below, preventing it from freezing and allowing life to exist under the frozen surface. Specific Thermal Electrical Metal Density Heat Conductivity Conductivity k cp g/cm3 1E6/Ωm watt/cm K cal/g° C Brass 0.09 1.09 8.5 Iron 0.11 0.803 7.87 11.2 Nickel 0.106 0.905 8.9 14.6 Copper 0.093 3.98 8.95 60.7 Aluminum 0.217 2.37 2.7 37.7 Lead 0.0305 0.352 11.2
  27. 27. Heat and Temperature Teaching Notes 1) Heat flows from hot to cold areas due to a temperature difference only. example: A small hot block of a material is placed next to a larger, cooler block. Heat flows from the small hot block to the larger block till equilibrium is reached. 2) Note the difference between the heat content and temperature. A lake may be cooler in temperature than a liter flask of water but the lake has a much greater heat content due to the vast number of particles and their associated motion. 3) Human perception of heat and temperature is not adequate for scientific work. So we must investigate tools that provide the accuracy and repeatability we need. 4) For temperature we will be using thermometers. There are other devices that allow you to measure temperature. 5) For heat we will be working with simple calorimeters. We will look at these devices when we look at the transfer of heat. Thermometers and Temperature Scales see comparisons charts
  28. 28. thermometers = instruments to measure temperature see drawing: gas (air) thermometers see drawing: liquid (Hg and alcohol) see drawing: solid (bimetallic) Mercury thermometer -- fill thin glass tube with Hg at a temp. greater than the maximum to be measured; tube is cut and sealed; the Hg cools and contracts leaving a partial vacuum above the Hg (eliminates effects of air resistance on expansion of Hg); then calibrate thermometer Hg freezes at -39° C so it cannot be used for low temperatures (use alcohol which freezes at -114° C) can use Hg for high temperature boils at 357° C (alcohol cannot be used for high temperature work due to its low boiling point - 78° C) use alcohol thermometers in schools unless extreme accuracy is needed due to safety factor. The alcohol may be inaccurate by 1 - 2 degrees but this may not be a problem if you are looking at changes in temp.
  29. 29. Calibration Both Celsius (Anders Celsius) and Fahrenheit (Gabriel Fahrenheit) scales are established by using the boiling and freezing point of water at 1 atmosphere of pressure. step 1 -- establish a mixture of ice and water in equilibrium (0° C) mark point of liquid in thermometer at 0° C establish a mixture of steam and water at equilibrium (both at a pressure of 1 atm.) steam condensing and water vaporizing) label this point of liquid as 100 °C step 2 -- divide the interval between 0° C and 100° C into 100 equal parts, each representing a change in temperature of 1° C. using this scale you can extend your marks below 0 °C and above 100 °C as far as ,you wish. The Fahrenheit scale labels the freezing point at 32° (his label for the temperature he could achieve with an ice and water mixture and labeled the boiling point temperature of water at 212 ° which was a number chosen for convenience apparently creating 180 divisions.
  30. 30. You might ask your self about the amount of heat energy need to cause a 1 degree change in temperature on a Celsius scale compared to that needed on a Fahrenheit scale. (more heat needed to cause change of 1 degree on Celsius scale.) KELVIN scale We know that gases decrease in volume 1/273 of its original volume for each degree drop in temperature. Thus at -273° C the volume of gas would shrink to zero and the gas would have no molecular motion. We know this is impossible (particles have zero- point energy). To label these very low temperatures a scale called the absolute or Kelvin scale is often used. It designates -273° C at the zero point, and is called called absolute zero. Extrapolation to absolute zero: A good research project for students or a quick review of graphing can be done by using the method to calculate absolute zero. Use a capillary tube with a trapped bubble of air between light machine oil. Measure the length of the bubble after placing it in different temperature mixtures. Plot the length versus temperature. Since it is not easy to obtain very cold temperatures, the linear series of points that you did obtain should allow you to extrapolate, (extend a curve beyond the known data points following the apparent pattern of the curve) until it intersects the temperature axis. See actual plot!
  31. 31. Transfer of Heat by Conduction, Convection, Radiation Conduction is a consequence of the kinetic behavior of matter. Faster vibrating particles collide with less energetic neighbors and transfer some of their kinetic energy to the slower moving particle. Through successive molecular collisions energy travels through a material without the average position of the particles being changed. There must be a temperature differential (one end of some object at a higher temperature than the other) for heat to be conducted. Gases are poor conductors of heat (compared to liquids and solids) because the molecules are relatively far apart and collisions are infrequent. Metals have the greatest ability to conduct heat (for the same reason as their high electrical conductivity). This is due to a significant number of electrons being able to move about freely instead of being bound permanently to particular atoms. Thermal conductivity of a material is a measure of its ability to conduct heat. Example: wood and metal (see class discussion) Convection involves the actual motion of a hot fluid from one place to another, displacing a colder fluid in its path and setting up a convection current. Convection is the chief mechanism of heat transfer in fluids.
  32. 32. Natural convection occurs when the buoyancy of heated fluids leads to motion. Heated fluids (gas or liquid) expand and becomes less dense than surrounding cooler fluids. It then rises. Radiation is defined as the energy that is transmitted by electromagnetic waves and requires no material medium for passage. All objects radiate electromagnetic waves with the higher the temperature of an object the shorter the predominating wavelength of its radiation. Example: see glass lined thermos bottle in heat packet in class. Transfer of Heat Conduction -- place iron rod in fire -- the end you are holding becomes warm due to conduction Convection -- stove heats room by convection Radiation -- heat the earth receives from the sun is radiation Natural direction of heat flow is from hot bodies to cold ones. Conduction -- conduction is a consequence of kinetic behavior of matter faster vibrating particle collide with less energetic neighbor and transfer some of their kinetic energy the slower moving particle
  33. 33. example: place hand on wood and metal sample at same temperature. The metal will seem cooler because it conducts the heat away much faster than the wood Convection -- actual motion of hot fluid from one place to another, displacing cold fluid in its path setting up a convection current = chief mechanism of heat transfer in fluids in most instances natural convection - the buoyancy of heated fluids leads to motion - heated fluid (gas or liquid) expands and becomes less dense than surrounding cooler fluids and rises Radiation -- energy that is transmitted by electromagnetic waves and requires no material medium for passage all objects radiate electromagnetic waves but the higher the temperature of an object the shorter the predominate wavelength of its radiation There are many examples you can use to demonstrate these three ideas but discussing a glass lined thermos bottle will allow you discuss them as well as High specific heat capacity material demonstrates relatively small change in temperature for a given change in internal energy content add 1 calorie of heat to 1 gram of water, helium, ice, gold temperature rises: water 1 °C
  34. 34. helium 1.3 °C ice 2 °C gold 33 °C Calorimetry: 1) Use 2 polystyrene cups (one within the other) -- the polystyrene will not absorb very much heat. 2) You may find if difficult for the students to be patient when working with the calorimeters. Your step-by-step procedure must be very simple and clear. This type lab is also a very dangerous time for your thermometers. 3) Several labs have been included in the packet involving the use of the calorimeters and thermometers: a) the heat of reactions and heat of solutions labs deal with endothermic and exothermic reactions (and you can incorporate the use of the mole as review) b) the specific heat capacity lab is an excellent demonstration of the principle of heat exchange as well as specific heat capacity of different metals and liquids. 1) for better results with this lab try to use as much metal as possible and as little fluid as possible and still cover the metal in the cups. Drain the metal samples quickly after removing them from the boiling water. This is a good time to have the students watch the boiling process which will be one of the final topics in this unit. 2) the math may be a bit difficult for you at first.
  35. 35. Thermal Expansion of water 1) From 0° C to 4° C the volume of water in a sample decreases (the greatest density is at 4° C) 2) We know that ice floats (less dense) so that a body of water in winter freezes from the top down. The ice is a poor conductor of heat so that the initial layer of ice that freezes impedes further freezing allowing fish and plant life to live through the winter. 3) The spaces between molecules in ice are greater than the same spaces in liquids. 4) Ice has what is called an Open Structure --> each water molecule can participate in 4 bonds with other water molecules, while other solid molecules can have as many as a dozen bonds with surrounding molecules resulting in a more compact substance. 5) As stated the density of the water increases from 0 °C to 4 °C. Large clusters of water molecules break into smaller clusters that occupy less space in the aggregate as the temperature rises to 4 °C. Only above 4 °C does the normal thermal expansion show a decreasing density with increasing temperature. Above 4° C , the normal thermal expansion of materials is seen. Here as the temperature rises the density decreases. Heat and Temperature
  36. 36. HEAT Heat is a form of internal energy which is transferred from one object to another due to a difference in temperature between the objects. Heat is the total energy of motion of all particles (the total kinetic energies of all the particles.) TEMPERATURE The temperature of a body of matter is a measure of the average kinetic energy of the random motion of its particles. Temperature is the kinetic energy divided by the number of particles. Temperature is that property of a substance which determines whether it is in thermal equilibrium with another object. Thermal Conductivity - a measure of the ability of a substance to conduct heat THERMAL EQUILIBRIUM This is the situation in which no heat moves from one object to another. CALORIE A 15° Calorie is the amount of heat energy needed to change the temperature of of 1 gram of water by 1° C (from 14.5° C to 15.5° C at 1 atmosphere of pressure). 1 calorie = 4.185 Joules and 1 kilocalorie = 1000 calories. SPECIFIC HEAT CAPACITY This is the amount of heat (in calories or Joules) that must be added or removed from a unit mass of that substance to change its temperature by one degree. Different substances have different capacities because they absorb and release heat at different rates.
  37. 37. WATER Water has a specific heat capacity of 1.00 cal/ g °C or 4.185 Joules/g °C. The SI unit would be 4185 J/kg °C. PRINCIPLE OF HEAT EXCHANGE The heat lost by an object must equal the heat gained by the object to which the heat is transferred. There must be a temperature difference for heat to be transferred. Q (heat energy) = m (mass) x At (temp.) x cp (specific heat capacity) (cal/Joules) (g) (°C) (cal/g °C) or (Joule/g °C) Problems: 1) How much heat energy is needed to raise the temperature of 100 grams of water from 0 degrees to 30° C? 2) A calorimeter contains 300 grams of water at 10° C. After a food sample is burned in the calorimeter the water temperature changes to 15° C. How much heat was given off by the food sample? LATENT HEAT Latent heat is the heat required to bring about a change in state. HEAT OF FUSION The heat of fusion is the amount of heat that must be supplied to change a unit mass of the substance at its melting point from solid to liquid. The heat of fusion of water is 80 calories per gram (80 kcal/kg).
  38. 38. HEAT OF VAPORIZATION The heat of vaporization is the amount of heat that must be supplied to change a unit mass of the substance at its boiling point from liquid to gas or vapor state. For water this is 540 cal/g or 540 kcal/kg. HEAT OF SUBLIMATION The heat of sublimation is the heat needed to change a solid to a gas. HEAT OF CONDENSATION The heat of condensation is the reverse of the heat of vaporization, it is the heat given off when a gas condenses to a liquid. Four States of Matter Matter is defined as any material that has mass, occupies volume, and exhibits inertia (resistance to movement). Solids definite shape and volume, resist deformation very close spacing of particles that make up the solid these particles appear to vibrate about fixed points particles vibrate faster at higher temperatures, slower at lower temperatures Crystalline solids particles are arranged in regular, repeated patterns - said to have "long-range order" to their structure -example would be NaCl (table salt) Amorphous solids solids that lack the definite arrangement in crystals are `amorphous' which means `without form'. Can be thought of as liquids whose stiffness is due to exaggerated viscosity. These solids are
  39. 39. said to have "short range order". Examples are pitch, glass, plastics. Polymers are flexible and some will change their structure when undergoing a physical change (examples are rubber bands, Saran Warp, Lucite, DNA, fats, cellulose, glycogen. Jello is a natural glucose polymer. Liquids definite volume, resist compression, will flow, takes the shape of its container greater spacing between molecules, liquid particles appear to travel in straight line paths between collisions but appear to rotate or vibrate about moving points Gases Have no definite shape or volume, takes the shape and volume of its container can be compressed or dispersed, the particles vibrate very rapidly, are relatively far far apart, and there are no forces holding them together Plasma very high temperature ionized gas (as high as 100 million degrees in some fusion reactors). These plasmas have no fixed volume or shape, most are mixtures that are not easily containable. They all have particles that are electrically charged and of low density. The Milky Way is a huge plasma. Energy Definitions Energy: having the ability to do work (move matter) Work: a push or pull over some distance (force x distance)
  40. 40. Force: a push or a pull Potential Energy: the energy a body possesses by virtue of its position, composition, and or condition stored energy or energy of position P.E. = mass x gravity x height examples: water behind a dam, stretched/compressed spring, explosives Kinetic Energy: the energy of motion (conserved in every elastic collision) K.E. = 1/2 mass x velocity2 heat energy flows from hot objects to cooler ones through transfer of K.E. when particles collide Momentum: mass time velocity (momentum is conserved in every collision where there is no friction) Linear momentum of a moving body is a measure of its tendency to continue in motion at a constant velocity. The conservation of linear momentum states that in the absence of forces from outside the system the total momentum of colliding particles cannot change but the distribution of the total momentum may change. Momentum is redistributed in a collision. Intermolecular Forces: potential energy forces that hold molecules together and in correct position in solids potential energy forces that hold molecules together in liquids the kinetic energy of the molecules in solids and liquids
  41. 41. cannot overcome intermolecular forces holding the molecules together (so they do not fly apart) gas molecules have enough kinetic energy to break free from intermolecular forces or to keep such forces from forming Kinetic – Molecular Theory of Gases gases are made up of molecules that are in continuous motion an increase in the temperature increases the speed of the molecules, thus increasing the kinetic energy of the substance All gases are compressible Gases display diffusion (random movement of molecules from one area to another with a net change in concentration – rate varies with temperature and molecular mass) Gases can be liquefied (called liquefaction) Closed System Criteria In using the above information we look at pressure, temperature, and volume in a closed system. 1. In a closed system nothing escapes or is allowed in (unless we choose to allow it) 2. all molecules are in motion (have K.E.) 3. molecules exert a uniform pressure on all surface areas of the walls of the container 4. Pressure = force/area (see examples given in class)
  42. 42. 5. Atmospheric pressure is the cumulative effect of the force generated by the weight of the atmosphere. Given values that must be used in problems include: 14.7 lb/in2, 101.3 kPa, 1 atmosphere, 760 mm of Hg, 1 033.6 g/cm2 6. Molecules exert pressure on other molecules inside container as they collide, push, and bounce off other molecules 7. The pressure a gas exerts on the walls of its container is the sum of the forces acting on the walls (equals the frequency of collisions with the walls of the container plus the force of each molecule as it pushes against the wall) due to the random collision of limitless numbers of these moving molecules. Collisions that occur between molecules are perfectly elastic, the particles bounce off each other and exchange energy, but there is no loss of energy * elastic atomic collisions: atoms (molecules) bounce back as far/fast as it would have had it not collided (no change in the total kinetic energy of the two particles before and after the collision) * inelastic collisions: the normal order in which the objects lose energy and slow down Momentum is conserved in every collision where there is no friction, energy is conserved only in elastic collisions. Gas Laws
  43. 43. 1. J.L. Gay-Lussac’s Law If the volume remains constant, the pressure is directly proportional to the absolute temperature: P ~ T P1 / T1 = P2 / T2 2. Boyle’s Law If the temperature remains constant, the volume of a gas varies inversely with the pressure: V ~ 1/P P1 V1 = P2 V2 3. Charles’ Law If the pressure is kept constant, the volume of a gas is directly proportional to its absolute temperature: V ~ T V1 / T 1 = V2 / T 2 For each degree increase in temperature, the volume increases 1/273 of its original volume 4. Combined gas law: P1 V1 / T1 = P2 V2 / T2 5. Ideal Gas Law: PV = nRT Overall conclusions: The temperature of a gas increases when it is compressed because the average energy of its molecules increases. The molecules rebound from the inward moving piston, traveling faster than before hitting the piston.
  44. 44. Molecules rebounding from fixed walls have unchanged speeds. The temperature of a gas decreases when its volume is expanded because the average energy of its molecules decreases. The molecules rebounding from outward moving piston move slower than before. Gas Law Problems: 1) An insulated system is known to have a temperature of 100.0° C at a pressure of 4.00 atm. If the absolute temperature is cut in half, what will be the new: ___________ atm, __________kPa, ______________° C, _______________ K 2) The volume is given as 27.0 L. If the pressure goes from 3.00 atm. to 9.00 atm., what is the new: ___________L, _____________ kPa 3) The volume is given as 5.00 L. If the absolute temperature goes from 273 to 819 K, what is the the new __________L (if new temperature was 800. K, what is the new volume in liters?) 4) The temperature is given as 25.0° C. If the volume is decreased from 100. mL to 10.0 mL, what is the new: ____________K, ____________° C
  45. 45. Thought Work -- Heat & Gas Laws 1. Devise a way to remove carbon dioxide (carbonation) from soda pop. This must be done quantitatively. How does temperature affect the solubility of gases being dissolved in liquids under pressure? 2. Explain what happens to a marshmallow when it is toasted. Why does this happen? What would happen to a frozen marshmallow? Why does a marshmallow float? 3. Find the pressure you exert when standing on both feet, on one foot, and lying flat on your back. 4. Explain why a toy balloon filled with hydrogen partially deflates overnight. 5. Using a steel ball and pieces of old pottery or modeling clay, devise an experiment that would demonstrate potential energy, kinetic energy, and momentum (all of which are involved with mass and velocity). 6. Suppose you had two identical sections of glass plate before you, one heated above body temperature and the other cooled by ice. What happens when you breathe on the two of them and why? Maybe try this at home first. 7. Devise an experiment to show the concept of diffusion, another to show cohesion, adhesion, and surface tension, and one to show buoyancy and Pascal's Law. 8. Changing ice to water requires 80 calories per gram of ice, but changing water to steam requires about 540 calories
  46. 46. per gram of water. What does this tell you about the intermolecular forces in ice and water, both qualitatively and quantitatively? Also explain why the chemical change of splitting or forming water requires about 5 times as many calories as the physical change of state. Chemical Properties of Matter Chemical properties are those properties of a substance that can be determined by a chemical test. Chemical properties are seen by the material's tendency to change, either alone or by interaction with other substances, and in doing so form different materials. 1.does the substance support combustion: examples are O2 and Cl2 2. does the substance burn (combustibility) 3. how does the substance react with acids (does it dissolve, evolve gases, explode, do nothing) 4. how does the substance react with oxygen (burn, form new compounds) 5. what is its reaction with electricity (usually it will be separated into simpler components) examples: alcohol burns, iron rust, wood decays, sodium explodes in water Physical Properties of matter
  47. 47. Physical properties are those properties used in identifying substances when we use our senses. These do not require chemical analysis. 1. color - reaction of eye and brain in recognizing combinations of certain wavelengths of visible light. 2. hardness - a measure of the ability of a substance to resist abrasion (see Mows Scale of Hardness) 3. density - the mass divided by its volume (often reported as specific gravity which is a unitless relationship between the density of the substance and the density of water) 4. texture - how object feels to touch; usually rough or smooth 5. magnetic attraction - is the material attracted to a magnet or can it be magnetized (must contain Fe, Co. Ni, or steel) 6. solubility - the amount of a substance which will dissolve in a known amount of solvent at a given temperature 7. taste - reaction of taste buds to stimuli along with the brain's recognition of the pattern 8. light transmission - is the substance transparent, translucent, or opaque
  48. 48. 9. viscosity - a measure of the internal resistance (friction) to flow in a liquid (molasses and tar would have high viscosities) 10. refractive index amount a ray of light is bent as it passes through a substance (technically the ratio of the speed of light in that substance to the speed of light in a vacuum) 11. specific heat capacity - the amount of heat energy (calories or Joules) required to change the temperature of 1 gram of a substance by 1 °C 12. atomic radius - the distance from the center of an atom's nucleus to the outermost orbital electron 13. boiling point - the temperature at which the liquid's vapor pressure equals atmospheric pressure during the boiling of a pure substance the temperature remains constant as long as both liquid and vapor are present 14. melting-freezing point - temperature at which solid- liquid phase is in equilibrium - during melting of pure solid the temperature remains constant; when all solid is melted and only liquid is present, further heating results in a steady increase in temperature to the boiling point 15. odor - olfactory nerves are stimulated by certain molecule and send messages to the brain which remembers the pattern 16. expansion - contraction coefficients - materials expand or contract a known amount when heated or cooled
  49. 49. Collapsing Can Demo area of surface of Coke can = 0.031 m2 pressure on this area = 3.1 E 3 N (about 680 lb) 1 atm = 1.0 E 5 N/m2 if we can reduce pressure by as little as 75% there would be a 500 lb difference between pressure inside and outside the can 1) when can is inverted in water bath - the water seals the opening and cools the can 2) as can cools vapor condenses - reduction of pressure inside can 3) can is sufficiently weak and water sufficiently viscous that can collapses before it fills with water 4) must use all aluminum can Physical Changes in State a change in the physical properties of a substance without a change in the chemical composition the arrangement of molecules may be changed but the molecular make-up remains the same
  50. 50. these changes deal with intermolecular forces which increase or decrease during the change. Problems: 1. 72 grams of ice + 51 840 calories yield 72 grams of water vapor. How many calories must be removed from water vapor to condense it back to ice? 2. If 1 gram of water at room temperature evaporates, about 600 calories are taken from the surroundings to convert the liquid to a gas. How many calories are `needed' to change 1001 grams of gas to a liquid? 3. If 50 grams of water vapor loses 36 000 calories in turning to ice, how many calories would 1 gram of water vapor have to lose to be turned back to ice? ice (0° C) + heat à water vapor (100° C) 36 g 25 920 cal 36 g water vapor (100° C) à ice (0° C) + heat 36 g 36 g 25 920 cal 2 H2 + O2 à 2H2O + heat energy released 4g 32 g 36 g 136 600 cal 2H2O + energy à 2H2 + O2
  51. 51. 36 g 136 600 cal 4g 32 g Chemical Changes in State (Phase) The molecular make-up (the specific arrangement of atoms) is changed, resulting in new substances being formed and energy changes occurring. EXOTHERMIC - any chemical change that releases energy is exothermic the amount of heat released is greater than the amount of heat used to start the reaction bond making is exothermic (energy is released into surroundings) example: oxidation à wooden splint burning ( heat, light, gases like CO2 and H2O being given off with carbon and ashes left over) other examples: burning H2 in O2, body reactions, dissolving metals in strong acids, mixing acid and water, homogenization, plaster of Paris in water, sugar dehydration ENDOTHERMIC - any chemical change that absorbs energy is endothermic energy continues to be absorbed as long as the reaction continues bond breaking is endothermic (energy is absorbed from surroundings) example: electrolysis à splitting some compound
  52. 52. (usually water) by running an electric current through it other examples: photosynthesis, pasteurization, canning vegetables Sugar dehydration demo here The chemical change involving splitting or forming water takes about 5 times as many calories as the physical change of state. The reason is that atoms (or molecules) are bonded together in a compound; the stronger the bond the more energy holding the parts together, thus more energy required to break these bonds. A physical change needs far less energy to overcome intermolecular forces holding groups of molecules together. Much more energy is needed to break bonds within molecules than to overcome the forces between molecules. physical change -- strength of intermolecular forces increased or decreased chemical change -- bonds formed or broken energy absorbed -- bonds broken or intermolecular forces overcome energy released -- bonds formed or intermolecular forces strengthened Problems: Tell whether each of the following is a chemical or physical change and further describe each chemical change as
  53. 53. endothermic or exothermic and the physical changes as absorbing or releasing energy. dry ice sublimates CO2 + H2O + sunlight à glucose air in heated tire expands burning coal water frozen into ice acid dissolves metal Endothermic vs Exothermic Reactions All chemical reactions involve bond breaking and bond making. Bond breaking is endothermic (energy is absorbed from surroundings) Bond making is exothermic (energy is released into surroundings) Imagine stretching a rubber band until it breaks. You must do work to stretch the band because the tension in the band opposes your efforts. You lose energy; the band gains it. Something similar happens when bonds break in a chemical reaction. The energy required to break the bonds is absorbed from the surroundings. Energy is absorbed or released when the heat capacities of the products and reactants differ. Usually this is
  54. 54. small. Remember that heat capacity is best thought of with a penny and specific heat best thought of as copper metal. Neutralization reactions are usually exothermic but when you add baking soda to vinegar it is slightly endothermic. The neutralization reaction actually does release heat: HC2H3O2 + NaHCO3 à CO2 + NaC2H3O2 (aq) + H2O This is because there is net bond formation. The products collectively have lower energy than the reactants. But evaporation of the liquid occurs as the carbon dioxide escapes from solution. Evaporation absorbs heat, cooling the liquid. (The expansion of the carbon dioxide gas bubbles as they are released also helps to cool the surroundings by Joule-Thomson cooling). The net result is an endothermic reaction. Mixing a strong acid with water is exothermic. Breaking a chemical bond requires energy (remember that stretching a spring until it breaks requires energy). Forming a chemical bond will release energy. So in a reaction that releases heat (exothermic) there must be net bond formation. Lets looks at HCl dissolved in water: HCl à H+ (aq) + Cl1- (aq) You would think at first this would be a heat absorbing (endothermic) process, because it looks like the bond between H and Cl is broken. But there is another reaction hiding here. The hydrogen ion reacts with water to form a complex of the form: H3O·(H2O)+n where n is a number
  55. 55. between 1 and 9. It is much easier just to write H+ (aq). Because the hydrogen ion is so tiny, a large amount of charge is concentrated in a very small area, and the polar water molecules are strongly attracted to it. This "hydration" of the hydrogen ion involves the formation of a covalent bond to one of the waters and a large number of strong hydrogen bonds, so it’s a strongly exothermic process. This causes the mixing of a strong acid with water to be strongly exothermic overall. Exothermic processes Endothermic processes making ice cubes melting ice cubes formation of snow in clouds conversion of frost to water vapor condensation of rain from evaporation of water water vapor a candle flame forming a cation from an atom in the gas phase mixing sodium sulfite and baking bread bleach rusting iron cooking an egg burning sugar producing sugar by photosynthesis forming ion pairs separating ion pairs combining atoms to make a splitting a gas molecule apart gas molecule mixing strong acids and mixing water and ammonium water nitrate nuclear fission melting solid salts
  56. 56. Heating Curve Information: This graph will aid in understanding the following information. Phase Change Diagram Solid - Gas Phase Change This change involves sublimation which is the direct change of a solid to a gas (deposition is the opposite). Examples include: moth balls (naphthalene), paradichlorobenzene, camphor, iodine crystals, and CO2 fire extinguishers (advantages: does not conduct electricity, colder than water, replaces O2 since CO2 is heavier and settles on ground area and the CO2 does not combust), will sublime away reducing cleanup - disadvantages include difficulty in keeping container pressurized over time, fact that you cannot use on living things due to extreme cold, and cost). Liquid - Solid Phase Change This discussion deals with melting-freezing point. A complete discussion of this concept using ice and heat units will be completed in class. See class discussion of ice cube. The addition of 1 calorie of heat to the ice cube at 0° C does not cause a change in the temperature of the ice cube though 1 calorie would change the temperature of 1 gram of water at 0° C.
  57. 57. It will take 80 calories just to melt the ice cube. That heat that is consumed in melting the solid is converted into potential energy. Freely moving molecules in liquids, with respect to intermolecular attraction, possess more energy than similar molecules bound rigidly in solids at the same temperature. Remember that temperature is a measure of the average kinetic energy only while heat content is a measure of the total kinetic energy plus potential energy possessed by that body. See class examples of the heat energy needed to change ice at any temperature to steam at any temperature. The melting-freezing point is defined as the temperature at which the solid and liquid phases are in equilibrium. This is the temperature at which a change of state between the solid and liquid phase can occur. Some of the solid will be melting and some of the liquid will be freezing. When a solid is heated to its melting point, its atoms or molecules acquire enough energy to shift the bonds holding them together so they form separate clusters. This clustering in liquids is confirmed with X-ray studies but the clusters are constantly shifting their arrangement unlike the permanent arrangement in solids. When heat is added to a solid the temperature of the solid will increase till it reaches the melting-freezing point. It will remain there until all the solid has melted and only
  58. 58. then can the temperature of the liquid rise according to its specific heat. Water molecules at 0° C contain more energy than the ice molecules at 0° C , not in the form of a faster more rapid motion but in the form of an ability to resist the attractive forces tending to pull them together. Melting points also depend on pressure (though not as much as boiling points.) Ice is strange in that its melting point decreases with increasing pressure. Almost all other materials show increasing melting points with greater pressures. The pressure an ice skater exerts on the ice due to the small area of the skate blade is usually enough to melt the ice creating a thin film of water that acts as a lubricant. On unusually cold days the pressure may not be enough to melt the ice and thus skating would be impossible. Liquid - Gas Phase Change The change from a gas to la liquid is condensation. This is due to cooling and/or a pressure change. In liquids, the energy of the particles is raised by adding heat. When some molecules have enough K.E. they break away from the liquid surface and become vapor. If the temperature falls, there is a decrease in the energy of the moving molecules and the liquid may eventually freeze to the solid phase.
  59. 59. Process of EVAPORATION: Molecules that have enough energy of motion (K.E.) break free from intermolecular forces and escape into the air as vapor. Some may return to the liquid is their energy is lost to other atoms. The liquid surface left behind is cooled. In evaporation the molecules that escape are the ones with the greatest velocity (heat) thus the average velocity and K.E. of the remaining particles is reduced. This results in cooling effects. Heat must be absorbed from the surroundings to continue the evaporation process. Adding heat increases evaporation because the VAPOR PRESSURE is increased. This is the pressure exerted by the vapor (gas) of a substance when it is in equilibrium with liquid or solid phase. The system is in equilibrium when the rate of evaporation equals the rate of condensation. The temperature at which the liquid's vapor pressure is equal to outside (atmospheric) pressure is that liquid's BOILING POINT. At this temperature the pressure of the vapor escaping from liquid equals the outside pressure. When the vapor pressure equals outside pressure bubbles of vapor form and push through to the surface. As they move into the gas phase we say this is boiling. Conduction of heat creates the gas, which rises because it is less dense than the liquid, as it strives for equilibrium.
  60. 60. The boiling point varies with atmospheric pressure. In mountains, the boiling point is below 100°C because the pressure of the atmosphere is less. Cooking requires longer times at high altitudes because of low boiling point Pressure cookers make food cook more rapidly because the foods can be heated above the normal boiling point without actually boiling. Intermolecular Forces and Latent Heat if we heat a mixture of ice and water, we find that no matter how much heat is transferred to the mixture, the temperature remains at 0° C until the last of the ice is melted. Only after all the ice is melted is heat converted into kinetic energy, and only then can the temperature of the water begin to rise. Experiment shows that 80 calories of heat must be absorbed from the outside world in order that 1 gram of ice might be melted, and that no temperature rise takes place in the process. The ice at 0° C is changed to water at 0° C. But if the heat gained by the ice is not converted into molecular kinetic energy, what does happen to it? If the Law of Conservation of Energy is valid, we know it cannot simply disappear. The water molecules in ice are bound together by strong attractive forces that keep the substance a rigid solid. In order to convert the ice to liquid water (in which the molecules, as in all liquids, are free of mutual bonds to the
  61. 61. extent of being able to slip and slide over, under, and beside each other) those forces must be countered. As the ice melts, the energy of heat is consumed in countering those intermolecular forces. The water molecules contain more energy than the ice molecules at the same temperature, not in the form of a more rapid motion or vibration but in the form of an ability to resist the attractive forces tending to pull them rigidly together. The Law of Conservation of Energy requires that the energy change in freezing be the reverse of the energy change in melting. If liquid water at 0° C is allowed to lose heat to the outside world, the capacity to resist the attractive forces is lost, little by little. More and more of the molecules lock rigidly into place, and the water freezes. The amount of heat lost to the outside world in this process of freezing is 80 calories for each gram of ice formed. In short, 1 gram of ice at 0° C, absorbing 80 calories, melts to 1 gram of water at 0° C; and 1 gram of water at 0° C giving off 80 calories, freezes to 1 gram of ice at 0° C. The heat consumed in melting ice or any solid, is converted into a sort of potential energy of molecules. Just as a rock at the top of a cliff has, by virtue of its position with respect to gravitational attraction more energy than a similar rock at the bottom of the cliff, so do freely moving molecules in liquids, by virtue of their position with respect to intermolecular attraction, possess more energy than similar molecules bound rigidly in solids.
  62. 62. It is the kinetic and potential energies of the molecules that together make up the internal energy that represents the heat content. It is kinetic energy only that is measured by the temperature. By changing the potential energy only, as in melting or freezing, the total heat content is changed without changing the temperature. In converting a gram of liquid water at 100° C to a gram of steam at 100° C what remains of the intermolecular attractions must be completely neutralized. Only then are molecules capable of displaying the typical properties of gases -- that is, virtually independent motion. In the earlier process of melting, only a minor portion of the intermolecular attractive force was countered, and the major portion remains to be dealt with. The latent heat of vaporization of water (the amount of heat required to convert 1 gram of water at 100° C to 1 gram of steam at 100° C) is 540 calories, almost seven times the earlier 80 calories needed in changing ice to water. The energy content of steam is thus surprisingly high. 100 grams of water at 100° C can be made to yield 10 000 calories as it cools to the freezing point. 100 grams of steam at 100° C can be made to give up 54 000 calories merely by condensing it to water. The water produced can then give up another 10 000 calories if it is cooled to the freezing point. It is for this reason that steam engines are so useful and hot water engines would never do as a substitute. If we boil water in a kettle its temperature remains at 100° C, no matter how fast we boil it, but we have to keep
  63. 63. adding heat to keep it boiling. Heat is absorbed by the molecules as they escape their liquid state and become a gas. The amount of heat needed to pull apart liquid molecules is called heat of vaporization (calories/gram). The heat of vaporization which a molecule must absorb before it can become a gas molecule is released by it when it cools again to liquid, as heat of condensation. Liquids with low boiling points, such as alcohol or ether, chill the hand as the molecules pick up their heat of vaporization and become a gas. The same is true for a glass of water, it will be cooler than room temperature. The kinetic molecular theory states that the kinetic energy depends on heat energy, which can be measured as temperature. A thermometer in boiling water and a thermometer in the vapor just above the boiling surface will read the same; 100° C at sea level. Therefore the average kinetic energy of the liquid molecules must be the same as the average kinetic energy of the gas molecules above it. An average molecule in the liquid state will be moving as fast as an average molecule in the gaseous state. Gas particles move in a straight line until they collide with another bit of matter, then they bound away in some other direction but always in a straight line, and without losing any of their energy to friction in the collision. The particles have perfect resilience. However, they will change their kinetic energies in the familiar way of all normal matter, as, for instance, do billiard balls. A slow-moving particle hit from behind by a fast one is speeded up, while the fast one is slowed down, but the sum total of their kinetic energies
  64. 64. remains the same. In the world of normal matter, perfect elasticity is unknown, as there is friction between surfaces. Two billiard balls when they collide will change each other's speed and direction of motion, and they will also roll to a stop in a short time. The ultimate particles of matter lose not a bit of their energies in collisions. They simply exchange speeds. If two particles collide, their total heat before and after is the same, but the originally slower particle after the collision is traveling faster and is therefore hotter, while the formerly speedier particle is now cooler and moving more slowly that it was. Heat and molecular motion, according to the theory, are two ways of speaking about the same thing. Heat/Temperature Problems: 1. Suppose a piece of iron (mass = 21.50 g at a temperature of 100.° C) is dropped into an insulated container of water (mass of water = 132 g and the temperature before adding iron was 20.0° C). What will be the final temperature of the system (at thermal equilibrium)? The specific heat of iron is 0.113 cal/g° C. 2. If 200. grams of water is to be heated from 24.0° C to 100.0° C to make a cup of tea, how much heat must be added? 3. Which is more effective in cooling a drink, 10 grams of water at 0° C or 10 grams of ice at 0° C? Explain your answer quantitatively.
  65. 65. 4. A 3.00 kg lead bar at 100.0° C is placed in 4.00 kg of water at 20.0° C. The final temperature of the lead bar would be ___________. (cp of lead is 0.0305 cal/g° C) 5. A 0.60 kg copper kettle holds 1.70 kg of water at 30.0° C. A 0.10 kg iron ball at 120.0° C is dropped into the water. What is the final temperature of the water? (cp of copper = 0.377 J/g° C and iron is 0.448 J/g° C) 6. A piece of iron with a mass of 20.50 grams at a temperature of 100.0° C is dropped into 140.00 grams of water at 40.0° C. What will be the final temperature of the system. The cp of iron is 0.45 J/g° C) 7. A sample of mercury metal is heated from 25.5° C to 52.5° C. In the process, 187 cal of heat are absorbed. What mass of mercury was in the sample? The specific heat of mercury is 0.033 cal/ g° C 8. A block of aluminum weighing 140. g is cooled from 98.4° C to 62.2° C with the release of 1080 cal of heat. From these data, calculate the specific heat of aluminum. 9. A cube of gold weighing 192.4 g is heated from 30.0° C to some higher temperature, with the absorption of 226 cal of heat. The specific heat of gold is 0.030 cal/ g ° C. What was the final temperature of the gold? 10. A total of 54.0 cal of heat are absorbed as 58.3 g of lead is heated from 12.0° C to 42.0° C. From these data, what is the specific heat of lead?
  66. 66. 11. A piece of erbium metal weighing 100.0 g and heated to 95.0° C is dropped into 200.0 g of water initially at 20.0° C. The final temperature of the mixture is 21.5° C. What is the specific heat of erbium metal? 12. A block of rhenium metal (specific heat = 0.0329 cal/ g° C) is heated to 88.2° C and then dropped into 100.0 g of water initially at 26.4° C. The final temperature of the mixture is 32.4° C. What was the mass of the block of rhenium? 13. When 258.6 g of benzene vapor is condensed to a liquid at its boiling point, 33 875 cal of heat are released. What is the heat of vaporization for benzene? 14. A sample of ethyl alcohol is converted from a liquid to a vapor with no temperature change. In the process 30 640 cal of heat are absorbed. What mass of ethyl alcohol was in the sample? The heat of vaporization of ethyl alcohol is 210. cal/g. 15. The heat of combustion of methane is 212.8 kcal per mole. How much heat will be produced in the combustion of 100.0 g of methane? 16. The heat of combustion of toluene is 934.2 kcal per mole. How much heat will be released during the combustion of 250.0 g of toluene? The formula for toluene is C6H5CH3 17. Copper has a density of 8.94 g/cm3 and a specific heat of 0.090 cal/g° C. A cube of copper is heated from 10.5° C
  67. 67. to 214° C. The cube of copper has dimensions of 5.00 cm. How much heat would the copper cube absorb? 18. The specific heat of water is 4.185 J/g° C (1.00 cal/g ° C). A piece of a pure metal with a mass of 24.0 g at a temperature of 45.0° C is added to 55 mL of water at 60.0° C. The final equilibrium temperature of the mixture is 95.4° C. Find the specific heat of the pure metal in both cal/g° C and J/g° C. 19. The specific heat of ice is 2.03 J/g° C. How much heat is needed to convert 550. g of ice at –15.0° C to 10.0° C? 20. What is the total amount of heat needed (in calories and joules) to convert 2.25 kg of ice at 0.0° C to steam at 200.0° C. 21. The specific heat of silicon is 0.057 cal/g° C and the density of silicon is 4.4 g/cm3. The volume of a cylinder formula is given as p r2L. The addition of 6000. calories raises the temperature of the silicon cylinder 55.5° C. Find the radius of the cylinder. The length of the cylinder is given as 6.00 cm. 22. A piece of metal with a mass of 75.5 g is heated to 84.5° C and added to 100.0 mL of water at 5.0° C. The final temperature of the mixture is 75.0° C. Find the specific heat of the metal. 23. Granite has a specific heat of 800. J/g° C. What mass of granite is needed to store 1.50 E 6 J of heat if the temperature of the granite is to be increased by 15.5° C?
  68. 68. 24. A 55 kg block of granite has an original temperature of 15.0° C. What will be the final temperature of this granite if 4.5 E 4 kJ of heat energy are added to the granite? The above curve attempts to demonstrate the addition of 100 calories of heat energy per minute to 1 gram of water. A thorough class discussion will attempt to identify the amount of time needed for each change to occur.
  69. 69. Knowledge of specific heat capacity, latent heat of fusion and vaporization is needed to determine these values.
  70. 70. The above chart shows a graph of the contraction of an air bubble in a capillary tube filled with oil. As the tube was cooled the length of the air bubble was measured and plotted as dots. When it could no longer be cooled to a lower temperature the graph line was extrapolated to find Absolute Zero (a temperature unobtainable in the actual world). Short Answer Questions Directions: Answer each question fully. Use complete sentences. Skip two lines of notebook paper (double double space if typed) between each question. Use only the front of paper.
  71. 71. 1. Explain latent heat (use water as an example). Give quantitative examples and clearly explain how intermolecular forces are involved during these changes. 2. Discuss how specific heat capacity might be used to identify an unknown metal sample. 3. Compare and contrast endothermic and exothermic chemical reactions. 4. Compare and contrast kinetic and potential energy, using as many practical examples as possible. 5. Using examples, explain fully elastic and inelastic collisions. Try to include some of the problems these concepts cause in science classes. Includemomentum in your discussion. 6. Describe how a liquid thermometer is made and how it could be calibrated without another thermometer 7. Explain how chemical and physical properties might be used to identify an unknown substance. Pick one common substance and give physical/chemical details for it. 8. Discuss the way in which gas exerts pressure on the walls of its container. 9. Explain the Kinetic-Molecular Theory of Gases as it relates to ordinary life. Include examples of Boyle’s Law and Charles’ Law in your explanation. 10. Compare intermolecular forces with chemical bonds within molecules. 11. Discuss atmospheric pressure and how it relates to gas laws, boiling points, and vapor pressures. 12. Explain how conduction, convection, and radiation might occur when using our
  72. 72. calorimeters. Compare our calorimeters with good Thermos bottles. 13. Explain fully melting-freezing point theory. 14. Explain fully boiling point theory and related phenomena such as elevation changes, pressure cookers, super heating, and vapor pressure. Molecular Motion Demonstrator This demonstration tool allows us to model a variety of atomic behavior. The following are notes (these will save the student having to make them during the demonstration and will allow them to study the modeling better. General observations 1) some molecules move faster than others, with constant random motion 2) molecules collide with each other and the walls 3) near elastic collisions 4) rarity of 3 body collisions 5) generally there is a large amount of space between the molecules 6) random motion in straight lines between collisions 7) pressure exerted on whatever they hit Diffusion - as the particles move across the barrier they demonstrate diffusion: (movement of molecules from one area to another with a net change in concentration) Temperature
  73. 73. 1) related to average speed of molecules 2) would 2 different gases at same temperature have same average speeds? 3) observe different speeds of different gas particles - > temperature relates to average K.E. of molecules 4) increase vibration rate: average speed increases, frequency of collisions increases, mean free path decreases Similarities with real world 1) model and real gases involve particles in continual, random motion, with straight line paths between collisions 2) particles in model occupy only a small fraction of total volume 3) changes in amount of movement is associated with change in temperature 4) rebounding forces of particles produce pressure in both model and real gases Dissimilarities with real world 1) model uses particles approximately ten million times the diameter of the particles in real gases 2) distance between collisions in real gases much greater than the distance traveled by plastic balls in relation to size 3) the speeds of the balls in the model are at most a few miles per hour while gas molecules travel at speeds of hundreds of miles per hour 4) real gas molecules collide elastically while our
  74. 74. model will run down due to friction is energy is not applied 5) real gases are three dimensional, model is two dimensional 6) most gas molecules are not spherical 7) spaces between particles in model taken up by air, in gases empty spaces occupies the space between molecules Liquids and Solids attractive forces between gas molecules are called van der Waal forces - positive charge of nucleus is not completely shielded by electron cloud so at short distances the nucleus may be attracted to electron cloud of another atom - these gas molecules may come together, slow down, and allow attractive forces to form Boyle’s Law Charles’ Law Relative Humidity When molecules from the vapor above a liquid surface impinge on the surface, they may be trapped there, so that a constant two-way traffic of molecules to and from the liquid occurs. If the density of the vapor above the liquid is sufficiently great, as many molecules return as leave it at any time, a situation that is described by saying
  75. 75. that the region is saturated with the substance. The higher the temperature, the greater the maximum vapor density: at 0° C the density of water vapor at saturation is 5 g/m3, at 20° C it is 17 g/m3, at 100° C it is 598 g/m3, and at 300° C it is all the way up to 45.6 kg/m3. If for any reason (such as a sudden drop in temperature) the vapor density exceeds the saturation value, condensation will be more rapid than evaporation until equilibrium is reestablished. It is for this reason that on a hot day moisture condenses on the outside of a glass that contains a cold drink. The relative humidity of a volume of air describes its degree of saturation with water vapor. Relative humidities of 0, 50%, and 100% mean respectively that no water vapor is present, that the air contains half as much moisture as the maximum possible, and that the air is saturated. On a hot day the evaporation of sweat from the skin is the chief means by which the human body dissipates heat, and a high relative humidity is uncomfortable because it impedes the process. A low relative humidity is also undesirable because it leads to the drying of the skin and mucous membranes. The regulation of relative humidity is as important a function of a heating or of an air- conditioning system as the regulation of temperature. We know that heating air decreases its relative humidity and cooling air increases its relative humidity. For instance, between 10° C and 20° C the saturated vapor density (which corresponds to 100% relative humidity) just about doubles. This means that if outside air at 10° C whose relative humidity is say, 70% is taken inside a house
  76. 76. and heated to 20° C, the relative humidity indoors will only be 35% since the actual vapor density stays the same. A way to humidify heated air in winter is clearly desirable. If the outside air is at 30° C with 70% relative humidity, then cooling it down to 24° C is enough to bring it to saturation, which is 100% relative humidity; further cooling will cause water to condense out. An air-conditioning system therefore should incorporate means to remove water vapor from the air being cooled. Relative humidity is water vapor density relative to saturation density. Changing the temperature of a body of air also changes its relative humiditiy. Principle of Heat Exchange - The Coffee-Cream Problem Newton's law of cooling: The rate of heat conduction is proportional to the temperature difference between an object and its surroundings. The Stefan-Boltzmann law of radiation: The rate of heat lost by radiation is proportional to the fourth power of the absolute temperature. The historic problem: Ah, you see, there is this business man who likes a large amount of cream in his coffee, and he wants the resultant mixture as hot as possible. (Alas, there is no microwave oven available).
  77. 77. He has just prepared his boiling coffee when he is called by the boss for a quick conference of ten minutes duration. The boss tolerates no coffee in his presence. What to do? To keep the coffee as hot as possible should he add the cream now or wait until after the conference? Which do you think he should do? _____________________________________ Using beakers for the coffee cups and water for the coffee and cream, design an experiment to test the problem. Use the computer and temperature probe to measure the temperature over time. How would you set up a graph(s) to help prove the best solution. Formulas K.E. = ½ m v2 P.E. = m g h Q = mass x D t x cp Q = mass x Hf Q = mass x Hv K = ° C + 273 ° C = K - 273 1 calorie = 4.185 joules
  78. 78. Pressure = Force area Temperature Scale Comparisons Fahrenheit Celsius Kelvin Boiling p. of 212° 100° 373 water Body 98.6° 37° 310 Temperature Freezing p. 32° 0° 273 water Coincidence - 40° - 40° 233 p. Absolute -273.16 - 460° 0 Zero ° Heat of Solution Reactions Lab Chemical and physical changes are usually accompanied by the liberation or absorption of energy. If energy is evolved, the reaction is said to be EXOTHERMIC. If the energy is absorbed, the reaction is ENDOTHERMIC.
  79. 79. Heat is a form of energy. The calorie or joule is the unit used to express heats of reaction. The calorie is defined as the amount of heat required to raise the temperature of 1 gram of water one Celsius degree. For conversions, 1 calorie = 4.185 joules. Energy may be transformed from one kind to another within an isolated (or closed) system but the total energy does not change. If the change in energy of a system can be measured and if this change is due solely to a chemical reaction, then the energy change must be equal to that of the chemical reaction itself. In this experiment a simple calorimeter will be used and the change in energy of the system will be measured by observing the temperature of a given weight of water before and after the reaction occurs. The specific heat capacity of water (i.e., the energy required to raise the temperature 1o C of 1 gram of the material) is very nearly 1.00 cal/goC for temperatures between 0o C and 100o C. Thus, if a calorimeter contained 100 grams of water at 23.0o C initially and after the reaction took place there were still 100. grams of water and the temperature was now 30.0o C, the energy liberated in the reaction would be Q (cal) = mass (g) x Dt (oC) x cp (cal/goC) 700. cal = 100. g x 7.00o C x 1.00 cal/goC This calculation assumes that no energy was required to raise the temperature of the calorimeter itself and that no