Simple Experiments in Thermochemistry

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1 Simple Experiments in Thermochemistry Purpose: To demonstrate the law of conservation of energy and propose a method for making a chemical heat pack using the heats of solution of sodium bicarbonate and calcium chloride. Introduction: Thermochemical measurements are a simple yet precise way to watch energy changes in chemical reactions as well as obtain a good conceptual feel of the meaning of temperature and heat capacity. Almost every chemical reaction that occurs will release energy (an exothermic reaction) or need energy (an endothermic reaction) as reactants are converted into products. However, understanding exactly how and why these energy changes come about is often a source of confusion for many students. For students to obtain a complete grasp of the role of energy in chemical reactions, it is essential that they understand the roles of temperature and heat capacity. Essentially, temperature is measuring the kinetic energy of a set of particles such that, on a molecular level, particles with high temperature are moving around quickly, while identical particles with low temperature have a much lower speed. An excellent analogy is modeling particles as a set of billiard balls that can float in the room and that can also collide with the walls of the room or other particles. Such collisions can exchange speeds between the particles in the container and hence transfer energy from one particle to another. However, at a given temperature, the average energy of the particles is constant. A thermometer is simply a device that receives some of this energy when placed in a container. As this energy is transferred to atoms in the thermometer, they move more violently, with the macroscopic effect being that the average distance between the atoms or molecules increases. We see this increase as expansion of the material in the thermometer What happens when two objects of different temperature are allowed to interact with one another? We know that when you mix a sample of hot water with a sample of cold water, warm water will result. If we think about this on a molecular level, this makes perfect sense because the hot molecules were moving quickly and at some point will collide with a slow moving cold molecule. As a consequence of the collision, speeds will be changed and you will likely get two molecules that are moving closer to the same speed. These myriad collisions will cause the average speed of the new mixture of molecules to be something in between the hot and cold molecules. Consequently, the temperature will drop, because the average speed of the molecules has decreased. We call this transfer of energy heat. Of course, molecules can store energy in other ways besides simply moving around. Pieces of molecules can vibrate or rotate as well. However, if energy is placed into these motions, it has been stored by the molecule in different ways and hence will not show up in the temperature of the substance. Indeed, we know that if I place a metal pan on a stove or a metal pan filled with some water and add energy to both via heating for the

2 same period of time, the temperature of the metal will rise much faster than the temperature of the metal and water combination. This is a direct result of the fact that water is an excellent energy storer. Specifically, it can store the energy of collisions in more ways than just moving quickly from point A to point B. Numerically, we can assign a numerical value to this ability to store energy and we call it the specific heat, denoted by the letter c. The specific heat can be easily measured, as it is defined as the amount of heat energy needed to raise the temperature of a substance one degree Celsius (or Kelvin). Thus, to measure c, we simply need to add a known amount of heat to a substance and watch how much its temperature rises. Mathematically, the specific heat can then be obtained by: heat added c = (1.1) mass * temperature change or, using the proper symbols for heat, mass, and temperature: q c = m*!t (1.2) A large c means a substance can absorb a lot of heat energy internally and not change its speed substantially. In contrast substances with small c s can only store energy by moving quickly. Water has one of the largest specific heats one will see, with a value of 4.18 J/(gram* o C). In contrast the specific heats of most metals are less than 1 J/(gram* o C). Measuring the specific heat of a substance is a common IPC or first-year chemistry lab experiment. Equation (1.2) can be rearranged to calculate the amount of heat a substance has gained via a chemical or physical process: q = m*c*!t (1.3) This is the essential equation of calorimetry, the process of measuring the energy changes of chemical reactions. If we perform a chemical reaction in an environment such as water, simply observing the temperature change of the water gives us energetic insight into what is happening in the chemical reaction. However, can we be sure that the energy gained or lost by the reaction is accurately reflected by the change of temperature of the water? Ideally the answer to this question is yes as conservation of energy guarantees that the total energy change of the reaction plus the environment is zero. However, calorimetry allows us to test conservation of energy very simply while reinforcing a molecular level picture of what occurs while hot and cold water are mixed. Using the small-scale calorimeters (more information about these is attached in the Flinn document), one can quantitatively demonstrate this critical law while gaining practice with (1.3).

3 EXPERIMENT I: VERIFYING THE LAW OF CONSERVATION OF ENERGY Purpose: To Verify the Law of conservation of energy while gaining experience with the relationship between temperature, specific heat, and energy change Materials: Two small scale calorimeters Two Styrofoam cups Hot and Cold Tap Water One 20 ml syringe One digital thermometer Digital Balances (0.01 g) Procedure: 1) Label the two calorimeters and the two Styrofoam cups hot and cold with a pen or pencil 2) Place cool tap water (room temp water is fine) on the cold cup. Place heated water in the hot cup. Fill each cup about halfway with water. A great way to make the hot water is to place the filled hot cup in the microwave for sec. Make sure that the temperature of the hot cup is no greater than 50 degrees C 3) Find the mass of each calorimeter to the nearest 0.01 g. Add 5-10 ml of the cold water to the cold calorimeter and remass to 0.01 g. Do the same with the hot calorimeter, adding 5-10 ml of hot water. However, do not exceed 15 ml of total volume between the two calorimeters (doing so will cause them to overflow when they mix). Record the mass of the hot calorimeter and water. 4) Using the digital thermometer, obtain the stable temperature of both the cold and hot water in the respective calorimeters, wiping off the thermometer with a paper towel between checks. Record the temperature of the room temperature water first and then record the temperature of the hot water. 5) As soon as you have a stable temperature recorded for the hot water, pour the water from the cold calorimeter to the hot calorimeter. Make sure every drop is transferred. 6) Carefully stir the mixture with the thermometer and record the stable final temperature of the water. Be careful not to poke or damage the calorimeter. 7) Repeat the experiment two more times, making sure you dry the calorimeters with a paper towel between each run. Data Table I: Hot and Cold Water Calorimeter Data Trial I Trial II Trial III Mass of Hot Water Calorimeter (g) Mass of Cold Water Calorimeter (g) Mass of Calorimeter + Hot Water (g) Mass of Calorimeter + Cold Water (g) Initial Temperature of Cold Water ( o C) Initial Temperature of Hot Water ( o C) Final Temperature of mixture ( o C)

4 Calculations: (All calculations are summarized in Table II). Use the space below the table to show a sample calculation of each type and record your answers in the appropriate spot in the table Table II: Energy Conservation Results Trial I Trial II Trial III Mass of Cold Water (g) Mass of Hot Water (g) Temperature Change of Cold Water ( o C) Temperature Change of Hot Water ( o C) Energy Gained by Cold Water (J) Energy Lost by Hot Water (J) Sum of Energy Lost and Gained. (J) Percentage of Energy Lost 1) Calculate the mass of the cold water and the mass hot water in each calorimeter by subtracting the mass of each calorimeter + water from the mass of each calorimeter. 2) Calculate the temperature change of the water in each calorimeter by subtracting the temperature of the hot water from the temperature of the mixed water, and the temperature of the mixed water from the temperature of the cold water. Note that your change in temperature for the hot water will be negative

5 3) Using q=m*c*δt, and the fact that the specific heat of water is 4.18 J/(g* o C), calculate the energy change of the cold water. 4) Using q=m*c*δt, calculate the energy change of the warm water. Again, this change will be negative. 5) Add the energy gained by the cold water to the energy lost (again negative) by the hot water. 6) Calculate the percentage of energy lost by dividing the sum of the energy lost and gained by the energy change of the cold water and multiplying by 100 to convert to a percent.

6 Discussion: 1) Based on your data would you say that energy is conserved when the hot water and cold water are mixed? Explain 2) It is very likely that the sum of the energy lost and gained was not zero. Where do you think this energy went? 3) Do you think you would have seen the same results in this experiment if you had the solutions in aluminum soft drink cans as opposed to the Styrofoam calorimeters? Explain 4) What do you think happens to the atoms of aluminum in a soft drink can when you add hot water to the can? Illustrate with a diagram as part of your explanation

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