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UNIT 3 CONTENTS CHAPTER 5 Energy and Change CHAPTER 6 Rates of Chemical Reactions UNIT 3 PROJECT Developing a Bulletin About Catalysts and Enzymes e ——rra— UNIT 3 OVERALL EXPECTATIONS m What energy transformations and mechanisms are involved in chemical change? m What skills are involved in determining energy changes for physical and chemical processes and rates of reaction? m How do chemical technologies and processes depend on the energetics of chemical reactions? Look ahead to the project at the end of Unit 3. Start preparing for the project now by listing what you already know about catalysts and enzymes. Think about how catalysts and enzymes affect chemical reactions. As you work through the unit, plan how you will investigate and present a bulletin about the uses of catalysts and enzymes in Canadian industries. In the nineteenth century, railway tunnels were blasted through the Rocky Mountains to connect British Columbia with the rest of Canada. Workers used nitroglycerin to blast through the rock. This compound is so unstable, however, that accidents were frequent and many workers died. Alfred B. Nobel found a way to stabilize nitroglycerin, and make it safer to use, when he invented dynamite. What makes nitroglycerin such a dangerous substance? First, nitro- glycerin, C3H5(NQO3)3(, gives off a large amount of energy when it decomposes. In fact, about 1500 k] of energy is released for every mole of nitroglycerin that reacts. Second, the decomposition of nitroglycerin occurs very quickly—in a fraction of a second. This fast, exothermic reaction is accompanied by a tremendous shock wave, which is caused by the expansion of the gaseous products. Finally, nitroglycerin is highly shock- sensitive. Simply shaking or jarring it can cause it to react. Thus, nitroglycerin’s explosive properties are caused by three fac- tors: the energy that is given off by its decomposition, the rate at which the reaction occurs, and the small amount of energy that is needed to initiate the reaction. In this unit, you will learn about the energy and rates of various chemical reactions. Energy Changes and Rates of Reaction

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The Energy of Physical, Chemical, and Nuclear Processes Most physical changes, chemical reactions, and nuclear reactions are accompanied by changes in energy. These energy changes are crucial to life on Earth. For example, chemical reactions in your body generate the heat that helps to regulate your body temperature. Physical changes, such as evaporation, help to keep your body cool. On a much larger scale, there would be no life on Earth without the energy from the nuclear reactions that take place in the Sun. The study of energy and energy transfer is known as thermodynamics. Chemists are interested in the branch of thermodynamics known as thermochemistry: the study of energy involved in chemical reactions. In order to discuss energy and its interconversions, thermochemists have agreed on a number of terms and definitions. You will learn about these terms and definitions over the next few pages. Then you will examine the energy changes that accompany chemical reactions, physical changes, and nuclear reactions. Studying Energy Changes The law of conservation of energy states that the total energy of the universe is constant. In other words, energy can be neither destroyed nor created. This idea can be expressed by the following equation: AEuniverse =0 Energy can, however, be transferred from one substance to another. It can also be converted into various forms. In order to interpret energy changes, scientists must clearly define what part of the universe they are dealing with. The system is defined as the part of the universe that is being studied and observed. In a chemical reaction, the system is usually made up of the reactants and products. By contrast, the surroundings are everything else in the universe. The two equations below show the relationship between the universe, a system, and the system’s surroundings. Universe = System + Surroundings AEyniverse = AEsystem + AEsurroundings =0 From the relationship, we know that any change in the system is accompanied by an equal and opposite change in the surroundings. AEsystem = _AEsurroundings Look at the chemical reaction that is taking place in the flask in Figure 5.1. A chemist would probably define the system as the contents of the flask—the reactants and products. Technically, the rest of the universe is the surroundings. In reality, however, the entire universe changes very little when the system changes. Therefore, the surroundings are considered to be only the part of the universe that is likely to be affected by the energy changes of the system. In Figure 5.1, the flask, the lab bench, the air in the room, and the student who is carrying out the reaction all make up the surroundings. The system is more likely to significantly influence its immediate surroundings than, say, a mountaintop in Japan (also, technically, part of the surroundings). Section Preview/ Specific Expectations In this section, you will m write thermochemical equa- tions, expressing the energy change as a heat term in the equation or as AH m represent energy changes using diagrams m compare energy changes that result from physical changes, chemical reac- tions, and nuclear reactions m communicate your under- standing of the following terms: thermodynamics, thermochemistry, law of conservation of energy, system, surroundings, heat (Q), temperature (T), enthalpy (H), enthalpy change (AH), endothermic reaction, exothermic reac- tion, enthalpy of reaction (AH. ), standard enthalpy of reaction (AH’ ), thermo- chemical equation, mass defect, nuclear binding energy, nuclear fission, nuclear fusion i EETZXEN The solution in the flask is the system. The flask, the laboratory, and the student are the surroundings. Chapter 5 Energy and Change - MHR 221

Heat and Temperature Heat, Q, refers to the transfer of kinetic energy. Heat is expressed in the same units as energy—joules (J). Heat is transferred spontaneously from a warmer object to a cooler object. When you close the door of your home on a cold day to “prevent the cold from getting in,” you are actually preventing the heat from escaping. You are preventing the kinetic energy in your warm home from transferring to colder objects, including the cold air, outside. Temperature, 7, is a measure of the average kinetic energy of the particles that make up a substance or system. You can think of temperature as a way of quantifying how hot or cold a substance is, relative to another substance. K B uc 110 :(\ GETEEE] Celsius degrees and Kelvin degrees are the same size. The Kelvin scale begins at absolute zero. This is the temperature at which the particles in a substance have no kinetic energy. Therefore, Kelvin temperatures are never negative. By contrast, 0°C is set at the melting point of water. Celsius temperatures can be positive or negative. Temperature is measured in either Celsius degrees (°C) or kelvins (K). The Celsius scale is a relative scale. It was designed so that water’s boiling point is at 100°C and water’s melting point is at 0°C. The Kelvin scale, on the other hand, is an absolute scale. It was designed so that 0 K is the temperature at which a substance possesses no kinetic energy. The relationship between the Kelvin and Celsius scales is shown in Figure 5.2, and by the following equation. Temperature in Kelvin degrees = Temperature in Celsius degrees + 273.15 Enthalpy and Enthalpy Change Chemists define the total internal energy of a substance at a constant pressure as its enthalpy, H. Chemists do not work with the absolute enthalpy of the reactants and products in a physical or chemical process. Instead, they study the enthalpy change, AH, that accompanies a process. That is, they study the relative enthalpy of the reactants and products in a system. This is like saying that the distance between your home and your school is 2 km. You do not usually talk about the absolute position of your home and school in terms of their latitude, longitude, and elevation. You talk about their relative position, in relation to each other. The enthalpy change of a process is equivalent to its heat change at constant pressure. 222 MHR - Unit 3 Energy Changes and Rates of Reaction

1 A Hj(g) + 202(g) (reactants) — | AH® =-285.8 kJ enthalpy, H — H,0¢) - (products) IS In an exothermic reaction, the enthalpy of the system decreases as energy is released to the surroundings. A MgO(s) + COZ(g) (products) " \I\l’ AH"=117.3 k] enthalpy, H e MgCO3(S) (reactants) XX In an endothermic reaction, the enthalpy of the system increases as heat energy is absorbed from the surroundings. You can also indicate the enthalpy of reaction as a separate expression beside the chemical equation. For exothermic reactions, AH" is always negative. Hag) + %Oz(g) — Hy0) AH'xy =-285.8K] A third way to represent the enthalpy of reaction is to use an enthalpy diagram. Examine Figure 5.4 to see how this is done. Representing Endothermic Reactions The endothermic decomposition of solid magnesium carbonate produces solid magnesium oxide and carbon dioxide gas. For each mole of magne- sium carbonate that decomposes, 117.3 kJ of energy is absorbed. As for an exothermic reaction, there are three different ways to represent the enthalpy change of an endothermic reaction. You can include the enthalpy of reaction as a heat term in the chemical equation. Because heat is absorbed in an endothermic reaction, the heat term is included on the reactant side of the equation. 117.3 kJ + MgCO3() — MgOy) + COxg) You can also indicate the enthalpy of reaction as a separate expression beside the chemical reaction. For endothermic reactions, the enthalpy of reaction is always positive. MgCOj35) = MgOg) + COz(g) AH = 117.3 K] Finally, you can use a diagram to show the enthalpy of reaction. Figure 5.5 shows how the decomposition of solid magnesium carbonate can be represented graphically. Stoichiometry and Thermochemical Equations The thermochemical equation for the decomposition of magnesium carbonate, shown above, indicates that 117.3 k] of energy is absorbed when one mole, or 84.32 g, of magnesium carbonate decomposes. The decomposition of two moles of magnesium carbonate absorbs twice as much energy, or 234.6 kJ. MgCO35) — MgOgs) + COyg AH i, = 117.3 K] 2MgCO3(5) — 2MgO(g) + ZCOZ(g) AH’, = 234.6 k] Enthalpy of reaction is linearly dependent on the quantity of products. That is, if the amount of products formed doubles, the enthalpy change also doubles. Figure 5.6 shows the relationship between the stoichiometry of a reaction and its enthalpy change. Because of this relationship, an exothermic reaction that is relatively safe on a small scale may be extremely dangerous on a large scale. One of the jobs of a chemical engineer is to design systems that allow exothermic reactions to be carried out safely on a large scale. For example, the blast furnaces used in steel making must withstand temperatures of up to 2000°C, produced by the exothermic combustion reaction of coal with oxygen. amount of varies directly with amount of varies directly with | heat absorbed compound A compound B or released (mol) (factor: molar ratio) (mol) (factor: AHxn) (kdJ) A A A A XX This diagram summarizes the relationship between the stoichiometry of a reaction and AH. 224 MHR - Unit 3 Energy Changes and Rates of Reaction

I Sample Problem Stoichiometry and Thermochemical Reactions Problem Aluminum reacts readily with chlorine gas to produce aluminum chloride. The reaction is highly exothermic. 2Al) + 3Clyg — 2AlCl3) AH'rxn = —1408 k] What is the enthalpy change when 1.0 kg of Al reacts completely with excess Cl,? What Is Required? You need to calculate the enthalpy change, AH, when the given amount of Al reacts. What Is Given? You know the enthalpy change for the reaction of two moles of Al with one mole of Cl;. From the periodic table, you know the molar mass of Al. ZAl(S) + 3C12(g) - 2A1C13(S] AH’ ., = —1408 k] M1 = 26.98 g/mol Plan Your Strategy Convert the given mass of Al to moles. The enthalpy change is linearly dependent on the quantity of reactants. Therefore, you can use a ratio to determine the enthalpy change for 1.0 kg of Al reacting with Cl,. Act on Your Strategy Determine the number of moles of Al in 1 kg. Remember to convert to grams. m nmol Al = AL Mai _ 1.0x10%g ~ 26.98 g/mol = 37 mol Use ratios to compare the reference reaction with the known enthalpy change (AH,) to the reaction with the unknown enthalpy change (AH;). AH, ny; mol Al AH; n;mol Al AH, 37 mol Al —1408k] 2 mol Al AH, = -2.6 x 10* kJ Check Your Solution The sign of the answer is negative, which corresponds to an exothermic reaction. The 1 kg sample contained about 20 times more moles of Al. Therefore, the enthalpy change for the reaction should be about 20 times greater, and it is. N \ Chapter 5 Energy and Change * MHR 225

100°C until all the water has been vaporized. If you add heat to the vapour, the temperature of the vapour will increase steadily. When you heat ice that is colder than 0°C, a similar process occurs. The temperature of the ice increases until it is 0°C (the melting point of water). If you continue to add heat, the ice remains at 0°C but begins to melt, as the bonds between the water molecules in the solid state begin to break. Figure 5.7 shows the relationship between temperature and heat for a solid substance that melts and then vaporizes as heat is added to it. A gas boiling | ) point : ! g liquid i i As heat is added to a Q ! ! substance, the temperature of the £ melting | AHyap | substance steadily increases until it F point | . i . | reaches its melting point or boiling solid i i i i i point. The temperature then remains ! AH met ! ! ! steady as the substance undergoes ' ' ' : > a phase change. Heat added » You can represent the enthalpy change that accompanies a phase change—from liquid to solid, for example—just like you represented the enthalpy change of a chemical reaction. You can include a heat term in the equation, or you can use a separate expression of enthalpy change. For example, when one mole of water melts, it absorbs 6.02 kJ of energy. \-J CHEM The process of melting is also H,Or) — HoO) AH =6.02K] known as fusion. Therefore, Normally, however, chemists represent enthalpy changes associated with you will sometimes see the phase changes using modified AH symbols. These symbols are described enthalpy of melting referred to below. as the enthalpy of fusion. N Y e enthalpy of vaporization, AH,,y : the enthalpy change for the phase change from liquid to gas ¢ enthalpy of condensation, AH,opnq: the enthalpy change for the phase change of a substance from gas to liquid e enthalpy of melting, AHpe1: the enthalpy change for the phase change of a substance from solid to liquid e enthalpy of freezing, AHy,: the enthalpy change for the phase change of a substance from liquid to solid Vaporization and condensation are opposite processes. Thus, the enthalpy changes for these processes have the same value but opposite signs. For example, 6.02 k] is needed to vaporize one mole of water. Therefore, 6.02 kJ of energy is released when one mole of water freezes. AI_Ivap = —AH_ong Similarly, melting and freezing are opposite processes. AHmelt = ~AHge Several enthalpies of melting and vaporization are shown in Table 5.1. Notice that the same units (kJ/mol) are used for the enthalpies of melting, vaporization, condensation, and freezing. Also notice that energy changes associated with phase changes can vary widely. Chapter 5 Energy and Change « MHR 227

Table 5.1 Enthalpies of Melting and Vaporization for Several Substances halpy of melti Ipy of vaporizati AH i (kJ/mol) AH 5, (kJ/mol) argon 1.3 6.3 diethyl ether 7.3 29 ethanol 5.0 40.5 mercury 23.4 29 methane 8.9 0.94 sodium chloride 27.2 207 water 6.02 40.7 Hot Packs and Cold Packs: Using the Energy of Physical Changes You just learned about the enthalpy changes that are associated with phase changes. Another type of physical change that involves a heat transfer is dissolution. When a solute dissolves in a solvent, the enthalpy change that occurs is called the enthalpy of solution, AHg,. Dissolution can be either endothermic or exothermic. Manufacturers take advantage of endothermic dissolution to produce cold packs that athletes can use to treat injuries. One type of cold pack contains water and a salt, such as ammonium nitrate, in separate compartments. When you crush the pack, the membrane that divides the compartments breaks, and the salt dissolves. This dissolution process is endothermic. It absorbs heat for a short period of time, so the cold pack feels cold. Figure 5.8 shows how a cold pack works. ~ soluble salt membrane of water pack A typical cold pack has two separate chambers. One chamber contains a salt. The other chamber contains water. Crushing the pack allows the salt to dissolve in the water—an endothermic process. 228 This person’s shoulder was injured. Using a cold pack helps to reduce the inflammation of the joint. Some types of hot packs are constructed in much the same way as the cold packs described above. They have two compartments. One compartment contains a salt, such as calcium chloride. The other compartment contains water. In hot packs, however, the dissolution process is exothermic. It releases heat to the surroundings. MHR « Unit 3 Energy Changes and Rates of Reaction

GETEXX] This graph shows the relative stability of nuclei. It indicates whether nuclei are more likely to split or fuse in a nuclear reaction. Notice that the helium-4 nucleus is unusually stable. Electronic Learning Partner To learn more about nuclear fission, go to the Chemistry 12 Electronic Learning Partner. 9 34g OFg 84Kr 119g,, S 812¢ UL § region of very < 74, stable nuclei 238 c 4He 5 5 67 3 E 5| 6L 2 fusion fission > 4 — - 2 (] S 3l 3y 2 5| °He 2 . |, m 11 °H 0 I I I I I I I I I I I I I 20 40 60 80 100 120 140 160 Mass number, A 180 200 220 240 260 The difference between the nuclear binding energy of the reactant nuclei and the product nuclei represents the energy change of the nuclear reaction. Nuclear Fission A heavy nucleus can split into lighter nuclei by undergoing nuclear fission. Nuclear power plants use controlled nuclear fission to provide energy. Uncontrolled nuclear fission is responsible for the massive destructiveness of an atomic bomb. The most familiar fission reactions involve the splitting of uranium atoms. In these reactions, a uranium-235 atom is bombarded with neutrons. The uranium nucleus then splits apart into various product nuclei. Two examples of fission reactions that involve uranium-235 are shown in Figure 5.10. ) > + energy ;n zggu > \\ i > + energy L\ ;n ¥ zggu > ETEEETY Uranium can undergo fission in numerous different ways, producing various product nuclei. Two examples are shown here. 230 MHR - Unit 3 Energy Changes and Rates of Reaction

Fission reactions produce vast quantities of energy. For example, when one mole of uranium-235 splits, it releases 2.1 x 10'3 J. By contrast, when one mole of coal burns, it releases about 3.9 x 10° J. Thus, the combustion of coal releases about five million times fewer joules of energy per mole than the fission of uranium-235. Nuclear Fusion Two smaller nuclei can fuse to form a larger nucleus, in what is called a nuclear fusion reaction. You and all other life on Earth would not exist without nuclear fusion reactions. These reactions are the source of the energy produced in the Sun. One example of a fusion reaction is the fusion of deuterium and tritium. 2H+3H - jHe + in The seemingly simple reaction between deuterium and tritium produces 1.7 x 10'? J of energy for each mole of deuterium. This is about 10 times fewer joules of energy than are produced by the fission of one mole of uranium. It is still, however, 500 000 times more energy than is produced by burning one mole of coal. Scientists are searching for a way to harness the energy from fusion reactions. Fusion is a more desirable way to produce energy than fission. The main product of fusion, helium, is relatively harmless compared with the radioactive products of fission. Unfortunately, fusion is proving more difficult than fission to harness. The fission of one mole of uranium-235 produces more energy than the fusion of one mole of deuterium with one mole of tritium, What if you compare the energy that is produced in terms of mass of reactants? Calculate a ratio to compare the energy that is produced from fusion and fission, per gram of fuel. What practical consequences arise from your result? ) , _ 1024 ) — Fusion will not proceed at a reasonable rate without an initial i t of This i t bl — daily solar energy enormous initial input of energy. This is not a problem 102" ) - falling on Earth in the core of the Sun, where the temperature ranges | energy of strong earthquake W\/\[W\M from 7 500 000°C to 15 000 000°C. It is a problem in 108y — industry. Scientists are working on safe and economical daily electrical output of . ] .. Canadian dams at Niagara Falls ways to provide the high-temperature conditions that 108y — . — 1000 t of coal burned are needed to make fusion a workable energy source. 1012J — Comparing the Energy of Physical, Chemical, o5, - 1tof TNT exploded 4 and Nuclear Processes | - 7KW h of electrical energy 7 In thi i 1 d that physical ch 64 — n this section, you learned that physical changes, 1080 9\ , , . ] ] heat released from combustion chemical reactions, and nuclear reactions all involve of 1 mol glucose energy changes. You also learned that the energy 10°J — changes have some striking differences in magnitude. heat required to boil 1 mol of water . . 0 — Figure 5.11 shows energy changes for some physical, 10°J chemical, and nuclear processes. Some other interesting , energy statistics are included for reference. 107J 106J 102 J —— heat absorbed during division of one bacterial cell 10-12 ) —f— energy from fission of one 235y atom £84 s 107150 — e GETEEREN The energy changes of physical, nuclear, and chemical ‘(&f?": processes vary widely. In general, however, chemical reactions are 108y — o ' associated with greater energy changes than physical changes. §5s Nuclear reactions are associated with far greater energy changes than 10-21 —— average kinetic energy of a molecule chemical reactions. in air at 300 K Chapter 5 Energy and Change - MHR 231

(b) What is the enthalpy change for the reaction of 46.7 g of graphite with excess calcium oxide? (e} 1.38 x 10?* formula units of calcium oxide react with excess graphite. How much energy is needed? © © Acetylene, C,H,, undergoes complete combustion in oxygen. Carbon dioxide and water are formed. When one mole of acetylene undergoes complete combustion, 1.3 x 102 k] of energy is released. (a) Write a thermochemical equation for this reaction. (b) Draw a diagram to represent the thermochemical equation. (¢) How much energy is released when the complete combustion of acetylene produces 1.50 g of water? @ © Write an equation to represent each phase change in Table 5.1 on page 228. Include the enthalpy change as a heat term in the equation. @ @D When one mole of gaseous water forms from its elements, 241.8 kJ of energy is released. In other words, when hydrogen burns in oxygen or air, it produces a great deal of energy. Since the nineteenth century, scientists have been researching the potential of hydrogen as a fuel. One way in which the energy of the combustion of hydrogen has been successfully harnessed is as rocket fuel for aircraft. (a) Write a thermodynamic equation for the combustion of hydrogen. (b) Describe three reasons why hydrogen gas is a desirable rocket fuel. (c) Suggest challenges that engineers might have had to overcome in order to make hydrogen a workable rocket fuel for aircraft. (d) Use print and electronic resources to find out about research into hydrogen as a fuel. Create a time line that shows significant events and discoveries in this research. © @D A healthy human body maintains a temperature of about 37.0°C. Explain how physical, chemical, and nuclear processes all contribute, directly or indirectly, to keeping the human body at a constant temperature. Chapter 5 Energy and Change * MHR 233

Section Preview/ Specific Expectations In this section, you will m determine the heat that is produced by a reaction using a calorimeter, and use the data obtained to calcu- late the enthalpy change for the reaction m communicate your under- standing of the following terms: specific heat capacity (c), heat capacity (C), calorimeter, coffee-cup calorimeter, constant- pressure calorimeter GETEXEE] On a sunny day, you would probably prefer to siton a bench made of wood rather than a bench made of aluminum. Wood has a higher heat capacity than aluminum. Therefore, more heat is needed to increase its temperature. CONCEPT CHECK Explain why water is some- times used as a coolant for automobile engines. Determining Enthalpy of Reaction by Experiment Chemical and physical processes, such as the ones you studied in section 5.1, are associated with characteristic enthalpy changes. How do chemists measure these enthalpy changes? To measure the enthalpy of a chemical or physical process, chemists insulate the system from the surroundings. They can then determine the heat change by measuring the temperature change of the system. What is the relationship between the heat change and the temperature change? As you learned in previous science courses, each substance has a charac- teristic property that dictates how its temperature will change when heat is lost or gained. Specific Heat Capacity The amount of energy that is needed to raise the temperature of one gram of a substance 1°C (or 1 K) is the specific heat capacity, c, of the substance. Specific heat capacity is usually expressed in units of J/g« “C. The specific heat capacities of several substances are given in Table 5.2. Figure 5.12 shows that you can often predict the relative specific heat capacities of familiar substances. Table 5.2 Specific Heat Capacities of Selected Substances (TR (T CE T T LT AW T IO 25"(:)‘ Element aluminum 0.900 carbon (graphite) 0.711 hydrogen 14.267 iron 0.444 Compound ammonia (liquid) 4.70 ethanol 2.46 ethylene glycol 2.42 water (liquid) 4.184 Other material air 1.02 concrete 0.88 glass 0.84 wood 1.76 You can use the specific heat capacity of a substance to calculate the amount of energy that is needed to heat a given mass a certain number of degrees. You can also use the specific heat capacity to determine the amount of heat that is released when the temperature of a given mass decreases. The specific heat capacity of liquid water, as shown in Table 5.2, is 4.184 J/g« °C. This relatively large value indicates that a considerable amount of energy is needed to raise or lower the temperature of water. 234 MHR - Unit 3 Energy Changes and Rates of Reaction

www.mcgrawhill.ca/links/ chemistry12 The famous French scientist Antoine Laviosier (1743-1794) is considered by many to be the first modern chemist. Lavoisier created a calori- meter to study the energy that is released by the metabolism of a guinea pig. To learn about Lavoisier’s experiment, go to the web site above and click on Web Links. What do you think about using animals in experiments? Write an essay to explain why you agree or disagree with this practice. Measuring Heat Transfer in a Laboratory From previous science courses, you will probably remember that a calorimeter is used to measure enthalpy changes for chemical and physical reactions. A calorimeter works by insulating a system from its surroundings. By measuring the temperature change of the system, you can determine the amount of heat that is released or absorbed by the reaction. For example, the heat that is released by an exothermic reaction raises the temperature of the system. Qreaction = _Qinsulated system In your previous chemistry course, you learned about various types of calorimeters. For instance, you learned about a bomb calorimeter, which allows chemists to determine energy changes under conditions of constant volume. In section 5.1, however, you learned that an enthalpy change repre- sents the heat change between products and reactants at a constant pressure. Therefore, the calorimeter you use to determine an enthalpy change should allow the reaction to be carried out at a constant pressure. In other words, it should be open to the atmosphere. To determine enthalpy changes in high school laboratories, a coffee-cup calorimeter provides fairly accurate results. A coffee-cup calorimeter is composed of two nested polystyrene cups (“coffee cups™). They can be placed in a 250 mL beaker for added stability. Since a coffee-cup calorimeter is open to the atmosphere, it is also called a constant-pressure calorimeter. As with any calorimeter, each part of the coffee-cup calori-meter has an associated heat capacity. Because these heat capacities are very small, however, and because a coffee-cup calorimeter is not as accurate as other calorimeters, the heat capacity of a coffee-cup calorimeter is usually assumed to be negligible. It is assumed to have a value of 0 J/°C. Using a Calorimeter to Determine the Enthalpy of a Reaction A coffee-cup calorimeter is well-suited to determining the enthalpy changes of reactions in dilute aqueous solutions. The water in the calorimeter absorbs (or provides) the energy that is released (or absorbed) by a chemical reaction. When carrying out an experiment in a dilute solution, the solution itself absorbs or releases the energy. You can calculate the amount of energy that is absorbed or released by the solution using the equation mentioned earlier. Q=mec+AT The mass, m, is the mass of the solution, because the solution absorbs the heat. When a dilute aqueous solution is used in a calorimeter, you can assume that the solution has the same density and specific heat capacity as pure water. As you saw above, you can also assume that the heat capacity of the calorimeter is negligible. In other words, you can assume that all the heat that is released or absorbed by the reaction is absorbed or released by the solution. The following Sample Problem illustrates how calorimetry can be used to determine AH of a chemical reaction. 236 MHR - Unit 3 Energy Changes and Rates of Reaction