12.3 Heats of Reaction

Transcription

1 12.3 Heats of Reaction All chemical reactions involve energy changes. Some reactions like combustion (burning) are obviously exothermic. You can feel the heat and see the light emitted from a burning campfire or fireplace. In an exothermic reaction, energy is a product just like the chemical products of the reaction. Exothermic reactions always release energy into the surroundings: reactants products energy Endothermic reactions are less familiar in everyday life, but not unknown. If you put an Alka-Seltzer tablet in a glass of water, a chemical reaction takes place, as evidenced by the gas bubbles produced. If you feel the glass, you will notice that it feels cold. The reaction is absorbing heat from the surroundings (water, glass, air). In an endothermic reaction, energy is consumed like a reactant. Endothermic reactions always absorb energy from the surroundings: reactants energy products thermochemical equation: a balanced chemical equation that includes the heat transferred to or from the surroundings heat of reaction: the heat transferred in a reaction based on the amounts given by the coefficients of the balanced chemical equation (in units of, for example, kj) ar heat of reaction: the quantity of heat transferred in a reaction per e of a specified substance (in units of, for example, kj/); represented by H r You have used calorimetry to determine the specific heat of reaction for a particular substance. Now we can use this experimental value to write a chemical equation that includes the heat transferred. This chemical equation is called a thermochemical equation, and the quantity of heat transferred is called the heat of reaction. A chemical equation communicates the relative amounts, in es, of all reactants and products. To be able to relate experimentally measured specific heats to chemical equations, the energy must be expressed as a quantity of heat per e, such as kilojoules per e; this quantity is called the ar heat of reaction, H r, for a particular substance. For example, the specific heat of combustion of octane, from a calorimetry experiment, is an exothermic kj/g. How can this be converted into a ar heat of combustion? First, a conversion factor to convert kilojoules per gram into kilojoules per e must be found. The mass in grams must be converted into an amount in es. As you will recall from Unit 2, the conversion factor that was created by chemists to complete this conversion is the ar mass of the chemical. Using the ar mass as a conversion factor, we can convert the specific heat of combustion of octane, h c, into the ar heat of combustion of octane, : = kj g g 1 C 8 H 18(l) C 8 H 18(l) = kj This means that the combustion of one e of octane releases 5074 kj of heat. Notice how the units cancel in the above calculation. Cancelling units is always a good procedure to check to see if you are calculating correctly. Molar heats of reaction must always specify the type of reaction (in this case the subscript used is c for combustion) and the substance reacted (in this case octane, C 8 H 18(l) ). 582 Chapter 12

2 12.3 Sample Problem 1 Ethane is the second largest component of natural gas. If its specific heat of combustion is kj/g, what is the ar heat of combustion of ethane? Solution h c = kj/g C 2 H 6(g) = kj g g 1 C 2 H 6(g) C 2 H 6(g) = kj The ar heat of combustion of ethane is an exothermic 1560 kj/. Thermochemical Equations To combine a calculated ar heat of reaction with a chemical equation, you need to pay particular attention to the balancing of the equation. For example, the ar heat of combustion of octane (a major component of gasoline) is 5074 kj/. Because combustion reactions are always exothermic, this means that 5074 kj of heat is transferred to the surroundings during the combustion of one e of octane. If we balance the chemical equation using 1 as the coefficient for octane, we can write the heat of reaction immediately without any calculations: 1 C 8 H 18(l) O 2 2(g) 8 CO 2(g) + 9 H 2 O (g) kj This equation is read as one e of octane reacts with twelve and a half (25/2) es of oxygen to produce eight es of carbon dioxide, nine es of water, and five thousand and seventy-four kilojoules of heat. Of course, you may still write the balanced chemical equation with whole-number coefficients, but the heat of reaction must be doubled because 2 of octane are shown: 2 C 8 H 18(l) + 25 O 2(g) 16 CO 2(g) + 18 H 2 O (g) kj The heat term is part of the balanced equation, and if the coefficients are changed, then the heat of reaction is changed in the same way. In the previous example, if the coefficients are doubled, then the energy is doubled. This is necessary so that the equation agrees with the empirical ar heat of reaction, as shown below: = kj 2 C 8 H 18(l) = kj 1 C 8 H 18(l) To write a thermochemical equation, it is necessary to know the balanced chemical equation, the ar heat of reaction for one substance in the equation, and whether the reaction is endothermic (takes in energy) or exothermic (gives out energy). If a reaction is endothermic, remember to put the heat of reaction on the reactant side of the chemical equation. This heat is consumed, or used up, like the reactants and is listed along with the reactants. For example, the decomposition of Energy from Hydrocarbons 583

3 water is endothermic and has a ar heat of reaction of kj/ of water. The thermochemical equations may be written as H 2 O (l) kj H 2(g) O 2(g) or 2 H 2 O (l) kj 2H 2(g) + O 2(g) Whether a particular reaction is endothermic or exothermic is not easy to recognize, so you will be given this information as part of any question. The only exception is a combustion reaction, which you are expected to know is exothermic. Sample Problem 2 A student is experimenting with different substances to make a cold pack. From a calorimetry experiment she determines that the specific heat of solution of potassium bromate, KBrO 3(s), is an endothermic 0.25 kj/g. Calculate the ar heat of solution and write the thermochemical equation for the dissolving process. Solution h s = 0.25 kj/g KBrO 3 H s = 0.2 5kJ g g 1 KBrO 3 H s 4 2 kj KBrO 3 The ar heat of solution is an endothermic 42 kj/ and the thermochemical equation is written as KBrO 3(s) 42 kj KBrO 3(aq) Sample Problem 3 Magnesium is commonly used in flares and fireworks. The combustion of magnesium (Figure 1) releases 24.7 kj/g of magnesium. Calculate the ar heat of combustion of magnesium and write the thermochemical equation for this reaction. Solution h c = 24.7 kj/g Mg = 24. 7kJ g Mg g 1 = 6 00 kj Mg Figure 1 The energy released by the combustion or burning of magnesium is obvious from the bright light produced by this sparkler. A considerable quantity of heat is also produced in this exothermic reaction along with the white solid, magnesium oxide. The ar heat of combustion is an exothermic 600 kj/ and the thermochemical equation is written as or Mg (s) O 2(g) MgO (s) kj 2 Mg (s) + O 2(g) 2 MgO (s) kj 584 Chapter 12

4 12.3 Finally, you may be asked to use a given thermochemical equation to determine a ar heat of reaction for a particular substance. The next sample problem shows an example of this kind of question. Sample Problem 4 Hexane is a component of some naphtha fuels used in camping stoves. Use the following thermochemical equation to calculate the ar heat of combustion of hexane: 2 C 6 H 14(l) + 19 O 2(g) 12 CO 2(g) + 14 H 2 O (g) kj Solution = kj 2 C 6 H 14 = kj C 6 H 14 The ar heat of combustion of hexane is an exothermic 3543 kj/. Practice Understanding Concepts 1. How are heats of reaction determined? 2. What does the term specific mean as part of a scientific quantity? 3. What are the units for specific heat of reaction and ar heat of reaction? 4. What conversion factor is used to convert between specific and ar heats of reaction? 5. Are combustion reactions endothermic or exothermic? State some common examples. 6. The specific heat of combustion of acetylene is determined calorimetrically to be kj/g. (a) What is the ar heat of combustion of acetylene? (b) Use the ar heat of combustion of acetylene to write a thermochemical equation. 7. Coal gasification converts coal (assume pure carbon) and water into a combustible mixture of carbon monoxide and hydrogen. The product mixture is called coal gas and is a cleaner fuel than coal for electric power-generating stations. From calorimetry, the specific heat of reaction for carbon is an endothermic 10.9 kj/g. (a) What is the ar heat of reaction of carbon? (b) Write a thermochemical equation for the gasification reaction. 8. If carbon monoxide is not recovered and recycled in an industrial process, it can be burned to form carbon dioxide. This same conversion occurs in the catalytic converter of a car: 2 CO (g) + O 2(g) 2 CO 2(g) kj (a) Rewrite this chemical equation using one e as the coefficient for carbon monoxide. (b) What is the ar heat of combustion of carbon monoxide? Answers 6. (a) 1299 kj/ 7. (a) 131 kj/ 8. (b) 283 kj/ Energy from Hydrocarbons 585

5 Figure 2 In the human body, exothermic chemical reactions occur as fats and carbohydrates are metabolized. 9. Translate the ar heats of reaction given below into a balanced chemical equation, including the energy change as a term in the equation. (a) The ar heat of combustion for methanol is an exothermic kj/. (b) The ar heat of formation for liquid carbon disulfide is an endothermic 89.0 kj/. The carbon disulfide is formed from its elements in a synthesis reaction. (c) The ar heat of roasting (complete combustion) for zinc sulfide is an exothermic kj/. (d) The ar heat of simple decomposition for iron(iii) oxide to its elements is an endothermic kj/. 10. Pentane is a volatile component of gasoline. During the winter, more pentane is added to gasoline so that automobiles start more easily. The specific heat of combustion of pentane is determined calorimetrically to be kj/g. Use this information to write a thermochemical equation representing the complete combustion of pentane. Reflecting 11. Do endothermic and exothermic only apply to chemical reactions? What are some other changes that can also be endothermic or exothermic? A Theoretical Perspective on Energy Change Energy transfer is an important factor in all chemical changes. Exothermic reactions, such as the combustion of gasoline in a car engine or the metabolism of fats and carbohydrates in a human body (Figure 2), release energy into the surroundings. Endothermic reactions, such as photosynthesis (Figure 3) or the decomposition of water into hydrogen and oxygen, remove energy from the surroundings. Knowledge of energy and energy changes is important to society and to industry, and the study of energy changes provides chemists with important information about chemical bonds. How do we explain the energy changes measured in calorimeters in terms of ecules, atoms, and bonds? Just as glue holds objects together, electrical forces hold atoms together. In order to pull apart objects that are glued together, you have to supply some energy. Similarly, if atoms or ions are bonded together, energy is required to separate them. Separated atoms or ions release energy when they bond together again: bonded particles + energy separated particles separated particles bonded particles + energy Figure 3 Plants use energy from the Sun in a series of endothermic reactions called photosynthesis. bond energy: the energy required to break a chemical bond; the energy released when a bond is formed The stronger the bond holding the particles together, the greater the energy required to separate them. Bond energy is the energy required to break a chemical bond. It is also the energy released when a bond is formed. Even the simplest of chemical reactions may involve the breaking and forming of several individual bonds. The terms exothermic and endothermic are empirical descriptions of overall changes that can be explained by knowledge of bond changes. Consider the decomposition of water, for example: 2H 2 O (l) energy 2H 2(g) + O 2(g) The easiest way to supply the energy for this decomposition is electrically. This is known as the electrolysis of water. In this endothermic reaction, hydrogen oxygen bonds in the water ecules must be broken before the hydrogen hydrogen and oxygen oxygen bonds in the products can be formed. 586 Chapter 12

6 12.3 Breaking bonds requires energy and forming bonds releases energy. We know experimentally that the decomposition of water is endothermic, which means that it requires a constant input of energy to keep reacting. The theoretical interpretation is that it must take more energy to break the bonds in the reactants than is released when the new bonds in the products form (Figure 4). Energy Changes in the Decomposition of H 2 0 H H H H O O (atoms) Energy H O H (ecules) H O H H H H H O O (ecules) This height represents the net amount of energy absorbed in the reaction. Reaction Progress 2 H 2 O (l) + energy 2 H 2(g) + O 2(g) Figure 4 Since the overall change is endothermic, the energy required to break the O H bonds must be greater than the energy released when the H H and OO bonds form. For exothermic reactions, such as the formation of hydrogen chloride, the opposite is true (Figure 5). Exothermic reactions produce energy that transfers to the surroundings, usually as heat. In this case, more energy is believed to be released in forming new bonds than is absorbed in breaking bonds in the reactants. This is true for all exothermic reactions, including the combustion of fuels. Energy Changes in the Formation of HCl H H Cl (atoms) Cl Energy H H Cl Cl (ecules) H Cl H Cl } (ecules) Reaction Progress H 2(g) + Cl 2(g) 2 HCl (g) + energy This height represents the net amount of energy released in the reaction. Figure 5 Energy is absorbed in order to break the H H and Cl Cl bonds, but more energy is released when the H Cl bonds form. The overall result is an exothermic reaction. Energy from Hydrocarbons 587

7 SUMMARY Energy Changes specific heat of reaction, h r : ar heat of reaction, H r : heat of reaction, term in chemical equation thermochemical equation: reactants energy products reactants products energy Typical Units kj/g kj/ kj (endothermic) (exothermic) Practice Understanding Concepts 12. Define bond energy, and use the term in an example. 13. Hydrogen gas reacts with bromine vapour to form hydrogen bromide gas. (a) Write the balanced chemical equation for this reaction. (b) List the bonds that must be broken and the ones that must be formed during this reaction. (c) In a calorimetry experiment, the temperature of the surroundings increased as a result of this reaction. Is this reaction endothermic or exothermic? (d) Using your answer to (c), what interpretation can you make about the total energy required for the breaking of the reactant bonds versus the forming of bonds in the product? 14. Ordinary salt, sodium chloride, has an endothermic heat of solution. (a) If enough salt is dissolved in water, what happens to the temperature of the surroundings? (b) Write the dissociation equation for the dissolving of solid sodium chloride in water. (c) What type of chemical bonds must be broken in the dissolving process? (d) What bonds are formed to produce the aqueous ions? (e) Compare the energy required to break the bonds in the reactant to the energy released when the aqueous ions form. Reflecting 15. Describe how a fridge magnet might be used to illustrate the concepts of a chemical bond and bond energy. Questioning Hypothesizing Predicting Planning Conducting INQUIRY SKILLS Recording Analyzing Evaluating Communicating Investigation Combustion of Octane Research and development of gasoline are done by all oil companies that make this product. There are many characteristics of gasoline that are studied, but energy is one of the most important. The purpose of this investigation is to use the technological skills and energy concepts you have learned to determine the energy released from the combustion of octane, C 8 H 18(l), a component of gasoline. Complete the Analysis and Evaluation sections of the lab report. 588 Chapter 12

8 12.3 Question What is the balanced thermochemical equation for the complete combustion of octane? Experimental Design An alcohol-type burner containing octane is used and a measured quantity of water in a calorimeter is heated. Materials eye protection ring stand thermometer burner containing octane calorimeter (from Investigation , or metal can) balance matches bottle of distilled water Octane is highly flammable. Long hair should be tied back and loose sleeves should be contained by sleeve protectors or a lab coat. Be careful when handling hot glassware. Alcohol-type burners present a fire hazard. Follow the teacher s directions for use of burners. Procedure 1. Measure the mass of the burner and octane. 2. Measure the mass of a clean, dry calorimeter. 3. Fill the calorimeter about one-half full with distilled water and measure the total mass of the calorimeter and water. 4. Set up the calorimeter on the ring stand 2 3 cm above the wick of the burner. 5. Measure the initial temperature of the water in the calorimeter. 6. Light the burner and stir the water gently with the thermometer. 7. When the temperature of the water has increased about C, blow out the flame. 8. Continue to stir the water. Measure and record the highest temperature the water reaches after stirring. 9. Measure the final mass of the burner and remaining octane. 10. If time permits, empty the calorimeter and repeat steps 3 to 9 two more times. Analysis (a) Calculate the specific heat of combustion for octane for each trial and average the best results. (b) Convert the average specific heat into a ar heat of combustion. (c) Answer the Question. Evaluation (d) Evaluate the Evidence by considering the quality of the Materials and procedure. (e) Identify some sources of experimental error or uncertainty. (f) How confident are you about the answer you obtained? Discuss briefly, including any assumptions you made in the calorimeter calculations and the nature of the combustion. Energy from Hydrocarbons 589

9 Section 12.3 Questions Figure 6 Propane-fuelled vehicles are not allowed to park in underground parking lots. Propane is denser than air, and a dangerous quantity of propane could accumulate in the event of a leak. Figure 7 While the propane tank is being used, it gets sufficiently cold to freeze water vapour out of the air. Notice how the frost layer abruptly stops at the level of the liquid propane inside the tank. The same cooling process is used when hydrocarbons or other refrigerants are used in a refrigerator. Understanding Concepts 1. Hexane is a component of crude oil and gasoline. In a laboratory, the specific heat of combustion of hexane is determined calorimetrically to be kj/g. (a) Determine the ar heat of combustion of hexane. (b) Communicate the ar heat of combustion of hexane as a term in a thermochemical equation balanced for one e of hexane. (c) Communicate the above thermochemical equation balanced for two es of hexane. (d) Explain, using the concept of bond energy, why this reaction is exothermic. 2. For each of the following reactions, translate the given ar heat into a balanced chemical equation that includes the energy as a term in the equation. (a) Propane obtained from natural gas is used as a fuel in barbecues and vehicles (Figure 6). The ar heat of combustion for propane, as determined by calorimetry, is 2.04 MJ/. (b) Nitrogen monoxide forms from nitrogen and oxygen at the high temperatures inside an automobile engine. The ar heat of formation for nitrogen monoxide is an endothermic 90.2 kj/. (c) Some advocates of alternative fuels have suggested that cars could run on ethanol. The ar heat of combustion for ethanol is 1.28 MJ/. 3. For each of the following balanced thermochemical equations, calculate the ar heat of combustion for the substance that reacts with oxygen: (a) 2 H 2(g) + O 2(g) 2H 2 O (g) kj (b) 4 NH 3(g) + 7 O 2(g) 4NO 2(g) + 6 H 2 O (g) kj 4. The formation of hydrogen iodide from its elements is an endothermic reaction. (a) In this reaction, which bonds are broken and which bonds are formed? (b) What interpretation can be made about the energy required to break the reactant bonds and the energy released when the product bonds are formed? Applying Inquiry Skills 5. Many modern camping stoves use propane as the fuel. When the liquid propane vaporizes during use, the outside of the tank becomes very cold (Figure 7). (a) Is the vaporization of propane an endothermic or exothermic change? (b) Write a thermochemical equation for vaporization of propane showing energy on the appropriate side of the equation. (c) Design an experiment to determine the specific heat of vaporization of propane. Complete the Question, Experimental Design, Materials, and Procedure for this proposed investigation. 6. Scientists create new concepts by collecting as much evidence as possible and then looking for a pattern or general trend. The purpose of this exercise is to create a generalization about the energy released by different sizes of hydrocarbon ecules. Complete the Evidence and Analysis sections of the lab report. 590 Chapter 12

10 12.4 Question What is the relationship between the size of a hydrocarbon ecule and the ar heat of combustion? Experimental Design A Web site reference is used to obtain the ar heats of combustion,. These values are then compared to the number of carbon atoms in each ecule. Evidence (a) Using the Internet, find values for of various hydrocarbons. Analysis (b) According to the Evidence, is there a clear relationship between the size of the hydrocarbon ecule and the ar heat of combustion? Describe as specifically as you can what the relationship appears to be. Follow the links for Nelson Chemistry 11, GO TO Making Connections 7. The ideal fuel provides a minimum amount of carbon dioxide and a maximum quantity of energy. For a complete combustion, coal (assume carbon) has a ar heat of combustion of 395 kj/, and natural gas (assume methane) has a ar heat of combustion of 803 kj/. (a) Calculate the amount, in es, of carbon dioxide produced for every 1.0 MJ of energy in the complete combustion of coal and natural gas. (b) Based on the quantity of greenhouse gas, CO 2(g), released, which is the better fuel? 8. If you were designing an automobile and deciding what kind of hydrocarbon fuel to use, would you rather have information about the specific or the ar heats of combustion? Provide your reasoning Our Use of Fossil Fuels As you can see from Figure 1, Canadians are the world s second largest per capita consumers of energy. Approximately 45% of total energy production in Canada ends up as waste in the form of heat lost in the generation and transmission of electricity. The amount of energy lost is more than that available to many developing countries to support their populations and economies. Figure 2 (page 592) shows the energy consumption by Canadians from various sources between 1871 and As you can see, by 1900, fossil fuels had replaced wood as the main energy source. By 1950, fluid fossil fuels such as oil and natural gas had largely replaced coal. By 1985, fossil fuels accounted for about 87% of total energy use in Canada. Energy from hydroelectric sources (river dams) accounted for 11% and energy from nuclear reactors, 2%. Our dependence on fossil fuels for energy is likely to continue well into the 21st century. Mass of Oil Equivalent (t) Canada Sweden Australia United Kingdom United States Spain Argentina China Brazil Thailand Figure 1 Per capita energy consumption by country (1997) Energy from Hydrocarbons 591