ADDITIONAL RESOURCES Chemical changes occur around us, and inside us, all the time. When chemical reactions happen, one or more new substances are formed and energy is either given off or absorbed in the process. Spontaneous chemical reactions those that occur naturally under favourable conditions release free energy, meaning they are exergonic. Cellular respiration is an example. Endergonic reactions, by contrast, require net energy input to make the reaction happen. Photosynthesis is an example. This learning resource explores energy associated with chemical reactions. Areas covered include activation energy and product formation, standard heats of reaction, thermochemical equations and Hess Law, and specific heat capacity and calorimetry. It is essential viewing for any senior secondary level Chemistry student, providing clear explanations and easily understood visuals, including comprehensive chemical equations, about concepts that many students find challenging. Duration of resource: 21 Minutes Year of Production: 2013 Stock code: VEA12052 Resource written by: Rebecca Ross BSc.Hon.Grad.B.ed
For Teachers Introduction This program introduces students to the complex concepts involved in energy changes that occur in all chemical reactions. These include the differences between endergonic, exergonic, endothermic and exothermic reactions, the activation energy required by reacting particles in order for a reaction to proceed, in addition to a detailed evaluation of how bond dissociation energies within molecules can provide scientists with a means of calculating the overall energy changes that occur between reactants and products. Students are guided through Hess s Law, used in multi-step reactions to calculate the overall energy change, and, then shown how enthalpy is used in calorimetry with specific heat capacity and temperature changes to determine the fundamental chemistry of what is occurring within this closed system. Timeline 00:00:00 Chemical reactions and energy 00:04:24 Activation energy and production formation 00:08:41 Standard heats of reaction 00:13:09 Thermochemical equations and Hess's Law 00:16:20 Specific heat capacity and calorimetry 00:19:58 Credits 00:20:36 End of program Related Titles Chemical Analysis Techniques Water: A Unique Chemical Recommended Resources http://www.vcaa.vic.edu.au/pages/vce/exams/examsassessreports.aspx http://www.scientificamerican.com/article.cfm?id=what-is-an-exothermic-rea http://www.lahc.edu/classes/chemistry/arias/exp%2010%20-%20thermof11.pdf http://www.wisc-online.com/objects/viewobject.aspx?id=gch8705 2
Student Worksheet Initiate Prior Learning 1. In the table below are the results of four experiments that were carried out by a Year 11 Chemistry student. Use your prior knowledge of physical and chemical changes to identify the clue(s) that suggest whether a chemical or physical change has taken place in each case. Justify your answer. REACTION NUMBER REACTANTS DETAILED OBSERVATIONS 1 Ice cube on watch glass Ice cube melted (temperature of surroundings 27 o C) 2 Magnesium Metal + HCl Solution bubbles, test tube heats up and eventually the magnesium strip dissolved into the solution. 3 Silver Nitrate + HCl A white precipitate comes out of solution. 4 Sodium Hydroxide (pellets) + HCl Initial Temp: 25 o C Final Temp: 29 o C Temperature increase Reaction 1 Reaction 2 3
Reaction 3 Reaction 4 2. For a reaction to proceed, a certain amount of energy is needed for the whole process to begin. The energy needed to get the reaction started is called the activation energy (Ea). We know (from question 1 above) that when we place a piece of magnesium metal in a hydrochloric acid solution, we produce hydrogen gas (which we can prove using the pop test), the test tube heats up (quite vigorously, depending on the concentration of hydrochloric acid) and eventually, the magnesium disappears. But can we make this reaction more interesting and actually speed it up or deliberately slow it down? And if so, how can this be done? Identify two ways that you could increase the dissolution of a piece of magnesium metal into a solution of hydrochloric acid. Justify your response. 4
3. Balancing Chemical equations is a fundamental aspect of any senior Chemistry syllabus. Write the products of the following double replacement reactions. Remember to check the equations are balanced. a) NaOH + Al(CO 3 ) 3 b) (NH 4 ) 3 PO 4 + CaSO 4 Now write the products of the following single replacement reactions: c) K (s) + PbI (aq) d) (d) Mg (s) + ZnCl (aq) 5
Active Viewing Guide Chemical Reactions and Energy 1. How do you know when a chemical reaction has occurred? 2. Given that the generic equation for the oxidation of an alkane is: Alkane + O 2 CO 2 + H 2 O Write the balanced equation for the oxidation of: a) Methane b) Octane 3. Contrast an exergonic reaction to an endergonic reaction. Give one example of each. 4. Circle the correct response. An exothermic reaction: a) Releases heat energy to the surroundings; b) Has the energy (joules) written into the products side of the equation; c) Includes the combustion of propane; d) All of the above. 6
5. Endothermic processes require the input of energy in order for a reaction to proceed. In doing so, endothermic processes cool down the surroundings as heat is being removed. Explain briefly why ice melting is an endothermic process, using the following scenario: An ice cube is removed from the freezer and placed on a bench at 25 o C. 6. A closed system is required when conducting Calorimetry experiments to determine enthalpy changes. Why do you think this would be a necessity? Activation energy and product formation 7. List three ways that the kinetic energy of reacting particles can be increased in order for a reaction to be more likely to occur. 7
8. The rate determining step can be represented using an energy diagram as shown below. A + B ABC (2 steps) A + B AB (slow) Slow AB + C ABC (fast) Fast A + B ABC AB + C Which is the rate determining step in the following diagram. Explain. Activated complex B Activated complex A Activated complex C Reactants Intermediate A Intermediate B Products 8
Standard heats of reaction 9. What is bond dissociation energy? 10. What is the enthalpy change when each of the following reactions takes place? Identify each reaction as exothermic or endothermic. a) CH 4 (g) + 2 O 2 (g) CO 2 (g) + 2 H 2 O (g) b) CH 4 (g) + Cl 2 (g) CH 3 Cl (g) + HCl (g) The relevant bond energies are: C H 414 kj/mol O==O 502 kj/mol C==O 730 kj/mol O H 464 kj/mol Cl Cl 243 kj/mol H Cl 431 kj/mol C Cl 331 kj/mol 9
Thermochemical equations and Hess s Law 11. Explain Hess Law, also known as the law of constant heat summation, in your own words. 12. Find the ΔH for the overall reaction below, given the following reactions and subsequent ΔH values: PCl 5 (g) PCl 3 (g) + Cl 2 (g) P 4 (s) + 6Cl 2 (g) 4PCl 3 (g) ΔH = -2439 kj 4PCl 5 (g) P 4 (s) + 10Cl 2(g) ΔH = 3438 kj 10
Specific heat capacity and calorimetry 13. Calculate the amount of heat needed to increase the temperature of 250g of water from 20 o C to 46 o C. Specific heat Capacity of water = 4.184 J/g o C. 14. Calculate the mass of ice involved if a block of ice is heated from 0 C to 100 C using 440 kj of energy. 15. How much energy is needed to increase the temperature of 500 g of lead from 20ºC to 45ºC? The specific heat capacity of lead is 128 J/kg ºC. 11
Extension Activities 1. Prior to the development of the kinetic molecular theory, Arrhenius had proposed that the rate constant was related to the absolute temperature by the equation: k = A(10 -Ea/2.30 RT ) Where A is a constant characteristic of the particular reaction, E a is the activation energy, and R is the gas constant with a numerical value equal to 8.314 J/K/mole. Assuming that E a is a constant, when log k is plotted against T -1, the slope of the line is equal to - E a /(2.30 x R). An experiment was carried out to determine the activation energy of the following reaction: A + B C and the following data was collected: Trial Rate Constant k Temperature (K) 1 0.001128 100 2 0.010840 150 3 0.033575 200 4 0.066179 250 Calculate the activation energy for this reaction. 2. When 4.5 grams of ethanol, C 2 H 5 OH, was burnt in excess oxygen, 105 kj of heat was released. Write the equation for the reaction and calculate the molar enthalpy change H, for the reaction. 3. Consider the table below, which shows heat capacities of common materials. On a hot day a swimmer leaves the water, walks across the sand and, when entering a car sits on the metal buckle of the seat belt. Explain how the data in the table is related to the swimmer s experiences. Substance Specific Heat Capacity (J.g -1o C -1 ) Water 4.18 Cooking oil -2.2 Ethanol 2.46 Iron 0.45 Copper 0.39 Aluminum 0.90 sand 0.48 12
4. Calculate the heat of formation of methane given the following: C (s) + O 2(g) CO 2(g) H = -393 kj 2H 2(g) + O 2(g) 2H 2 O (l) H = -571 kj CH 4(g) + 2O 2(g) CO 2(g) + H 2 O (l) H = -882 kj 5. Design an experiment to measure the enthalpy change for the reaction of solid sodium hydroxide and dilute hydrochloric acid carried out in: a) one step: NaOH (s) + HCl (aq) NaCl (aq) + H 2 O (l) b) two steps: (i) NaOH (s) + H 2 O (l) NaOH (aq) (ii) NaOH (aq) + HCl (aq) NaCl (aq) + H 2 O (l) Be sure to address all the safety issues involved in this experiment. 13
Suggested Student Responses Initiate Prior Learning 1. In the table below are the results of four experiments that were carried out by a Year 11 Chemistry student. Use your prior knowledge of physical and chemical changes to identify the clue(s) that suggest whether a chemical or physical change has taken place in each case. Justify your answer. REACTION NUMBER REACTANTS DETAILED OBSERVATIONS 1 Ice cube on watch glass Ice cube melted (temperature of surroundings 27 o C) 2 Magnesium Metal + HCl Solution bubbles, test tube heats up and eventually the magnesium strip dissolved into the solution. 3 Silver Nitrate + HCl A white precipitate comes out of solution. 4 Sodium Hydroxide (pellets) + HCl Initial Temp: 25 o C Final Temp: 29 o C Temperature increase Reaction 1 (Melting ice cube) is a physical change. The reactant is solid H 2 O and the product is liquid H 2 O. Hence, only a change of phase has occurred, not a chemical reaction. Reaction 2 Mg + HCl) Mg + HCl H 2 + MgCl 2 Three results suggest here that this was a chemical change. Firstly, the reaction vessel bubbled, indicating a gas was being produced, then the temperature of the solution heated up and finally, the magnesium dissolved into solution. All these three things are indicators of a chemical change, as the products differed from the reactants. Reaction 3 (AgNO3 + HCl) AgNO 3 + HCl AgCl + HNO 3 The white precipitate that comes out of solution is the indicator that the products are different from the reactants and this is a chemical reaction. The white precipitate is AgCl as nitric acid (HNO 3 ) is an aqueous substance. 14
Reaction 4 (NaOH + HCl) When sodium hydroxide pellets are added to dilute hydrochloric acid, the reaction is highly exothermic, and the temperature of the solution begins to increase. This increase in temperature is an indicator that a chemical reaction has occurred, with the products being H 2 O and NaCl. 2. For a reaction to proceed, a certain amount of energy is needed for the whole process to begin. The energy needed to get the reaction started is called the activation energy (Ea). We know (from question 1 above) that when we place a piece of magnesium metal in a hydrochloric acid solution, we produce hydrogen gas (which we can prove using the pop test), the test tube heats up (quite vigorously too depending on the concentration of hydrochloric acid) and eventually, the magnesium disappears. But can we make this reaction more interesting and actually speed it up or deliberately slow it down? And if so, how can this be done? Identify two ways that you could increase the dissolution of a piece of magnesium metal into a solution of hydrochloric acid. Justify your response. There are only three ways in total that can be chosen in order to increase the rate of dissolution of magnesium into hydrochloric acid. (Students only needed to choose 2) Surface Area as an increase in surface area increases the rate of a reaction, then the piece of magnesium metal could be cut up in to many, many small pieces. This effectively allows for more 3-dimensional surfaces to be exposed and the likelihood of collisions would increase which would then increase the likelihood of increasing the rate of product formation. Concentration- An increase in concentration of the hydrochloric acid solution would also increase the rate of a reaction. This is because as more hydrochloric acid formula units are added to the solution, the density of them increases, which increases the number of likely collisions with the magnesium strip, hence increasing the likelihood collisions, will occur. Temperature An increase in the temperature of the hydrochloric acid solution will increase the kinetic energy of the hydrochloric acid formula units, hence resulting in increased movement and an increased likelihood of collisions with the magnesium strip. 3. Balancing Chemical equations is a fundamental aspect of any senior Chemistry syllabus. Write the products of the following double replacement reactions. Remember to check the equations are balanced. a) NaOH + Al(CO 3 ) 3 3Na 2 CO 3 + 2Al(OH) 3 b) (NH 4 ) 3 PO 4 + CaSO 4 3(NH 4 ) 2 SO 4 + Ca 3 (PO 4 ) 2 Now write the products of the following single replacement reactions: c) K (s) + PbI (aq) KI (aq) + Pb (s) d) (d) Mg (s) + ZnCl (aq) MgCl 2(aq) + 2Zn (s) 15
Active Viewing Guide Chemical Reactions and Energy 1. How do you know when a chemical reaction has occurred? One or more substances are formed in the products. 2. Given that the generic equation for the oxidation of an alkane is: Alkane + O 2 CO 2 + H 2 O Write the balanced equation for the oxidation of: a) Methane CH 4 + 2 O 2 CO 2 + 2H 2 O b) Octane 2C 8 H 18 + 25O 2 16CO 2 + 18H 2 O 3. Contrast an exergonic reaction to an endergonic reaction. Give one example of each. Exergonic reactions release free energy whereas endergonic reactions require an energy input. An example of an exergonic reaction is cellular respiration, which produces the ATP (adenosine triphosphate) molecule. An example of an endergonic reaction is photosynthesis which requires the presence of light energy for the reactions to proceed. 4. Circle the correct response. An exothermic reaction: a) Releases heat energy to the surroundings; b) Has the energy (joules) written into the products side of the equation; c) Includes the combustion of propane; d) All of the above. 5. Endothermic processes require the input of energy in order for a reaction to proceed. In doing so, endothermic processes cool down the surroundings as heat is being removed. Explain briefly why ice melting is an endothermic process, using the following scenario: An ice cube is removed from the freezer and placed on a bench at 25 o C. When an ice cube is removed from the freezer and the surrounding temperature is increased to 25 o C, energy is removed from the surroundings and transferred to the ice cube. In doing so, the ice cube heats up (hence melts) and in response to the transfer of energy, the temperature of the surroundings (particularly very close to the ice cube) feels very cool. 6. A closed system is required when conducting Calorimetry experiments to determine enthalpy changes. Why do you think this would be a necessity? Closed systems are required to ensure that the results obtained are accurate. Closed systems require insulation from the outside environment and these include lids. When a system is open heat can escape from the internal environment resulting in incorrect results. 16
Activation energy and product formation 7. List three ways that the kinetic energy of reacting particles can be increased in order for a reaction to be more likely to occur. Increase the pressure (of gaseous substances), increase the concentration, stirring, heating the reactants. 8. The rate determining step can be represented using an energy diagram as shown below. A + B ABC (2 steps) A + B AB (slow) Slow AB + C ABC (fast) Fast A + B ABC AB + C Which is the rate determining step in the following diagram? Explain. Activated complex B Activated complex A Activated complex C Reactants Intermediate A Intermediate B Products The rate determining step is the formation of activated complex B as it requires the most energy, hence the slower of the three activated complexes to be formed. 17
Standard heats of reaction 9. What is bond dissociation energy? Bond dissociation energy is the energy required to break a single bond in a molecule. 10. What is the enthalpy change when each of the following reactions takes place? Identify each reaction as exothermic or endothermic. a) CH 4 (g) + 2 O 2 (g) CO 2 (g) + 2 H 2 O (g) CH 4 + 2O 2 CO 2 + 2H 2 O Reactants - Products = H [(4 x 414) + (2 x 502)] - [(2 x 730) + (2 x (2 x 464))] =? (1656 + 1004) - (1460 + 1856) 2660-3316 H = - 656 kj/mol Negative sign indicates energy released hence exothermic. b) CH 4 (g) + Cl 2 (g) CH 3 Cl (g) + HCl (g) CH 4 + Cl 2 CH 3 Cl + HCl Reactants - Products = H [(4 x 414) + (243)] - [(3 x 414) + 331 + 431] =? (1656 + 243) - (1242 + 331 + 431) =? 1899-2004 H = - 105 kj/mol Negative sign indicates energy released hence exothermic. The relevant bond energies are: C H 414 kj/mol O==O 502 kj/mol C==O 730 kj/mol O H 464 kj/mol Cl Cl 243 kj/mol H Cl 431 kj/mol C Cl 331 kj/mol Thermochemical equations and Hess s Law 11. Explain Hess Law, also known as the law of constant heat summation, in your own words. Hess Law explains that the enthalpy in a chemical reaction is the same whether the reaction occurs in one step, or in two or more steps. So if a reaction occurs in several steps, the sum of standard enthalpies of those intermediate steps equals the standard enthalpy for the single, overall reaction. 18
12. Find the ΔH for the overall reaction below, given the following reactions and subsequent ΔH values: PCl 5 (g) PCl 3 (g) + Cl 2 (g) P 4 (s) + 6Cl 2 (g) 4PCl 3 (g) ΔH = -2439 kj 4PCl 5 (g) P 4 (s) + 10Cl 2(g) ΔH = 3438 kj PCl 5 (g) PCl 3 (g) + Cl 2 (g) Equation 1 P 4 (s) + 6Cl 2 (g) 4PCl 3 (g) ΔH = -2439 kj Equation 2 4PCl 5 (g) P 4 (s) + 10Cl 2(g) ΔH = 3438 kj Add equation 1 to equation 2 4 P 4 (s) + 6Cl 2 (g) + 4PCl 5 (g) 4PCl 3 (g) + P 4 (s) + 10Cl 2(g) Equation 3 4PCl 5 (g) 4PCl 3 (g) + 4 Cl 2(g) ΔH = -2439 kj + 3438 kj ΔH = 999kJ Divide equation 3 by 4 PCl 5 (g) PCl 3 (g) + Cl 2(g) ΔH = 249.8kJ Specific heat capacity and calorimetry 13. Calculate the amount of heat needed to increase the temperature of 250g of water from 20 o C to 46 o C. Specific heat Capacity of water = 4.184 J/g o C. q = m x C x T q = 250g x 4.18J/g o C x 26 o C q = 37,620J or 38kJ 14. Calculate the mass of ice involved if a block of ice is heated from 0 C to 100 C using 440 kj of energy. q = m x C x T 440 = m x 4184 x (100 0) 440 = m x 418400 m = 440/418400 = 1.05 x 10-3 kg m ~ 1gram 15. How much energy is needed to increase the temperature of 500 g of lead from 20ºC to 45ºC? The specific heat capacity of lead is 128 J/kg ºC. mass of lead = 500 1000 = 0.5 kg temperature change = 45 20 = 25ºC q = mc T q = 0.5 128 25 q = 1600 J (1.6 kj) 19
Extension Activities 1. Prior to the development of the kinetic molecular theory, Arrhenius had proposed that the rate constant was related to the absolute temperature by the equation: k = A(10 -Ea/2.30 RT ) Where A is a constant characteristic of the particular reaction, E a is the activation energy, and R is the gas constant with a numerical value equal to 8.314 J/K/mole. Assuming that E a is a constant, when log k is plotted against T -1, the slope of the line is equal to - E a /(2.30 x R). An experiment was carried out to determine the activation energy of the following reaction: A + B C and the following data was collected: Trial Rate Constant k Temperature (K) 1 0.001128 100 2 0.010840 150 3 0.033575 200 4 0.066179 250 Calculate the activation energy for this reaction. Trial Rate Constant k Log k Temperature T -1 (K) 1 0.001128-2.94769 100 0.01 2 0.010840-1.96497 150 6.66x10-3 3 0.033575-1.47398 200 5 x 10-3 4 0.066179-1.17928 250 4 x 10-3 Log k 0-0.5-1 -1.5-2 -2.5-3 -3.5 Log k versus 1/T 0 0.005 0.01 0.015 y = -294.7x - 0.000 R² = 1 1/T Slope of the line = mx + C y = -294.76x - 0.0006 m = slope = -294.76 therefore, -294.76 = -Ea/ (2.30 x R) where R = 8.314 So, Ea = 294.76 x (44.0642) Ea = 12988.36 Joules Ea = 12.9884 kj 20
2. When 4.5 grams of ethanol, C 2 H 5 OH, was burnt in excess oxygen, 105 kj of heat was released. Write the equation for the reaction and calculate the molar enthalpy change H, for the reaction. C 2 H 5 OH (l) + 3O 2(g) 2CO 2(g) + 3H 2 O (l) Need to calculate the heat released per mole of ethanol used Molar mass of ethanol = 2 x 12.01 + 6 x 1.01 + 16.0 = 46.1 g/mol Number of moles of ethanol burnt = 4.5 /46.1 = 0.098 moles Heat released by 0.098 moles = 105 kj Heat released per mole = 105/0.098 = 1071.43 kj/mol When a reaction releases heat (as is done here) we say that the heat absorbed is minus the quantity of heat released (in this case, - 1071.43 kj/mol). Because H is defined as heat absorbed, H = 1071.43 kj/mol. 3. Consider the table below, which shows heat capacities of common materials. On a hot day a swimmer leaves the water, walks across the sand and, when entering a car sits on the metal buckle of the seat belt. Explain how the data in the table is related to the swimmer s experiences. Substance Specific Heat Capacity (J.g -1o C -1 ) Water 4.18 Cooking oil -2.2 Ethanol 2.46 Iron 0.45 Copper 0.39 Aluminum 0.90 sand 0.48 Metal and sand have lower specific heats than water. This means that metal and sand reach higher temperatures than water when they absorb energy from the sun. The high thermal conductivity of metal allows the rapid conduction of heat from the hot metal buckle of the seat belt to the skin of the swimmer. 21
4. Calculate the heat of formation of methane given the following: C (s) + O 2(g) CO 2(g) H = -393 kj 2H 2(g) + O 2(g) 2H 2 O (l) H = -571 kj CH 4(g) + 2O 2(g) CO 2(g) + H 2 O (l) H = -882 kj Equation 1 C (s) + O 2(g) CO 2(g) H = -393 kj Equation 2 2H 2(g) + O 2(g) 2H 2 O (l) H = -571 kj Equation 3 CH 4(g) + 2O 2(g) CO 2(g) + H 2 O (l) H = -882 kj Reverse Equation 3 Equation 4 CO 2(g) + 2H 2 O (l) CH 4(g) + 2O 2(g) H = 882 kj Add Equations 1 & 4 1 C (s) + O 2(g) + CO 2(g) + 2H 2 O (l) CO 2(g) + CH 4(g) + 2O 2(g) H = (-393 + 882) kj Equation 5 C (s) + 2H 2 O (l) CH 4(g) + O 2(g) H =489 kj Add equations 5 & 2 2H 2(g) + O 2(g) + C (s) + 2H 2 O (l) 2H 2 O (l) + CH 4(g) + O 2(g) H =(489 571)kJ Equation 6 2H 2(g) + C (s) CH 4(g) H = 82kJ 5. Design an experiment to measure the enthalpy change for the reaction of solid sodium hydroxide and dilute hydrochloric acid carried out in: c) one step: NaOH (s) + HCl (aq) NaCl (aq) + H 2 O (l) d) two steps: (i) NaOH (s) + H 2 O (l) NaOH (aq) (ii) NaOH (aq) + HCl (aq) NaCl (aq) + H 2 O (l) Answers will vary, but one could be: Measure out 200 ml of 0.3M HCl into a measuring cylinder and record the temperature. Weigh out approximately 2 grams of sodium hydroxide pellets into a previously weighted Styrofoam cup; Place the Styrofoam cup inside another and then add the 300 ml of HCL. Stir well and record the highest temperature reached. Using a clean dry calorimeter, repeat the above three points using distilled water instead of acid. Pour 100 ml of 0.5 M NaOH into a clean dry calorimeter and record the temperature. Pour 100 ml of 0.6 M HCl into a measuring cylinder and record the temperature. Pour the 0.6 M HCl into the calorimeter containing the 0.5M NaOH solution. Stir continuously and record the highest temperature reached. Safety issues include corrosiveness of NaOH and HCl, PPE needs to be worn, glassware being used, hence care in carrying items in the lab etc. 22