Reaction Energy - Student Exploration

Name: Date: Logan Langenkamp Student Exploration: Reaction Energy Directions: Follow the instructions to go through the simulation. Respond to the questions and prompts in the orange boxes. Vocabulary: calorimeter, chemical bond, endothermic, enthalpy, exothermic, Hess’s law Prior Knowledge Questions (Do these BEFORE using the Gizmo.) 1.​ Two magnets are stuck together. What might you have to do to get them to separate? Pull them apart 2.​ Suppose you held two magnets a short distance apart, then let go. What would happen? They would get attached to eachother 3.​ Think about the magnets in terms of energy. In which case do you increase the potential energy of the magnets? In which case do you increase the kinetic energy of the magnets? Potential energy would increase when the magnets are being held apart, but they would increase energy when they were connected or moving towards each other Gizmo Warm-up Just like magnets, atoms of different elements are attracted together to form chemical bonds. Breaking these bonds requires energy. When a new bond forms, energy is released and temperatures rise. In the Reaction Energy Gizmo, you will explore how the energy of chemical bonding relates to temperature changes that occur during chemical reactions. To begin, check that Reaction 1 and Forward are selected. In this reaction, hydrogen (H2) and oxygen (O2) react to form water (H2O). The reaction takes place inside a device called a calorimeter. Inside the calorimeter, a small chamber holds the reactants. The rest of the calorimeter is filled with water. 1.​ Click Play ( ). What happens? ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ Hydrogen atoms pair up with oxygen atoms to form H2O ​ 2.​ How does the temperature change? ​ ​ ​ ​ The temperature increases from 21c to 27.4c Reproduction for educational use only. Public sharing or posting prohibited. © 2020 ExploreLearning™ All rights reserved​ Activity A: Energy of chemical bonds Get the Gizmo ready: ●​ Check that Reaction 1 and Forward are selected. ●​ Select the INVESTIGATION tab. Introduction: The heat energy stored in a chemical system is called the enthalpy (H) of the system. When atoms are joined by a chemical bond, energy must be added to pull them apart. This increases the enthalpy of the system. When a chemical bond forms, energy is released as shared electrons move into lower-energy orbitals. This causes the enthalpy to decrease. Question: How can you predict how much energy is released in a chemical reaction? 1.​ Predict: In the warm-up activity, you observed how the reaction inside the chamber affected the temperature of the surrounding water. Based on what happens to the surrounding water, do you think heat energy (enthalpy) is absorbed in the reaction or released? Explain. I think the energy was released since the hydrogen atoms created bonds with the oxygen, this released energy and caused the temperature to increase. ​ ​ ​ ​ ​ 2.​ Observe: In the Gizmo, the energy required to break a chemical bond is modeled by placing a molecule into a set of mechanical claws. Place one of the hydrogen (H2) molecules between the claws, and press Break bond. ​ A.​ What happens? ​ The energy absorbed changed to 436 KJ/mol ​ ​ ​ ​ ​ ​ ​ B.​ Look under the Energy absorbed column of the table. How much energy was required to break this bond?​ ​ 436 KJ Note: The energy is given here in units of kilojoules per mole (kJ/mol). This is the energy, in kilojoules, required to break all of the H–H bonds in one mole of H2 gas. C.​ Remove the hydrogen atoms from the claws and then break apart the other H–H molecule.​ What is the total energy absorbed so far? 872 KJ/mol 3.​ Measure: Notice that the oxygen atoms are connected by a double covalent bond. This is because the oxygen atoms share two pairs of electrons. Place the oxygen molecule in the claws and press Break bond. A.​ How much energy is required to break the first O–O bond? 349 KJ/mol B.​ Press Break bond. How much energy is needed to break both bonds? 495 KJ/mol ​Reproduction for educational use only. Public sharing or posting prohibited. © 2020 ExploreLearning™ All rights reserved​ C.​ What is the total energy required to break up two moles of H2 molecules and one mole of O2 molecules? ​ 1367 KJ/mol 4.​ Create: Remove the two oxygen atoms from the claws. Now the claws disappear and you see a template for creating a water molecule. Drag an oxygen and a hydrogen atom into the template. (If necessary, use the Key on the right-hand side as a reference.) A.​ Click Create bond. ​ What happens? ​ Energy is released when the bonds are created B.​ The “jiggling” animation you see represents the release of kinetic energy that occurs when a bond is formed. ​ How much energy was released? 436 KJ/mol C.​ Drag another hydrogen molecule into the template and click Create bond to make a water molecule. What is the total energy released so far? 926 KJ/mol D.​ Drag the first water molecule away from the template, then use the Gizmo to create a second water molecule. What is the total energy released now? 1852 KJ/mol 5.​ Calculate: Compare the energy absorbed in breaking up the molecules to the energy released when new bonds are formed. A.​ In this reaction, was more energy absorbed or released? More energy was released B.​ How does this relate to the change in temperature observed for this reaction? This relates to the changes in temperature because when energy is released, molecules speed up rising the temperature C.​ The change in enthalpy (∆H) of the system is equal to the total energy absorbed minus the total energy released. What is the ∆H value for this reaction? -485 KJ/mol Compare this value to the Theoretical ∆H listed on the right side. 6.​ Draw conclusions: The experimental ∆H value was determined by measuring how much heat the reaction produced inside the calorimeter. This is calculated based on the temperature change of the reaction, the ​Reproduction for educational use only. Public sharing or posting prohibited. © 2020 ExploreLearning™ All rights reserved​ amount of water inside the calorimeter, and the specific heat of the calorimeter. Compare the theoretical change in enthalpy to the experimental value. Are these values close? ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ The value of the theoretical is -485KJ/mol, while the experimental is 484KJ/mol, meaning they are almost opposite of each other ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ Get the Gizmo ready: Activity B: Reaction direction ●​ Select the Reaction tab. ●​ With Reaction 1 selected, choose Reverse. Introduction: Reactions that emit heat energy are called exothermic reactions. In an exothermic reaction, more heat energy is released by the formation of bonds than is absorbed in the breaking of bonds. In an exothermic reaction, the temperature of the surroundings will increase and the enthalpy of the system will decrease as energy is emitted from the system. In an endothermic reaction, more heat energy is absorbed in the breaking of bonds than is released in the formation of bonds. The temperature decreases and the enthalpy increases. Question: How can you predict the direction of a chemical reaction? 1.​ Observe: Notice that in the reverse reaction, the container is filled with water molecules. A.​ Click Play. What happens? The temperature and speed of the molecules don’t change B.​ Did the temperature change? No C.​ Is there any evidence that a reaction took place? No ​ Explain. There is no evidence because the reaction chamber still contains 0.2 mols of water 2.​ Investigate: Switch to the INVESTIGATION tab. Use the claws to break up the two water molecules, and then form two hydrogen molecules and one oxygen molecule. A.​ How much energy was absorbed in breaking up the water molecules? 1852 KJ/mol B.​ How much energy was released in forming the product molecules? ​ 1367 KJ/mol C.​ What is the total enthalpy change in the system? ​ 485 KJ/mol D.​ Based on this result, is this reaction exothermic or endothermic? ​ Endothermic ​ Explain. More energy was absorbed than released ​Reproduction for educational use only. Public sharing or posting prohibited. © 2020 ExploreLearning™ All rights reserved​ Most endothermic reactions do not occur unless there is a continuous input of energy. For example, water molecules can decompose into hydrogen and oxygen gas by adding a salt and passing an electrical current through the water. 3.​ Investigate: Select Reaction 2. On the INVESTIGATION tab, go through the reaction in both the forward and reverse directions. State the energy absorbed, energy released, and theoretical ∆H value in each direction: Direction Energy absorbed Energy released Theoretical ∆H Forward 2347 KJ/mol 2144 KJ/mol 203 KJ/mol Reverse 2144 KJ/mol 2347 KJ/mol -203 KJ/mol What do you notice about the theoretical ∆H values in the forward and reverse directions? The values are opposites of each other ​ ​ ​ 4.​ Predict: Based on your findings, do you think Reaction 2 is more likely to occur in the forward or reverse ​ direction? Reverse Explain.​ Most reactions are exothermic ​ 5.​ Test: Select the REACTION tab and test your prediction by running the reaction in the forward and reverse directions. ​Describe your findings. ​ ​ ​ ​ ​ The reaction occurred in reverse ​ When the reaction occurred, how did the theoretical ∆H value compare to the experimental value?​ They were different by 8. Theoretical value was -203 and experimental was -211 ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ 6.​ Investigate: On the INVESTIGATION tab, go through the remaining reactions, in either direction. (You will only need to look at each reaction once.) Fill in the table, then predict if the reaction will proceed in the forward or reverse directions. Reaction Direction Theoretical ∆H Forward or reverse? 3 Reverse 4 Forward ​Reproduction for educational use only. Public sharing or posting prohibited. © 2020 ExploreLearning™ All rights reserved​ 7.​ Test: Test your predictions on the REACTION tab. ​ What did you find? ​ Reaction 3 was reverse and reaction 4 was forward ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​Reproduction for educational use only. Public sharing or posting prohibited. © 2020 ExploreLearning™ All rights reserved​ Activity C: Bond enthalpy Get the Gizmo ready: ●​ Select Reaction 3 and Reverse. ●​ Select the INVESTIGATION tab. Introduction: Each chemical bond has a “bond enthalpy” that describes how much energy is absorbed to break a bond and how much energy is released when the bond is formed. (These values are the same.) A chart of bond enthalpies for some common bonds is shown below. Bond C–H C–C C=C C=O Enthalpy (kJ/mol) 413 348 614 799 Bond O–H O=O O–S O=S Enthalpy (kJ/mol) 463 495 265 523 Bond H–H N–H N≡N S–S Enthalpy (kJ/mol) 436 391 941 266 Question: How can you use bond enthalpy to predict the total enthalpy change of a chemical reaction? 1.​ Calculate: Consider the reaction CH4 + 2O2 → CO2 + 2H2O. A.​ In the reactants, how many C–H bonds are there? ​ 1 B.​ Using the chart above, what is the total bond enthalpy of these bonds? ​ 2174 C.​ How many O=O bonds are there? ​ 495 D.​ What is the total enthalpy of these bonds? ​ 1746 E.​ What is the total bond enthalpy of the reactants? ​ 2348 F.​ Do the same calculation for the products of the reaction, CO2 + 2H2O. Carefully count how many of each bond there is, and consider whether bonds are single or double bonds. Show your work and list the total bond enthalpy of the products below. C=O equals 1598 O-H= 463 x 2 =926 add both Total bond enthalpy of products: 2524 G.​ Based on the enthalpy of the reactants and products, what is the ∆H value for this reaction? (Recall that ∆H = energy absorbed – energy released.) ​ ​ 4994 KJ/mol H.​ Use the Gizmo to check your results, and correct any errors if necessary. ​Reproduction for educational use only. Public sharing or posting prohibited. © 2020 ExploreLearning™ All rights reserved​ Enthalpy (kJ/mol) 413 348 614 799 Bond C–H C–C C=C C=O Enthalpy (kJ/mol) 463 495 265 523 Bond O–H O=O O–S O=S Enthalpy (kJ/mol) 436 391 941 266 Bond H–H N–H N≡N S–S 2.​ On your own: Using the chart above, predict the enthalpy change for the following reactions. Show your work. A key of all relevant molecule structures is shown on the right. A.​ N2 + 3H2 → 2NH3​ N2: N≡N H2: H–H NH3: 2249 KJ/mol Total ∆H: B.​ 8SO2 → S8 + 8O2 Exothermic or endothermic: ​ SO2: O=S–O ​ S8: 8 S–S bonds O2: O=O 14680 KJ/mol Total ∆H: Exothermic Exothermic or endothermic: C.​ C2H4 + 3O2 → 2CO2 + 2H2O​ Exothermic O2: O=O CO2: O=C=O H2O: H–O–H C2H4: 8411 KJ/mol Total ∆H: Exothermic or endothermic: Exothermic 3.​ Analyze: In all of the reactions you investigated today, did it make any difference in which order you broke bonds or formed bonds? No ​ ​ ​ ​ Explain your answer. ​ ​ ​ ​ ​ ​ ​ ​ ​ According to Hess’s law, the overall enthalpy change depends only on the initial and final stakes. ​ ​ ​ ​ ​ ​ According to Hess’s law, the total enthalpy change for a reaction is simply the sum of all of the individual enthalpy changes in the reaction. The order in which you add these enthalpy changes does not matter. ​Reproduction for educational use only. Public sharing or posting prohibited. © 2020 ExploreLearning™ All rights reserved