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THERMODYNAMICS: REACTION ENERGY

Unit 12. THERMODYNAMICS: REACTION ENERGY. Tuesday, May 8 – EOC Field Testing for Chemistry Wednesday, May 9 – Unit 12 Test Review Thursday, May 10 – Unit 12 Test. Important Dates. The Flow of Energy Energy – the capacity to do work or supply heat

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THERMODYNAMICS: REACTION ENERGY

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  1. Unit 12 THERMODYNAMICS: REACTION ENERGY

  2. Tuesday, May 8 – EOC Field Testing for Chemistry Wednesday, May 9 – Unit 12 Test Review Thursday, May 10 – Unit 12 Test Important Dates

  3. The Flow of Energy • Energy – the capacity to do work or supply heat • Chemical Potential Energy – energy stored within the bonds of chemical compounds • Activation Energy – the minimum energy colliding particles must have in order to react Day 1 Notes

  4. Thermodynamics – the study of energy in chemical reactions; literally means “changes in heat.” • All chemical reactions either release or absorb energy when they occur. Another way to say this is that reactions either “give off” or “take in” energy. • The Law of Conservation of Energy – 1st Law of Thermodynamics; states energy can neither be created nor destroyed, only transformed • ALL energy is either work performed, stored potential, or heat lost. THERMODYNAMICS

  5. System – the chemical reaction under study • Surroundings – every place in the universe except the system • Universe – the system and the surroundings Reaction system, surroundings, and universe

  6. Reaction system, surroundings, and universe SYSTEM Surroundings Universe

  7. Exothermic Reaction – a reaction in which heat is released by the system to the surroundings; from the perspective of the system, “heat is given off” Types of Heat Flow

  8. Endothermic Reaction - is a reaction in which heat is absorbed by the system from the surroundings. From the perspective of the system, “heat is taken in.” Types of Heat Flow

  9. Qualitative Practice – Heat FlowMatch the reaction diagram with the appropriate description of the flow of energy. Energy level diagram for an exothermic chemical reaction without showing the activation energy; it could also be seen as quite exothermic with a highly unlikely zero activation energy, but reactions between two ions of opposite charge usually has a very low activation energy. Very endothermic reaction with a large activation energy. Moderately exothermic reaction with a moderately high activation energy. A small activation energy reaction with no net energy change; this is theoretically possible if the total energy absorbed by the reactants in bond breaking equals the energy released by bonds forming in the products. Very exothermic reaction with a small activation energy. Energy level diagram for an endothermic chemical reaction without showing the activation energy; it could also be seen as quite endothermic with zero activation energy. Moderately endothermic reaction with a moderately high activation energy.

  10. How to read the graph! Potential Energy – Read from the x-axis to the reaction line EX: b is the potential energy of ….. (answer to question 2) A+B and C+D are NOT answers on the questions. They represent the reactants (A+B) and the products (C+D). It might help to draw dashed lines so that you can visualize the potential energy all the way across the graph Activation Energy Reactants Products

  11. Is the above reaction endothermic or exothermic? What letter represents the potential energy of the reactants? What letter represents the potential energy of the products? What letter represents the change in energy forthe reaction? Assessment – potential energy diagram b f d

  12. What letter represents the activation energy of the forward reaction? What letter represents the activation energy of the reverse reaction (read the chart backwards)? What letter represents thepotential energy of theactivated complex? Is the reverse reactionendo or exothermic? If a catalyst were added,what letter(s) would change? Assessment – potential energy diagram a e c a and c

  13. Day 2 THERMODYNAMICS: REACTION ENERGY

  14. Measuring and Expressing Heat Changes • Calorimetry– the accurate and precise measurement of heat change for chemical and physical processes • Calorimeter – the insulated device used to measure the absorption or release of heat in chemical or physical processes • Enthalpy (H) – heat energy content of a system at constant pressure • Enthalpy is the heat absorbed or released by a system when pressure is constant. • It is impossible to record enthalpy directly, but change in enthalpy (ΔH ) can be measured. • Units of heat energy: calorie (cal), joules (J), or kJ (kJ). Reaction Energy

  15. For an exothermic reaction, the sign of ΔH is negative. • When a reaction is exothermic (ΔH is negative), that is a favorablecondition. Enthalpy is just one of the variables involved when predicting whether or not a reaction will occur, but, in general, reactions which release heat are more likely to occur than ones in which heat is required. • Heat EXITS the system so the energy (in kJ or J) is shown as a PRODUCT • AB + CD  AD + BC + ΔH Reaction Energy – Exothermic Rxn

  16. For an endothermic reaction, the sign of ΔH is positive. • When a reaction is endothermic (ΔH is positive), that is an unfavorable condition. Enthalpy is just one of the variables involved when predicting whether or not a reaction will occur, but reactions which absorb heat are less likely to occur than ones in which heat is released, all things being equal. • Heat is PUT INTO the system so the energy (in kJ or J) is shown as a REACTANT • AB + CD + ΔH  AD + BC Reaction Energy – Endothermic Rxn

  17. Thermochemical Equation – an equation that includes the heat change • Ex. CaO(s) + H2O(l) → Ca(OH)2(s) + 65.2kJ • Heat Change as a reactant means endothermic • Heat Change as a product means exothermic • Heat of Reaction – the heat of change for the equation exactly as it is written • ΔH = positive means endothermic • ΔH = negative means exothermic Reaction energy, con’t Heat Change

  18. 2NO(g) + O2(g) 2NO2(g) + 113.04 kJ endothermic or exothermic • 2H2(g) + O2(g) 2H2O(l); ΔH = -571.6 kJ endothermic or exothermic • 4NO(g) + 6H2O(l) 4NH3(g) + 5O2(g);    ΔH = +1170 kJ endothermic or exothermic • SO2 (g) +296 kJ  S(s)+ O2 (g) endothermic or exothermic Example Problems - EnthalpyDefine the following examples as either endothermic or exothermic based upon the change in heat. Heat is shown as a product Heat is negative Heat is positive Heat is shown as a reactant

  19. Energy WS: Problem #6 Activation energy = 100 kJ/mol = 450 kJ/mol 400 Potential energy of reactants = 350 kJ/mol 300 Potential Energy in kJ Potential energy of products = 250 kJ/mol 200 EXOTHERMIC REACTION ΔH = - 100 kJ 100 Reaction Pathway (timeline)

  20. Potential Energy diagram WS 1. Which of the letters a–f in the diagram represents the potential energy of the products? _______ 2. Which letter indicates the potential energy of the activated complex? ____ 3. Which letter indicates the potential energy of the reactants? ________ 4. Which letter indicates the activation energy? _____ 5. Which letter indicates the heat of reaction? ______ 6. Is the reaction exothermic or endothermic? ______ 7. Which letter indicates the activation energy of the reverse reaction? ________ 8. Which letter indicates the heat of reaction of the reverse reaction? ________ 9. Is the reverse reaction exothermic or endothermic? __________ e c a b f endo d f exothermic

  21. Reading a chart with numbers! 80 1. The heat content of the reactants of the forward reaction is about ______ kJ. 2. The heat content of the products of the forward reaction is about _______kJ. 3. The heat content of the activated complex of the forward reaction is about ______ kJ. 4. The activation energy of the forward reaction is about ______ kJ. 5. The heat of reaction (ΔH) of the forward reaction is about ______ kJ. 6. The forward reaction is _______________ (endothermic or exothermic). 7. The heat content of the reactants of the reverse reaction is about ________ kJ. 8. The heat content of the products of the reverse reaction is about _______ kJ. 9. The heat content of the activated complex of the reverse reaction is about _______kJ. 10. The activation energy of the reverse reaction is about _______ kJ. 11. The heat of reaction (ΔH) of the reverse reaction is about _______ kJ. 12. The reverse reaction is __________________ (endothermic or exothermic). 160 240 160 +80 endothermic 160 80 240 80 - 80 exothermic

  22. Day 3 THERMODYNAMICS: REACTION ENERGY

  23. Calculating Heat Changes • Standard Heat of Formation (ΔHf0) – the change in enthalpy that accompanies the formation of one mole of a compound from its elements • Heat of Reaction (ΔH0)– the heat released or absorbed during a chemical reaction, or Enthalpy • ΔH0 = ΔHf0 (products) - ΔHf0 (reactants) • READ AS: Heat of Rxn EQUALS the SUM of Heat of Formation of the Products minus the SUM of Heat of Formation of the Reactants • Standard Heats of Formation have been determined for many common pure substances, both elements and compounds. Elements in their natural state are understood to have a ΔHf0 = 0 Reaction Energy continued, Heat of Formation and Reaction

  24. In order to calculate ΔH0, the standard heats of formation of the reactants and products must be known; they can be found on the following table: Heat of Formation and Reaction

  25. Steps to Calculate Heat of Reaction: • Find the balanced chemical equation for the reaction; must have states of matter • Find the sum of the Heats of Formation for the reactants • Multiply the ΔHf0 for each reactant by its corresponding number of moles (coefficient) from the balance equation • Sum the reactants • Find the sum of the Heats of Formation for the products • Multiply the ΔHf0 for each product by its corresponding number of moles (coefficient) from the balance equation • Sum the products • Subtract the ΔHf0 (reactants) from the ΔHf0 (products) Calculating heat of Rxn

  26. Find the balanced chemical equation for the reaction; must have states of matter • Find the sum of the Heats of Formation for the reactants • Multiply the ΔHf0 for each reactant by its corresponding number of moles (coefficient) from the balance equation • Sum the reactants • Find the sum of the Heats of Formation for the products • Multiply the ΔHf0 for each product by its corresponding number of moles (coefficient) from the balance equation • Sum the products REMEMBER! Elements in their natural state are understood to have a ΔHf0 = 0 OTHERWISE, use the table provided to lookup individual ΔHf0 Example: 2CO(g) + O2(g) 2CO2(g)

  27. Subtract the ΔHf0 (reactants) from the ΔHf0 (products) ΔH0 = ΔHf0 (products) - ΔHf0 (reactants) (show work here) Example: 2CO(g) + O2(g) 2CO2(g)

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