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Chemistry 30 – Unit 1 Thermochemical Changes

To accompany Inquiry into Chemistry. PowerPoint Presentation prepared by Robert Schultz robert.schultz@ei.educ.ab.ca. Chemistry 30 – Unit 1 Thermochemical Changes. Preparation Info. Systems: Open, closed, and isolated - definitions

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Chemistry 30 – Unit 1 Thermochemical Changes

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  1. To accompany Inquiry into Chemistry PowerPoint Presentation prepared by Robert Schultz robert.schultz@ei.educ.ab.ca Chemistry 30 – Unit 1Thermochemical Changes

  2. Preparation Info • Systems: Open, closed, and isolated - definitions • First Law of Thermodynamics – Total energy of the universe is constant (energy can’t be created or destroyed) • Second Law of Thermodynamics – In the absence of energy input, a system becomes more disordered (its entropy increases)

  3. Preparation • Meaning? • A system at lower temperature will be more ordered as the particles have less average kinetic energy • Two systems in thermal contact will transfer energy such that the more ordered (cooler) one gains energy and becomes more disordered • Consequence: heat always flows from hotter systems to cooler ones

  4. Preparation Important Definitions: • Thermal Energy: the total kinetic energy of all particles of a system • Temperature: a measure of the average kinetic energy of the particles of a system • Heat: a transfer of thermal energy between 2 systems

  5. Chapter 9, Section 9.1 Questions: • Which has more thermal energy, a hot cup of coffee or an iceberg? • Which has a larger average thermal energy, a hot cup of coffee or an iceberg? • If an iceberg and a hot cup of coffee come into contact, in which direction will heat flow?

  6. Preparation • Heat energy transferred will be related to the temperature change of the system • It takes different amounts of heat energy to change the temperature of 1 g of a substance by 1°C • This number is called the specific heat capacity, c, and is measured in units of:

  7. Preparation • Water has a c value of • This means that it takes 4.19 J of heat to raise the temperature of 1 g of water by 1°C • Water has a very large c compared to most other common substances

  8. Preparation • To determine the amount of heat transferred the formula used is • Despite what your text says on page 337, I would always take ∆t as positive • If heat is absorbed, temperature of surroundings will decrease; if heat is released temperature of surroundings will increase • Examples: Practice Problems 1 and 4, page 337

  9. Preparation • Practice Problem 1, page 337 • Since 1 J is such a small amount of heat energy I start my questions in kJ as shown above • If necessary I move into MJ or GJ

  10. Preparation • Practice Problem 4, page 337 • Putting kilo top and bottom cancels out and c stays the same • The substance is granite • Worksheet: WS 43 (Nelson) then BLM 9.1.1 (back only)

  11. Chapter 9, Section 9.1 • Energy changes in chemical reactions crucial to life • Not just in photosynthesis, fuels, and batteries, but in the very way that your body metabolizes food and makes the energy available for life processes • Thermodynamics: the study of energy and energy changes

  12. Chapter 9, Section 9.1 • Recall the first law of thermodynamics: ∆Euniverse= 0 • If a system loses energy, the surroundings gain energy (get warmer) • If a system gains energy, the surroundings lose energy (get cooler)∆Esystem = - ∆Esurroundings

  13. Chapter 9, Section 9.1 • Energy types: • Kinetic energy, Ek, energy of motion of particles of a system • Temperature is a measure of the average Ek of the particles of a system • Potential energy, Ep, stored energy, usually in chemical bonds

  14. Chapter 9, Section 9.1 • Transfer of Ek: heat flows from hotter objects to cooler ones (Preparation section of notes) • Breaking bonds always requires energy (endothermic); forming bonds always releases energy (exothermic) • Chemical reaction: breaking bonds + energy1 forming bonds + energy2 • If energy1 > energy2, reaction is endothermic • If reverse is true, it is exothermic • Worksheet BLM 9.1.3 input output

  15. Chapter 9, Section 9.1 • New term: enthalpy (not entropy) • Enthalpy (change), ∆H: the difference in potential energy between reactants and products, measured at constant pressure – measured in kJ (or MJ, etc) • Molar Enthalpy (change), ∆rH: the enthalpy change for 1 mole of a specified substance – measured inkJ/mol (or MJ/mol etc) • In common usage the word change gets left out

  16. Chapter 9, Section 9.1 • Negative∆H’s are exothermic (think lose heat) and temperature of surroundings increases • Positive∆H’s are endothermic (think gain heat) and temperature of the surroundings decreases • Note: this increase → negative, and decrease → positive is a stumbling block for many students

  17. Chapter 9, Section 9.1 • Chemical reactions can be written using ∆H notation: C6H12O6(s) + 6 O2(g) 6 CO2(g) + 6 H2O(l) ∆H=-2802.5 kJ 4 NO(g) + 6 H2O(g) 4 NH3(g) + 5 O2(g)∆H=+906 kJ • They can also be written with the heat as a term in the equation: C6H12O6(s) + 6 O2(g) 6 CO2(g) + 6 H2O(l) + 2802.5 kJ 4 NO(g) + 6 H2O(g) + 906 kJ 4 NH3(g) + 5 O2(g) Do ∆H Worksheet! value for the reaction as written

  18. C6H12O6(s) + 6 O2(g) reactants H (kJ) H (kJ) 4 NO(g) + 6 H2O(g) ∆H = -2802.5 kJ products ∆H = +906 kJ 4 NH3(g) + 5 O2(g) 6 CO2(g) + 6 H2O(l) reactants products Chapter 9, Section 9.1 • Potential energy diagrams for the same 2 reactions are shown below:

  19. Chapter 9, Section 9.2 • Recalling that breaking bonds always endothermic and forming new bonds is always exothermic, more complete Ep diagrams might be shown as follows: Endothermic Exothermic intermediate intermediate products Ep (kJ) Ep (kJ) reactants ΔH ΔH products reactants

  20. Chapter 9, Section 9.1 • Alternate forms of potential energy diagram (from Chemistry 30 Diploma Exam Bulletin)

  21. H (kJ) C(s) + 2 H2(g) reactants ∆H = -74.6 kJ CH4(g) products Chapter 9, Section 9.1 • Example: Practice Problem 3, page 346 • C(s) + 2 H2(g) CH4(g) + 74.6 kJ • C(s) + 2 H2(g) CH4(g) ∆H = -74.6 kJ c) Do Ep diagrams for formation of Cr2O3(s), simple decomp* of AgI(s), and formation of SO2(g)

  22. Formation of Cr2O3(s) 2 Cr(s) + 3/2 O2(g) Chapter 9, Section 9.2 Ep (kJ) ΔH=ˉ1139.7 kJ Cr2O3(s) reaction coordinate formation of SO2(g) simple decomposition of AgI(s) Ag(s) + ½ I2(s) 1/8 S8(s) + O2(g) Ep (kJ) Ep (kJ) ΔH=+61.8 kJ ΔH=ˉ296.8 kJ AgI(s) SO2(g) reaction coordinate reaction coordinate

  23. Chapter 9, Section 9.1 • Molar enthalpy of combustion: the enthalpy change for the complete combustion of 1 mol of a substance • Complete combustions of fossil fuels always yields CO2(g) and H2O • Open systems – constant pressure – gases escape – H2O(g) • Isolated systems – H2O(l) • Human body – cellular respiration - H2O(l)

  24. C4H10(g) + 13/2 O2(g) 4 CO2(g) + 5 H2O(g) ∆H = -2657.3 kJ 2 C4H10(g) + 13 O2(g) 8 CO2(g) + 10 H2O(g) ∆H = -5314.6 kJ Chapter 9, Section 9.1 • Table of Molar Enthalpies of Combustions of alkanes, page 347 • Practice Problem 5b, page 347 (open system) OR: note change in units! • In thermodynamics it is acceptable to write equations with fractional coefficients – don’t do this elsewhere • Try question 5a, page 347

  25. C5H12(l) + 8 O2(g) 5 CO2(g) + 6 H2O(g) ∆H = -3244.8 kJ Chapter 9, Section 9.1 • Question 5a page 347 • Note that the value of ∆H varies directly as the number of moles of reacting substances • This formula gets used to calculate enthalpy changes for ∆Ep like phase changes, chemical reactions, and nuclear reactions

  26. Note: from table, page 347 - comment mol of pentane Chapter 9, Section 9.1 • Example Practice Problem 3a, page 349

  27. 4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g) ΔH = -906 kJ Chapter 9, Section 9.1 • Example Practice Problem 6, page 349 • molar enthalpy change for? • a) ammonia • b) oxygen • c) nitrogen monoxide • d) water

  28. Chapter 9, Section 9.1 • Do Worksheet BLM 9.1.6

  29. Chapter 9, Section 9.2 • Finding the value of energy changes experimentally: calorimetry • Device: calorimeter • The following diagrams show the principle behind calorimetry – note arrow directions

  30. Chapter 9, Section 9.2 • A simple calorimeter like the one you will use 2 nested styrofoam cups containing a measured volume of water sitting in a beaker so that it doesn’t fall over 3rd styrofoam cup inverted on top with hole for thermometer (stirrer)

  31. Chapter 9, Section 9.2 • Assumptions in styrofoam cup calorimetry: • Amount of energy transferred to cups and thermometer is small and can be ignored • The system is isolated • The solution produced has the same density and specific heat capacity as water • The process occurs at constant pressure

  32. and calorimetry equation system calorimeter “water” Chapter 9, Section 9.2 • The enthalpy change of a chemical reaction = energy lost or gained, and is indicated by the symbol ΔH • Energy gained or lost by the water causes a temperature change and is indicated by the symbol Q • In an ideal calorimeter ΔH = Q • But recall: • Therefore

  33. Chapter 9, Section 9.2 limiting reagent, if not stated, or substance question asks about • I will redo the example on page 354 using this formula • remember m c Δt is for the “water” and n (c v) for the CuSO4(aq) Since the temperature has gone up the process is exothermic Correct answer:

  34. Since temperature increases, answer is correctly expressed as Chapter 9, Section 9.2 • Practice Problem 9, page 355 • Note that question asks for molar enthalpy of reaction for sodium • n will be moles of sodium (question asks) Do Practice Problems 7, 10, 12, page 355

  35. Chapter 9, Section 9.2 • Investigation 9.A page 356 (goes with the questions you’ve been doing) • Molar enthalpy of combustion: Investigation 9.B, page 357

  36. Chapter 9, Section 9.2 • Bomb Calorimetry: a bomb calorimeter is used to make accurate and precise measurements Not on diploma exam

  37. Chapter 9, Section 9.2 • Reaction takes place inside an inner container called the “bomb” that contains pure oxygen • Chemicals are electrically ignited and heat is released to or absorbed from calorimeter water • Calorimeter materials: stirrer, thermometer, containers are not ignored • With calorimeter filled to a set level with water, all of their heat capacities are combined as shown: Not on diploma exam

  38. Chapter 9, Section 9.2 bomb calorimeter equation • Note that C contains the mass and specific heat capacity of each component of the calorimeter • How do you know when to use Heat capacity of calorimeter Not on diploma exam

  39. Chapter 9, Section 9.2 • Look for:- words “bomb calorimeter”- no mention of the mass or volume of water- words “heat capacity” rather than “specific heat capacity”- units J/°C rather than J/g°C • Question 2, Worksheet 46 • Since temperature increases, answer is -286 kJ/mol • Do rest of Worksheet 46 Not on diploma exam

  40. Chapter 9, Section 9.2 • More practice with • WS 9.1.5

  41. Chapter 9, Section 9.2 • Review: page 366-7 good questions: 1, 3, 4 (no actual calculation needed), 5c (data page 347), 6a (data page 347), 8, 10, 13, 15, 16, 17, 18, 19, 21

  42. Chapter 9, Section 9.2

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