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Thermochemistry

Thermochemistry. Chapter 17. Introduction. Thermochemistry is the chemistry associated with heat. Heat (q) is a form of energy that flows. Heat flow is a measurable quantity. Objects have different rates at which they absorb heat.

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Thermochemistry

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  1. Thermochemistry Chapter 17

  2. Introduction • Thermochemistry is the chemistry associated with heat. • Heat (q) is a form of energy that flows. • Heat flow is a measurable quantity. • Objects have different rates at which they absorb heat. • Objects also store and release a specific amount of heat, that is measurable. • Calorimetry is the technique used to measure heat flow.

  3. The Flow of Energy – Heat and Work(Section 17.1) • Energy Transformations • Exothermic & Endothermic Processes • Units for Measuring Heat Flow • Heat Capacity and Specific Heat

  4. I.) Energy Transformations • Energy = capacity to do work or supply heat. • Energy can be changed from one form to another. • There are different types of energy transformations.

  5. Chemical bonds also have potential energy. • During a chemical reaction a substance changes into another substance with a different PE in its bonds. • Thermochemistry is the study of energy changes during a chemical reaction and a change in state. • Energy changes occur as either heat transfer or work or a combination of both.

  6. What is heat? • Heat (q): Energy that transfers from one object to another because of temperature difference between them. • Heat always flow from a warmer object to a colder one until both objects are at the same temperature.

  7. II.) Exothermic & Endothermic Processes • Chemical reactions or changes in physical states involve either the absorption or release of heat. • Though energy moves from one place to another, overall energy does not change. • Law of Conservation of Energy • We can describe the movement of heat as either exothermic or endothermic depending upon our perspective.

  8. Chemical perspective: System and surroundings. • System: The part of the universe on which you focus your attention. • Surroundings: Everything else in the universe. • System + Surroundings = Universe

  9. Endothermic Processes • A process in which the system absorbs heat from the surroundings. • In these processes the system heats up while the surroundings cool down.

  10. Exothermic processes • A process in which the system releases heat to the surroundings. • In these processes the system loses heat as the surroundings heat up.

  11. Determining direction of heat flow. • Heat flowing out of a system = -q • Heat flowing into a system = +q

  12. Sample problem On a sunny winter day, the snow on a rooftop begins to melt. As the melt-water drips from the roof, it refreezes into icicles. Describe the direction of heat flow as the water freezes. Is the process endothermic or exothermic?

  13. Sample Problem Classify the following processes as exothermic or endothermic. • Condensing steam • Evaporating alcohol • Burning alcohol • Baking potato

  14. III.) Units for Measuring Heat Flow • Heat flow is measured in two common units, the calorie and the joule. • calorie(cal): The quantity of heat needed to raise the temperature of 1 g. of pure water 1oC. • Joule (J): The SI unit of energy. 1 J = 0.2390 cal 4.184 J = 1 cal

  15. IV.) Heat Capacity and SpecificHeat • Heat capacity: The amount of heat required to increase the temperature of an object exactly 1oC. • Heat capacity is dependent upon the mass and chemical composition of an object. • The greater the mass of an object the greater its heat capacity.

  16. Specific heat. • The specific heat capacity of a substance is the amount of heat it takes to raise the temperature of 1 g. of the substance 1oC.

  17. Calculating specific heat. q • C = specific heat • q = heat (joules or calories) • M = mass (grams) • ∆T = Change in temperature (Tf – Ti) • Tf = Final temperature • Ti = Initial temperature C = m x ∆T

  18. Sample problem The temperature of a 95.4 g. piece of copper increases from 25oC to 48.0oC when the copper absorbs 849 J of heat. What is the specific heat of copper?

  19. Measuring and Expressing Enthalpy Changes(Section 17.2) • Calorimetry • Thermochemical Equations

  20. I.) Calorimetry • The precise measurement of the heat flow into and out of a system for chemical and physical processes. • These measurements can be done at constant pressure or constant volume.

  21. Constant pressure calorimetry: coffee cup calorimetry • Most chemical reactions and physical changes are carried out under constant pressure. • The heat content of a system at constant pressure is the same as the enthalpy (H) of a system.

  22. Calculating enthalpy changes (∆H). • The heat released or absorbed by a reaction at constant pressure is the same as the change in enthalpy. • Therefore q = ∆H. • If q = m x C x ∆T, • And qsystem = -qsurroundings, • Then qsystem = ∆H = -qsurroundings = -m x C x ∆T • +∆H = endothermic reactions • - ∆H = exothermic reactions

  23. Sample problem When 25 mL of water containing 0.025 mol HCl at 25oC is added to 25.0 mL of water containing 0.025 mol NaOH at 25.0oC in a foam cup calorimeter, a reaction occurs. Calculate the enthalpy change (in kJ) during the reaction if the highest temperature observed is 32.0oC. Assume the densities of the solutions are 1.00 g/mL.

  24. Constant volume calorimetry: bomb calorimetry. • In this type of calorimetry a sample is burned at high pressure. • The heat released warms the water surrounding the chamber. • Measuring the temperature difference allows for the calculation of heat released.

  25. II.) Thermochemical Equations • Reactions can release or absorb heat. • The amount of heat absorbed or released can be treated as a reactant or product.

  26. Heat of Reaction • The heat of reaction (ΔH) is the enthalpy change for the chemical reaction exactly as it is written.

  27. Problem Solving Involving ΔH • These problems are solved in a manner similar to stoichiometry problems. • ΔH depends on the moles of products and reactants involved. • Again the physical states of the reactants and products must be stated.

  28. Sample problem When carbon disulfide is formed from its elements, heat is absorbed. Calculate the amount of heat (in kJ) absorbed when 5.66 g. of carbon disulfide id formed. (ΔH for the reaction is 89.3 kJ)

  29. Heat in Changes of State(Section 17.3) • Heats of Fusion • Heats of Vaporization

  30. I.)Heats of Fusion & Solidification • Phase changes also require heat. • No increase in temperature. • Melting: Endothermic Process • Molar heat of fusion (ΔHfus): The heat absorbed by one mole of a solid substance as it melts to a liquid. • Freezing: Exothermic Process • Molar heat of solidification(ΔHsolid): The heat lost when one mole of a liquid solidifies.

  31. Melting is the same as freezing. • The quantity of heat absorbed by a melting solid is exactly the same as the quantity of heat released when the liquid solidifies. • ΔHfus= ΔHsolid ΔHfus= ΔHsolid = 6.01 kJ

  32. Sample Problem How many grams of ice at 0o C will melt if 2.25 kJ of heat are added?

  33. II.)Heats of Vaporization and Condensation • Vaporization: Endothermic Process • Molar heat of vaporization (ΔHvap): The heat absorbed by one mole of a liquid substance as it vaporizes. • Condensation: Exothermic Process • Molar heat of condensation(ΔHcond): The heat released when one mole of a vapor condenses.

  34. Vaporizing is the same as condensing. • The quantity of heat absorbed by vaporizing a liquid is exactly the same as the quantity of heat released when the gas condenses. • ΔHvap= ΔHcond ΔHfus= ΔHsolid = 40.7 kJ

  35. Heating Curve for Water

  36. Sample Problem How much heat (in kJ) is absorbed when 24.8 g of liquid water at 100o C and 101.3 kPa is converted to steam at 100o C?

  37. Calculating Heats of Reaction(Section 17.4) • Hess’s Law • Standard Heats of Formation

  38. I.) Hess’s Law If you add two or more thermochemical equations to give a final equation, then you can also add the heats of reaction to give the final heat of reaction. • Hess’s law allows us to determine the heat of reaction indirectly. • Calculating heats of reaction using Hess’s Law requires specific steps.

  39. Calculating for heats of reaction. • Determine the reactants and products of the desired reaction. • Find reactions for which enthalpy values are known. • Set up the reactions such that the desired reactants are on one side and the desired products are on the other side of the arrow. • If a reaction is reversed then the sign of the enthalpy changes must be reversed.

  40. Sample Problem Find the enthalpy change for the conversion of diamond to graphite by using the combustion reaction for each.

  41. II.) Standard Heats of Formation • Enthalpy changes depend upon the conditions of the process. • Standard state: 25o C, 101.3 kPa, physical state of compound/element at 25o C. Standard Heat of Formation (ΔHf0) : The change in enthalpy that accompanies the formation of one mole of a compound from its elements with all the substances in their standard states.

  42. Using Standard Heats of Formations • ΔHf0 provide an alternative to Hess’s Law in determining heats of reactions. • For a reaction that occurs at standard conditions, we can use the ΔHf0 values to determine the heat of reaction (ΔH0) .

  43. Sample Problem What is the standard heat of reaction for the reaction of CO(g) with O2(g) to form CO2 (g)?

  44. Thermochemistry Chapter 17 The End

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