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Unit 3: Gases Thermochemistry

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Unit 3: Gases Thermochemistry

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    1. Unit 3: Gases & Thermochemistry Chapters 5 & 6

    3. Pressure Calculated using the equation: Units include: Pressure exerted by gases in the atmosphere can be demonstrated by boiling water in a large metal can, then turning off the heat and sealing the can.

    4. Yikes!

    5. Measuring Pressure Barometers Manometers

    6. Atmospheric Pressure

    7. Measuring Pressure

    8. Variables That Describe a Gas

    9. Boyle’s Law

    10. Charles’ Law

    11. Gay-Lussac’s Law

    12. Combined Gas Law

    13. Sample Problem 5.2

    14. Avogadro’s Law At the same temperature and pressure, equal ______________ of gases have equal number of _____________.

    15. Ideal Gas Law

    16. Sample Problems 5.3

    17. Density of a Gas

    18. Molar Mass of a Gas

    19. Gas Stoichiometry Same 4 steps! The only twist:

    20. Gas Stoich

    21. Dalton’s Law The total pressure in a mixture of gases is equal to the sum of the parts. The partial pressure of one gas in a mixture is dependent on its mole fraction.

    22. KMT

    23. Molecular Speed of a Gas

    24. Diffusion and Effusion

    25. Graham’s Law

    26. Real Gases

    27. Thermochemistry Chapter 6

    28. Key Terms Energy 1st Law of Thermodynamics Heat vs. Temperature Enthalpy System Surroundings Exothermic Endothermic

    29. Energy Diagrams

    30. Heat Transfer

    31. A Look at Enthalpy

    32. More About Enthalpy

    33. Calorimetry The science of measuring heat based on observing temperature changes in substances. The amount of heat absorbed/emitted also depends on the substance itself: __________________ is the amount of heat needed to raise the temp of a substance by 1oC.

    34. Calorimetry Calorimeter is used to measure the changes in temp. Then the relationship between heat, temp, mass, and specific heat can be described by the equation:

    35. Determining Specific Heat Suppose a 55.0 g piece of metal was heated in boiling water to 99.8oC and then dropped into water in an insulated beaker. There is 225 mL of water (D = 1 g/mL) in the beaker, and its temperature before the metal was dropped in was 21.0oC. The final temperature of the metal and the water is 23.1oC. What is the specific heat of the metal? Assume no heat transfers through the walls of the beaker or to the atmosphere.

    36. Coffee Cup Calorimetry Suppose you place 0.500 g of magnesium chips in a coffee cup calorimeter and then add 100.0 mL of 1.00 M HCl. The temperature of the solution increases from 22.2oC to 44.8oC. What is the enthalpy change for this reaction in kJ per mol of Mg? (c = 4.184 J/goC and D = 1 g/mL)

    37. 3 Other Ways to Find the Enthalpy of a Reaction Hess’s Law Standard Enthalpies of Formation Bond Energies

    38. Hess’s Law Enthalpy changes are state functions Hess’s Law: The overall enthalpy change for a reaction is equal to the sum of its steps N2 + 2O2 ? 2NO2 DH = 68kJ In steps: N2 + O2 ? 2NO DH1 = 180 kJ 2NO + O2 ? 2NO2 DH2 = -112 kJ N2 + 2O2 ? 2NO2 DH = 68 kJ

    39. Hess’s Law Hints to Remember The sum of the steps must equal the overall equation in terms of the reaction itself. If you must reverse a step to get it to match the overall equation, the sign of DH is also reversed If you must multiply a step by a coefficient to get it to match the overall equation, DH must also be multiplied Work backward from the overall reaction, make sure reactants and products are on the correct sides and that coefficients match

    40. Standard Enthalpies of Formation DHfo accompanies the formation of one mole of a compound from its elements at standard states Standard state conditions: The form an element exists at 25oC and 1 atm For a compound Gases at 1 atm Solutions at 1M DHfo for an element equals zero

    41. Standard Enthalpies of Formation Appendix 4 lists the standard enthalpies of formation for many compounds Using this information, the DH of a reaction can be found using the equation: DHrxn = SnDHfo (products) – SnDHfo (reactants)

    42. Bond Energies Takes energy to break bonds and energy is released when bonds are formed. Since the sign indicates whether energy is being absorbed or released, we can use the equation: DHrxn = S E (bonds broken) – S E (bonds formed) Use table 3.4 on p. 351 for bond energies

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