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Chapter 6: Thermodynamics

Chapter 6: Thermodynamics. Calorimetry and Hess’s Law. Calorimetry. Process of measuring quantities of heat Reactions are carried out in a calorimeter and the changes in temperature equate to changes in energy for the reaction . Heat Capacity.

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Chapter 6: Thermodynamics

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  1. Chapter 6: Thermodynamics Calorimetry and Hess’s Law

  2. Calorimetry • Process of measuring quantities of heat • Reactions are carried out in a calorimeter and the changes in temperature equate to changes in energy for the reaction

  3. Heat Capacity • Quantity of heat required to change the temperature of the system by 1oC (or by 1K) • C = q / DT • Units = J/oC or J/K

  4. Molar heat capacity & specific heat • Molar heat capacity is the heat capacity of one mole of a substance • Specific heat is the heat capacity of a 1-g sample Specific heat = heat capacity/mass = C/m

  5. Specific heats of some substances at 25oC

  6. Calculating specific heat and q • Specific heat = q/mDT • Q = mCpDT M =mass C =specific heat DT =change in temperature

  7. Example 6.7 • How much heat, in joules and in kilojoules does it take to raise the temperature of 225 g of water from 25.0 to 100.0oC?

  8. Example 6.8 • What will be the final temperature if a 5.00 g silver ring at 37.0 oC gives off 25.0 J of heat to its surroundings? Specific heat for silver is 0.235 J g-1 oC-1

  9. Measuring Specific Heats • Coffee Cup calorimeters • If hot solids are placed into water into a insulated cup, the heat lost by the solid is gained by the water in the cup.

  10. Coffee Cup Calorimeter

  11. Example 6.9 • A 15.5 g sample of metal alloy is heated to 98.9oC and then dropped into 25.0 g of water in a calorimeter. The temperature of the water rises from 22.5oC to 25.7oC. Calculate the specific heat of the alloy.

  12. Measuring Enthalpy Changes for Chemical Reactions • If we measure the a heat of reaction in an isolated system, the change in chemical energy will appear solely as a change in the thermal energy of the system.

  13. Example 6.11 • A 50.0 ml sample of 0.250M NaCl at 19.50oC is added to 50.0 mL of 0.250 M NaOH, also at 19.50oC in a calorimeter. After mixing the solution temperature rises to 21.1oC. Calculate the heat of reaction.

  14. Bomb Calorimetry: Reactions at Constant Volume • Bomb calorimeters are needed for combustion reactions and other reactions involving gases. • To complete calculations with Bomb Calorimeters the heat capacity for the bomb calorimeter must also be integrated into the calculations

  15. Bomb Calorimeter

  16. Example 6.13 • In a preliminary experiment, the heat capacity of a bomb calorimeter assembly is found to be 5.15 kJ/oC. In a second experiment, a 0.480 g sample of graphite is placed in the bomb with an excess of oxygen. The water, bomb, and other contents of the calorimeter are in thermal equilibrium at 25.00oC. The graphite is ignited and burned, and the water temperature rises to 28.05oC. Calculate DH for the reaction.

  17. Hess’s Law of Constant Heat Summation • The heat of a reaction is constant, whether the reaction is carried out directly in one step or indirectly through a number of steps.

  18. Example 6.14 • Calculate the enthalpy change for reaction (a) given the data in equations (b) (c) and (d). • (a) 2C(graphite) + 2H2(g)  C2H4(g) DH =? • (b)C (graphite) + O2(g)  CO2(g) DH= -393.5kJ • (c) C2H4(g) + 3O2(g)  2CO2(g) + 2H2O(g) DH= -1410.9kJ • (d) H2(g) + 1/2O2(g)  H2O(l) DH = -285.8kJ

  19. Standard Enthalpies of Formation • DH = Hproducts – Hreactants • DHo= the enthalpy change for a reaction in which the reactants in their standard states yield products in their standard states • DHofis the enthalpy change that occurs as 1 mole of the substance forms from its elements when both products and reactants are in their stand states

  20. Standard Enthalpies of Formation

  21. Example 6.15 • Synthesis gas is a mixture of carbon monoxide and hydrogen that is used to synthesize a variety of organic compounds. One reaction for producing synthesis gas is 3CH4(g) + 2H2O(l) + CO2(g) 4CO(g) + 8H2(g) Use standard enthalpies of formation form Table 6.2 to calculate the standard enthalpy change for this reaction.

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