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Thermochemistry

Thermochemistry. Exothermic reactions release heat to the surroundings. Fe 2 O 3 + 2 Al  2 Fe + Al 2 O 3 + 851.5 kJ. Or ΔH = -851.5 kJ. Thermite Reaction Demo. Potassium Permanganate Reaction Demo. Thermochemistry. Endothermic reactions absorb heat from the surroundings.

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Thermochemistry

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  1. Thermochemistry • Exothermic reactions release heat to the surroundings. • Fe2O3 + 2 Al  2 Fe + Al2O3 + 851.5 kJ Or ΔH = -851.5 kJ Thermite Reaction Demo Potassium Permanganate Reaction Demo

  2. Thermochemistry • Endothermic reactions absorb heat from the surroundings. • Ba(OH)2 + 2 NH4NO3 + 102.2 kJ  Ba(NO3)2 + 2 NH3 + 10 H2O Or ΔH = +102.2 kJ Barium Hydroxide Reaction Demo

  3. [http://cwx.prenhall.com/bookbind/pubbooks/hillchem3/medialib/media_portfolio/text_images/CH06/FG06_15.JPG][http://cwx.prenhall.com/bookbind/pubbooks/hillchem3/medialib/media_portfolio/text_images/CH06/FG06_15.JPG]

  4. Enthalpy • Energy that is gained or lost by substances during a reaction. • Symbolized by the letter H. [http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/FG07_14.JPG]

  5. Enthalpy • Example: • Given equal amounts (mass) of both, which produces more energy: Methane (Natural Gas!) or Octane (Gasoline!)

  6. Enthalpy • Combustion of Methane (Natural Gas!) • CH4(g) + O2(g)  CO2(g) + H2O(g) • ΔH = -890.4 kJ (Exothermic!) 2 2

  7. Enthalpy CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g) ΔH = -890.4 kJ Your furnace turns on and burns 32.0 g of methane, how much heat is produced? 32.0 g CH4 1 mol CH4 -890.4 kJ = -1780 kJ 16.0 g CH4 1 mol CH4 Taken from reaction ratio!

  8. Enthalpy • Combustion of Octane (Gasoline!) • C8H18(l) + O2(g)  CO2(g) + H2O(g) • ΔH = -5430 kJ (Exothermic!) 2 16 25 18

  9. Enthalpy 2 C8H18(l) + 25 O2(g)  16 CO2(g) + 18 H2O(g) ΔH = -5430 kJ Your car turns on and burns 32.0 g of octane, how much heat is produced? 32.0 g C8H18 1 mol C8H18 -5430 kJ 114.0 g C8H18 2 mol C8H18 = -762 kJ Taken from reaction ratio!

  10. Enthalpy • Methane produces more energy when burned…as long as mass is constant.

  11. Enthalpy of Formation • This is the amount of energy (enthalpy) that is involved in “creating” a compound. [http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/TB07_02.JPG]

  12. Enthalpy of Formation • The “ΔH” given in a reaction is calculated by using these values in this equation:

  13. Enthalpy of Formation • Example: • CH4 (g) + O2 (g)  CO2 (g) + H2O (l) • What is ∆H for this reaction? • 1st Step - Balance the Equation 2 2

  14. Enthalpy of Formation 2nd Step – Find Enthalpy Values Using Chart CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l) -75 kJ/mol -393.5 kJ/mol -286 kJ/mol Any element that is “pure”, has an enthalpy of 0 kJ/mol

  15. Enthalpy of Formation • 3rd – Plug into ΔH equation and solve: ΔH = (Hproducts) – (Hreactants) ΔH = [ (-286) + (-393.5)] – [(-74.8) + 0] 2 Coefficient from balanced equation!

  16. Enthalpy of Formation • Final Answer: -890.7 kJ/mol • No significant digits needed since all the numbers are found on tables!

  17. Hess’s Law • A series of reactions can be added together to find their overall enthalpy. N2 + O2 2 NO ΔH = +181 kJ 2 NO + O2 2 NO2 ΔH = -131 kJ + N2 + 2 O2 2 NO2 ΔH = +68 kJ

  18. Hess’s Law • Example: • C2H2 + H2 C2H4ΔH = -174.4 kJ • C2H6 C2H4 + H2ΔH = 137.0 kJ • What is the enthalpy for the net reaction: • C2H2 + 2 H2  C2H6

  19. If you cannot simply add to find the net reaction, “flip” a reaction to make it work! ΔH = -174.4 kJ C2H2 + H2 C2H4 C2H4 + H2  C2H6 ΔH = -137.0 kJ C2H6 C2H4 + H2 ΔH = 137.0 kJ + C2H2 + 2 H2 C2H6 ΔH = - 311.4kJ

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