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17-1 Thermochemistry

17-1 Thermochemistry. Transfers of energy as heat in chemical reactions and physical changes. Remember…. The First Law of Thermodynamics: energy cannot be created or destroyed only converted from one form to another. Heat vs.Temperature. Heat (Joule, J)

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17-1 Thermochemistry

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  1. 17-1 Thermochemistry Transfers of energy as heat in chemical reactions and physical changes

  2. Remember… The First Law of Thermodynamics: energy cannot be created or destroyed only converted from one form to another

  3. Heat vs.Temperature • Heat (Joule, J) • measure of energy change in a system. • Temperature (Celsius, °C or Kelvin, K) • measure of the kinetic energy (movement) of the particles in a system. • Gaining or losing heat energy in a substance can change its temperature. • Exothermic • System loses energy to surroundings • Endothermic • System gains energy from surroundings

  4. Specific Heat • is a property of matter, and different species have different Specific Heat. • The heat energy required to raise one gram of a pure substance 1° C or 1 K • The symbol we use is cp

  5. Specific Heat • Metals have very low cp, • which is why metals often feel cold to the touch. • Table on page 513 • Water has a very highcp, • 4.184 J/g·0C • Substances with lower cp will rise in temperature faster and require less energy to do so than do substances with high cp.

  6. Specific Heat 1 calorie= 4.184 Joules

  7. Specific Heat Equation cp = q/(m ΔT) OR q = cp x m x ΔT

  8. Heat of Reaction • Quantity of energy released or absorbed during a chemical reaction • Thermochemical equation – shows the quantity of heat • Example: 2H2 + O2 2H2O + 483.6 kJ • Energy released or absorbed?

  9. Heat of Reaction • Example: 2H2 + O2 2H2O + 483.6 kJ • Energy released (on product side) • EXOTHERMIC • What about when it’s on the reactant side?

  10. Heat of Reaction • Example: 2H2O + 483.6 kJ  2H2 + O2 • Energy absorbed (on reactant side) • ENDOTHERMIC

  11. Enthalpy Change (ΔH) • The amount of energy absorbed or lost ΔH = Hproducts - Hreactants • Thermochemical equations usually written this way 2H2 + O2 2H2O ΔH = -483.6 kJ/mol • When ΔH is negative, the system loses energy and it is EXOTHERMIC

  12. Enthalpy Change (ΔH) • The amount of energy absorbed or lost ΔH = Hproducts - Hreactants 2H2O  2H2 + O2 ΔH = +483.6 kJ/mol • When ΔH is positive, the system gains energy and it is ENDOTHERMIC

  13. Endothermic or Exothermic? C6H12O6 + 6O26CO2 + 6H20 ΔHrxn = -2870 kJ/mol

  14. Heat of Formation (ΔHf) • energy released or absorbed to form 1 mole of a compound from its elements • p. 902 • If a large amount of energy is released when compound is formed… • Endothermic or exothermic? • Positive or Negative ΔHf?

  15. Heat of Formation (ΔHf) • energy released or absorbed to form 1 mole of a compound from its elements • p. 902 • If a large amount of energy is released when compound is formed… • Endothermic or exothermic? • Positive or NegativeΔHf? • HIGH NEGATIVE ΔHf= VERY STABLE!

  16. Heat of Formation (ΔHf) • POSITIVE ΔHf= UNSTABLE! • Pure elements ΔHf= O • ΔHfof carbon dioxide = -393.5 kJ/mol • More stable than C and O alone • HgC2N2O2ΔHf= +226.7 kJ/mol • Very unstable, used in explosives

  17. Heat of Combustion (ΔHc) • Reactants? • Products? • Exothermic or Endothermic? • Positive or Negative?

  18. Heat of Combustion (ΔHc) • Reactants? C and H with O2 • Products? CO2 and H2O and heat and light • Exothermic or Endothermic? • Positive or Negative? • P. 896

  19. Stability of these compounds? • Al2O3 (s) -1676.0 kJ/mol • CaCO3 (s) -1206.92 kJ/mol • NO (g) 90.29 kJ/mol • O3 142.7 kJ/mol

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