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Thermochemistry

Thermochemistry. Lesson 2. Introduction. Thermochemistry is the study of heat changes that accompany chemical reactions and phase changes. Enthalpy (H) is the heat content of a reaction at constant pressure.

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Thermochemistry

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  1. Thermochemistry Lesson 2

  2. Introduction • Thermochemistry is the study of heat changes that accompany chemical reactions and phase changes. • Enthalpy (H) is the heat content of a reaction at constant pressure. • Enthalpy is not constant, but changes in enthalpy such as heat lost or absorbed, can be used to determine the energy of a reaction. • The change in enthalpy for a reaction is called the enthalpy of reaction and is written as ΔHrxn. This is the difference between the enthalpy of the products and the enthalpy of the reactants. • The enthalpy values for exothermic reactions are always negative while the enthalpy values for endothermic reactions are always positive.

  3. Thermochemical Equations • A thermochemical equation is a balanced chemical equation that includes the physical states of all reactants and products and the energy change expressed as the change in enthalpy, ΔH. • C6H12O6 (s) + 6 O2 (g)  6 CO2 (g) + 6 H2O (l) ΔHcomb= -2808 kJ • The enthalpy of combustion of a substance is the enthalpy change for the complete burning of one mole of the substance. • ΔHocomb= standard conditions of 1 atmosphere and 298K

  4. Changes of State • Molar enthalpy of vaporization (ΔHvap) is the heat required to vaporize one mole of a liquid. • Molar enthalpy of fusion (ΔHfus) is the heat required to melt one mole of a solid substance. • Both of these processes are endothermic so their enthalpy values are positive. • The enthalpy values for condensation and solidification are negative values of vaporization and fusion.

  5. Standard Enthalpies of Vaporization and Fusion

  6. Calculating Enthalpy Change • Hess’s Law states that if you can add two or more thermochemical equations to produce a final equation for a reaction, then the sum of the enthalpy changes for the individual reactions is the enthalpy change for the final reaction.

  7. Example Problem • S (s) + O2 (g)  SO2(g) ΔH = -297 kJ • 2 SO3 (g)  2 SO2(g) + O2 (g) ΔH = 198 kJ Rearrange the equations: 2 S (s) + 2 O2 (g)  2 SO2 (g) 2( ΔH = -297 kJ) = -594 kJ 2 SO2 (g) + O2 (g)  2 SO3 (g) -1(ΔH = 198 kJ) = -198 kJ 2 S + 2 SO2+ 3 O2  2 SO2+ 2 SO3ΔH = -792 kJ 2 S (s) + 3 O2 (g)  2 SO3ΔH = -792 kJ

  8. Standard Enthalpy of Formation • The standard enthalpy of formation is the change in enthalpy that accompanies the formation of one mole of the compound in its standard state from its constituent elements in their standard states. • The standard enthalpy of formation for any element is ΔHof = 0.

  9. Standard Enthalpies of Formation for Selected Compounds

  10. Example Problem • H2S + 4 F2 2 HF + SF6ΔHorxn = ? ½ H2 + ½ F2  HF ΔHof= -273 kJ S + 3F2  SF6ΔHof = -1220 kJ H2S  H2 + S ΔHof = 21 kJ 2(½ H2+ ½ F2) + S + 3F2 + H2S  2 HF + SF6 + H2+ S ΔHof= -546 kJ + -1220 kJ + 21 kJ = -1745 kJ

  11. Standard Enthalpies of Formation Equation • ΔHorxn = ΣΔHof (products) - ΣΔHof(reactants) • CH4 + 2 O2 CO2 + 2 H2O ΔHof(CO2) = -394 kJ ΔHof(H2O) = -286 kJ ΔHof(CH4) = -75 kJ ΔHof(O2) = 0 kJ ΔHorxn= ( -394 kJ + 2 (-286) kJ) – (- 75 kJ + 2 (0 kJ)) ΔHorxn=-966 + 75 = -891 kJ

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