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Lecture12A - Enthalpy and Chemical Reactions(19.10-12...

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Lecture 12 1 Enthalpy and Chemical Reactions (19.10-12) Most (but not all) chemical reactions are conducted in open vessels, under constant atmospheric conditions. Changes in enthalpies can be measured for these reactions. (sometimes called thermochemistry, at least in the “old” days, i.e. 20 th century) products reactants - Enthalpy change for reaction r H H H - r r When 0, exothermic - excess enthalpy is released as heat 0, endothermic - enthalpy is absorbed by product as heat H H < Fig 19.8 Since enthalpy is a state function, the path that we take for a given overall reaction doesn’t matter. As you learned in Gen. Chem. we can use Hess’s Law to compute enthalpy of reactions from individual reaction steps.
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Lecture 12 2 You have seen examples like this many times before (I hope): S(s) + O 2 (g) SO 2 (g) Δ r H=-296.83 kJ (1) SO 2 (g) + ½ O 2 (g) SO 3 (g) Δ r H= -98.89 kJ (2) If we consider (1) and (2) as separate steps, the net reaction: S(s) + O 2 (g) SO 3 (g) Δ r H= Δ r H(1)+Δ r H(2) =-395.72 kJ 3 2 Of course, reversing a reaction is taking the reverse path: S(s) + O 2 (g) SO 2 (g) Δ r H=-296.83 kJ (1) SO 2 (g) S(s) + O 2 (g) Δ r H=? (2) Since the net reaction is: S(s) + O 2 (g) S(s) + O 2 (g) Δ r H=0 Δ r H(1)+Δ r H(2) =0 or - Δ r H(1)= Δ r H(2) This should be (very) old stuff for you.
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Lecture 12 3 This applies to any type of chemical reaction (e.g. combustion), transformation (e.g. sublimation, fusion, vaporization) or process (e.g.
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