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lecture 03 2010

# lecture 03 2010 - 12 In a laboratory calorimeter combustion...

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12 In a laboratory calorimeter, combustion is often carried out. The Δ T is measured, and the heat evolved when M grams of substance undergoes combustion is calculated from the heat capacity of the calorimeter and its contents. Note that the units are J per degree. No moles involved in this calculation. Example : Suppose 1.00 grams of a substance with a formula weight of 100 g /mole undergoes combustion in a calorimeter. The heat capacity of the calorimeter and its contents is 19,000 J per degree. A Δ T of 2.20 ° C is observed. What is the Δ E for combustion? q V = 19,000 J/deg × 2.20 ° C = 41,800 J, so this is a change in internal energy. This heat is per 1g/100 g/mol = 0.01 moles of compound So Δ E = -q V /moles = -4180 kJ/mole. Why is the sign for the reaction negative? Positive q is absorbed by the surroundings. Therefore negative q is absorbed by the system.

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13 Properties of Enthalpy The enthalpy is an especially useful quantity for calorimetry, since most chemical reactions occur at constant pressure, 1 atm. The units of Δ H are J mol -1 . So, for the reaction CO( g ) + ½ O 2 ( g ) CO 2 ( g ) Δ H = -283.0 kJ. 283.0 kJ of energy are liberated when one mole of CO 2 is formed. So, 2 × 283.0 = 566.0 kJ is liberated when 2 moles of CO 2 are formed. H is an extensive thermodynamic variable. H = E + PV. E, P, V are thermodynamic functions of state (they are not interactions between the system and the surroundings) so H is also a state function. Δ H is independent of path. This is a very important concept . We can make CO
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lecture 03 2010 - 12 In a laboratory calorimeter combustion...

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