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Exp6 - Experiment 6 THERMOCHEMISTRY I Learning Objectives...

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6-1 Experiment 6 THERMOCHEMISTRY I. Learning Objectives… To introduce the concept of thermochemistry. To illustrate the additivity of heats of reaction ( Hess’s Law ). To define and investigate exothermic and endothermic reactions. II. Background Information Thermochemistry is the study of the heat energy involved in chemical reactions and changes of physical state . Heat energy ( thermal energy) is always spontaneously transferred from hotter to colder matter. The First Law of Thermodynamics (Law of Energy Conservation) states that the total energy of the universe must remain constant. Therefore, all energy transferred between a system and its surroundings must be accounted for as heat or work . The standard S.I. unit for heat energy is the joule ( J ). It takes 4.184 joules (1 calorie) to raise the temperature of one gram of water by 1 ° C. The kilojoule (kJ) is commonly used (1000 joule = 1 kilojoule) in many applications. Exothermic and Endothermic Reactions When a chemical reaction takes place at roughly constant temperature and pressure (bench-top conditions), the system defined by the reactants and products either absorbs or releases heat energy. If the reacting system releases heat energy to its surroundings, a concurrent increase in surroundings temperature is observed, and the reaction is exothermic . If the system absorbs heat energy from its surroundings, a decrease in the surroundings temperature is observed, and the reaction is endothermic. A measure of the amount of heat given off or absorbed in any chemical reaction is called the enthalpy change or heat of reaction, and is given the symbol Δ H . The relationship of enthalpies of
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6-2 reactants and products to one another ( e.g . enthalpy change) in both endothermic and exothermic reactions is shown below in Figure 6.1. When thermodynamic measurements are carried out at standard-state conditions (constant pressure of 1 atm and constant temperature of 298 K), the reaction enthalpy is designated as the standard enthalpy change or Δ H o . It is important to have standardized values because the enthalpy of a reaction can vary with different reaction conditions. The following reaction for the formation of water from its constituents is exothermic: H 2 (g) + ½ O 2 (g) H 2 O (l) Δ H ° f = -286 kJ For every mole of H 2 O (l) formed at standard-state conditions, 286 kilojoules of heat energy are released . When the standard enthalpy change of reaction describes the formation of 1 mol of compound directly from its elements in their standard states (as in this example), the value of Δ H o is called the standard heat of formation and is given the symbol Δ H ° f . To determine the enthalpy change for a given reaction ( Δ H ° rxn ), the summation of the heats of formation ( Δ H ° f ) for the reactants are subtracted from the summation of the heats of formation ( Δ H ° f ) for the products.
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6-3 Δ H ° rxn = [n Δ H ° f (products)] - [n Δ H ° f (reactants)] Tables containing the standard heats of formation for a number of compounds are available in the appendices of any general chemistry textbook.
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