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Chemistry 121 Chapter 5 Study material for this Chapter: Textbook sections 5.1 through 5.8 THERMOCHEMISTRY 5.1 The Nature of Energy Thermodynamics is the study of energy and its transformation Thermochemistry is the relationships between chemical reactions and energy changes Kinetic Energy and Potential Energy Energy of motion is kinetic energy E k = ½ mv 2 Potential energy is the energy that an object possesses by virtue of its position Potential energy can be converted to kinetic energy, e.g. an object dropping off a high building Units of Energy SI unit is joule , J From E k = ½ mv 2 ; 1 joule = kg x m 2 /s 2 Traditionally, the calorie was the unit of energy 1 cal = 4.184 J (exactly) The nutritional Calorie, Cal = 1000 cal Systems and Surroundings A system is part of the universe we are interested in studying Surroundings are the rest of the universe (the surroundings are the portions of the universe not involved in the system) If we are interested in the reaction between hydrogen and oxygen in a container, then the hydrogen and oxygen in the container form the system Transferring Energy: Work and Heat Some definitions o Force is a push or pull on an object 1
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o Work is the amount of force applied to an object over a distance W = F x d o Heat is energy transferred from a hotter object to a colder object o Energy is the capacity to do work or to transfer heat Homework Problems: 5.13, 5.15, 5.17 5.2 The First Law of Thermodynamics The first law of thermodynamics states that energy cannot be created or destroyed The first law of thermodynamics deals with the conservation of energy o The energy is constant of the system and surroundings o Any energy transferred from the system must be transferred to the surroundings and vice versa Internal Energy The total energy of a system is called the internal energy o Internal energy = potential + kinetic energy Absolute internal energy cannot be measured, only changes in the internal energy Change in internal energy: ∆E = E final - E initial A mixture of H 2 (g) and O 2 (g) has a higher internal energy than H 2 O (g) H 2 (g) + O 2 (g) → 2 H 2 O (g) ; ∆E ‹ 0 ; negative change in energy 2 H 2 O (g) → H 2 (g) + O 2 (g); ∆E › 0 ; positive change in energy Relating ∆E to Heat and Work From the first law of thermodynamics ∆E = q + w In other words, if a system undergoes physical or chemical change, the change in internal energy is given by the heat added to or absorbed by the system plus the work done on or by the system. Heat flowing from the surrounding to the system is positive, q › 0 Work done by the surroundings on the system is positive, w › 0 Endothermic and Exothermic Processes An endothermic process is one that absorbs energy from the surroundings An exothermic process is one that transfers heat to the surroundings State Functions A state function depends on the initial and final states of a system 2
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e.g. The altitude difference between Denver and Columbus does not depend on
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121lec5 - <?xml version="1.0"...

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