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Thermochemistry PowerPoint Presentation

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Published on : Apr 27, 2018 | Views : 1940 | Downloads : 1

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Slide 1-

Thermochemistry By: Ahmed M Elghamry
Slide 2-

Thermochemistry The Nature of Energy Enthalpy and Calorimetry Hess’s Law
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The Nature of Energy Energy is the capacity to do work or to produce heat.
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The Nature of Energy Energy is conserved. The law of conservation of energy states that energy can be converted from one form to another but can be neither created nor destroyed. The total energy is always constant. Heat is a one of the forms of energy. Heat and temperature are decidedly different. Temperature is a property that reflects the random motions of the particles in a particular substance. Heat, on the other hand, involves the transfer of energy between two objects due to a temperature difference.
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The Nature of Energy There are two ways to transfer energy: through work and through heat. Work is defined as force acting over a distance. This energy transfer is divided between work and heat depends on the specific conditions—the pathway. A state function refers to a property of the system that depends only on its present state. A change in this function (property) in going from one state to another state is independent of the particular pathway taken between the two states.
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This energy transfer also happens in chemical systems. We divide the universe into two parts: the system and the surroundings. The system is the part of the universe on which we wish to focus attention; the surroundings include everything else in the universe. In chemical reactions, we define the system as the reactants and products of the reaction. The surroundings consist of the reaction container (a furnace, for example), the room, and anything else other than the reactants and products. The Nature of Energy
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When a reaction results in the evolution of heat, it is said to be exothermic (exo- is a prefix meaning “out of”). Reactions that absorb energy from the surroundings are said to be endothermic. This energy generates from the difference in potential energy between reactants and products. The total energy is conserved. The energy gained by the surroundings must be equal to the energy lost by the system. The Nature of Energy
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In any exothermic reaction, some of the potential energy stored in the chemical bonds is being converted to thermal energy (random kinetic energy) via heat. the bonds in the products are stronger (on average) than those of the reactants. That is, more energy is released by forming the new bonds in the products than is consumed to break the bonds in the reactants. For an endothermic reaction, the situation is reversed, Energy that flows into the system as heat is used to increase the potential energy of the system. In this case the products have higher potential energy (weaker bonds on average) than the reactants. The total energy is conserved. The energy gained by the surroundings must be equal to the energy lost by the system. The Nature of Energy ∆PE represents the difference between the energy required to break the bonds in the reactants and the energy released when the bonds in the products are formed.
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The internal energy E of a system can be defined most precisely as (where ∆E represents the change in the system’s internal energy, q represents heat, and w represents work.) If a quantity of energy flows into the system via heat (an endothermic process), q is equal to +x, where the positive sign indicates that the system’s energy is increasing. On the other hand, when energy flows out of the system via heat (an exothermic process), q is equal to -x, where the negative sign indicates that the system’s energy is decreasing. The Nature of Energy
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Enthalpy and Calorimetry
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Enthalpy A less familiar property of a system is its enthalpy H, which is defined as (where E is the internal energy of the system, P is the pressure of the system, and V is the volume of the system.) ∆H = q only at constant pressure.
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Enthalpy For a chemical reaction, the enthalpy change is given by the equation At constant pressure, exothermic means ∆H is negative; endothermic means ∆H is positive.
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Calorimetry The device used experimentally to determine the heat associated with a chemical reaction is called a calorimeter. Calorimetry, the science of measuring heat, is based on observing the temperature change when a body absorbs or discharges energy as heat. One substance might require a great deal of heat energy to raise its temperature by one degree, whereas another will exhibit the same temperature change after absorbing relatively little heat. The heat capacity C of a substance, which is a measure of this property, is defined as
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Calorimetry Specific heat capacity: the energy Required to raise the temperature of One gram of a substance by one degree Celsius. Molar heat capacity: the energy Required to raise the temperature of One mole of a substance by one degree Celsius. Water his high heat capacity of 4.18 J/°C · g.
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Hess’s Law
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Hess’s Law In going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.
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Consume Energy well! And Good Bye!