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Thermochemistry

Thermochemistry. Thermodynamics. Study of energy transformations Thermochemistry is a branch of thermodynamics which describes energy relationships in chemical reactions. Energy. Capacity to do work or to transfer heat

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Thermochemistry

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  1. Thermochemistry

  2. Thermodynamics • Study of energy transformations • Thermochemistry is a branch of thermodynamics which describes energy relationships in chemical reactions

  3. Energy • Capacity to do work or to transfer heat • Mechanical work (w) is the product of force (F) operating on an object and the distance (d) through which it moves • W = F x d • Energy is required to do work

  4. Heat (Q) • Heat is the energy transferred from one object to another due to a difference in temperature

  5. Forms of Energy • Kinetic Energy – energy of motion - magnitude depends on the mass of the object and its velocity - EK = ½ m v2 - both mass and speed determine how work it can do

  6. Potential Energy – stored energy • Other forms of energy are simply types of kinetic or potential on an atomic or molecular level

  7. Energy Units • Joule (J) • 1J = 1 kg m2/ s2 • A calorie (cal) is the amount of energy required to raise the temp of 1 g of water 1 ºC 1 cal = 4.184 J

  8. Example • A 145 g baseball is thrown with a speed of 25 m/s. Calculate the kinetic energy in Joules. • What is the kinetic energy in calories?

  9. Systems • Portion we single out for study • Surroundings is everything else outside the system • When studying energy changes in a chemical reaction, the reactants and products are the system and everything else is the surroundings

  10. Law of Conservation of Energy • Energy can be converted from one form to another but cannot be created or destroyed • Also called “First Law of Thermodynamics”

  11. Internal Energy • Total energy of a system – sum of kinetic and potential energies • Cannot determine exact internal energy • Can only determine a change in internal energy ΔE = Efinal – Einitial

  12. If ΔE is positive there is a gain in internal energy in the system • If ΔE is negative the system lost energy to its surroundings • Higher energy systems tend to lose energy and are therefore less stable

  13. Heat and Work • Any system can exchange energy with surroundings in two ways – as heat or work • Internal energy increases as heat is added to or work is done on a system ΔE = Q + w Q is positive if heat is added to system w is positive if work is done on the system

  14. Heat Changes • Exothermic Reactions – when heat is given off by the reaction (-Q) • Endothermic Reactions – when heat is used by the reaction (+Q)

  15. Example • As a combustion reaction occurs the system loses 550 J of heat to its surroundings and it does 240 J of work in moving a piston. What is the change in its internal energy?

  16. State Function • These are systems for whom the value of ΔE does not depend on the previous history of the sample, only on the present condition • Energy is a state function • Work and heat are not state functions

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