Thermochemistry

1. 0 Thermochemistry Chapter 6 Dr. Victor H. Vilchiz

2. Thermochemistry 0 • Thermodynamics is the science that studies heat and its transformations to other forms of energy. • Thermochemistry is the study of the quantity of heat absorbed or evolved in the case of chemical reactions.

3. Energy 0 • Energy is defined as the capacity to do work. • In every situation in which energy is transferred there are two components. • System – The component in which we are interested (it can be an object or a chemical reaction). • Surroundings- Whatever surrounds our system.

4. Energy 0 • There are three broad concepts of energy: • Kinetic Energy (Ek) is the energy associated with an object by virtue of its motion. • Potential Energy (Ep) is the energy an object has by virtue of its position in a field of force. • Internal Energy (U) is the sum of the kinetic and potential energies of the particles making up a substance. We will look at each of these in detail.

5. Energy 0 • Kinetic Energy: An object of mass m and speed or velocity v has kinetic energy Ekequal to • This shows that the kinetic energy of an object depends on both its mass and its speed.

6. A Problem to Consider 0 • Consider the kinetic energy of a person whose mass is 130 lb (59.0 kg) traveling in a car at 60 mph (26.8 m/s). • The SI unit of energy, kg.m2/s2, is given the name Joule.

7. Energy 0 • Potential Energy: This energy depends on the “position” (such as height) in a “field of force” (such as gravity). • For example, water of a given mass m at the top of a dam is at a relatively high “position” h in the “gravitational field” g of the earth.

8. A Problem to Consider 0 • Consider the potential energy of 1000 lb of water (453.6 kg) at the top of a 300 foot dam (91.44 m).

9. Energy 0 • Internal Energy is the energy of the particles making up a substance and it contains kinetic and potential energy which is not included in Ek and Ep • The total energy of a system is the sum of its kinetic energy, potential energy, and internal energy.

10. Energy 0 • The Law of Conservation of Energy:Energy may be converted from one form to another, but the total quantities of energy remain constant. Also known as the First Law of Thermodynamics. DEUniv=0 this can be rewritten as follows: DEUniv=Efinal-Einitial=0 In the case of chemistry DEUniv=Eproducts-Ereactants=0

11. Energy Transfer 0 • In chemical reactions, energy is often transferred from the “system” to its “surroundings,” or vice versa. EUniv=Esystem+Esurroundings EUniv=Esystem+Esurroundings but from conservation of energy… 0=Esystem+Esurroundings this means that Esystem=-Esurroundings In other words: any energy/heat lost by the system is gained by the surroundings or vice-versa.

12. Heat of Reaction 0 • Heatis denoted by the symbol q. • The sign of q is positive if heat is absorbed by the system. • The sign of q is negative if heat is evolved by the system. • Heatof Reaction is the value of q required to return a system to the given temperature at the completion of the reaction.

13. Heat of Reaction 0 • Anexothermic processis a chemical reaction or physical change in which heat is evolved (q is negative). • Anendothermic processis a chemical reaction or physical change in which heat is absorbed (q is positive).

14. Heat of Reaction Exothermicity “out of” a system Dq < 0 Endothermicity “into” a system Dq > 0 0 Surroundings Surroundings Energy Energy System System

15. Enthalpy and Enthalpy Change 0 • The heat absorbed or evolved by a reaction depends on the conditions under which it occurs. • Usually, a reaction takes place in an open vessel, and therefore at the constant pressure of the atmosphere. • The heat of this type of reaction is denoted qp, the heat at constant pressure.

16. Heat and Work 0 • Energy can be separated into two parts a)thermal energy and b) work E=q+w • The values of q and w depend on the energy flow. • In chemistry, we define the sign of energy transfer from the point of view of the system. (Engineering does it from the point of view of the surroundings and hence the signs are switched.)

17. Changes in Energy hot water Energy (heat) lost to surroundings. q<0 DE<0 E RT water RT water Energy (heat) gained from surroundings. q>0 DE>0 E ice water 0 • There are four cases in energy change. • Heat flowing out of the system with no work. • Heat gained by the system with no work.

18. Changes in Energy 0 • Work done by system no heat. • Work done on the system no heat. HCl(aq) and Zn initial state Energy lost as the system moves the piston up. w<0 DE<0 E ZnCl2 and H2 gas final state

19. Energy Changes and Energy Units 0 • In Chem 101 we will deal only with systems in which we have energy changes in the form of heat only. Therefore E=q since w=0 • Units of Energy • SI Unit (Joule) • 1J = 1kg●m2/s2 • 1 cal = 4.184J • 1Calorie =1 Cal = 1000 cal • 1 BTU= 1055 J

20. Enthalpy and State Functions DV 0 • A state function, a property of a system that depends only on its present and initial states and is independent of the path followed. If we have the piston from earlier where a reaction forces a change in Volume but the Pressure is kept constant then the system performs some work. Work= -PDV work done on surroundings If the reaction takes place at constant pressure then the thermodynamic part is described as enthalpy (H). DH=DE+PDV; DE=qp+PDV DH=qp Change in enthalpy = the change in heat at constant p

21. Enthalpy 0 • Enthalpy, denotedH, is an extensive property of a substance that can be used to obtain the heat absorbed or evolved in a chemical reaction. • Anextensive propertyis one that depends on the quantity of substance.

22. Enthalpy 0 • There are three cases when DE=DH • The chemical reactions do not involve gases • The reaction does involve gases but there is no net change in the moles of gases between reactants and products. (no work is done) • There is work done but q>>PDV

23. Enthalpy Change 0 • Since H depends on E, P, and V and these are state functions then H must also be a state function. • The change in enthalpy for a reaction at a given temperature and pressure (called the enthalpy of reaction) is obtained by subtracting the enthalpy of the reactants from the enthalpy of the products.

24. Thermochemical Equations 0 • Athermochemical equationis the chemical equation for a reaction (including phase labels) in which the equation is given a molar interpretation, and the enthalpy of reaction for these molar amounts is written directly after the equation. If DH is (-) the reaction is exothermic If DH is (+) the reaction is endothermic

25. Thermochemical Equations 0 • In athermochemical equationit is important to note phase labels because the enthalpy change, DH, depends on the phase of the substances.

26. Thermochemical Equations 0 • The following are two important rules for manipulating thermochemical equations: • When a thermochemical equation is multiplied by any factor, the value of DH for the new equation is obtained by multiplying the DH in the original equation by that same factor. • When a chemical equation is reversed, the value of DH is reversed in sign.