Chapter 6 Energy Transfer

1. Chapter 6Energy Transfer CHM 108 Suroviec Spring 2014

2. I. Principles of Heat Flow • Energy is the capacity to do work • Work is the result of a force acting through a distance

3. A. Different Types of Energy • Kinetic energy – energy associated with motion of an object • Thermal energy – energy associated with temperature of an object • Potential energy – energy associated with position of an object • Chemical energy – energy associated with position of an electron around an atom

4. B. Law of Conservation of Energy • Energy cannot be created or destroyed • Energy can be transferred from one kind to another

5. II. First Law of Thermodynamics • Total energy of the universe is constant and energy is neither created or destroyed • Internal Energy (IE) • Sum of kinetic and potential energies of all the particles that compose the system • Is a state function

6. A. Internal Energy • Internal energy change = DE

7. B. Exchanging Energy • A system can exchange energy with the surroundings through heat and work

8. B. Exchanging Energy

9. III. Quantifying Heat and Work • Heat • Heat = exchange of thermal energy between a system and its suroundings

10. B. Heat Capacity • When a system absorbs heat (q) its temperature changes by DT. • Heat capacity = quantity of heat required to raise T by 1oC. (J/oC) • Specific heat capacity = quantity of heat required to raise T of 1 gram of material by 1oC. (J/goC)

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12. Example • Suppose that you have the idea to start making stained glass. Given that glass turns to a pliable liquid around 1500oC and you are starting at 25oC, how much heat does a 2.50 g piece of glass absorb?

13. C. Pressure – Volume work • How does work change with the volume changes?

14. Example • Inflating a balloon requires P-V work 1. If I inflate a balloon from 0.200L to 0.985L at an external pressure of 1.00 atm, how much work in Latm is done? 2. Given that 101.3 J = 1 Latm, what is this value of work in Joules?

15. IV. Measuring ΔE • The way to measure this is with constant volume calorimetry

16. A. Calorimetry Calorimetry measures change in thermal energy between a system and its surroundings. At constant volume the change in temperature is related to heat absorbed by entire calorimeter

17. Example When 2.09g of sucrose (C12H22O11) is combusted the temperature rose from 23.89 oC to 28.45 oC. What is the DErxn for this in kJ and then in kJ/mole?

18. V. Enthalpy • In most cases we cannot hold the volume constant • We are interested with how much energy is given off at constant pressure

19. V. Enthalpy • The signs of DH and DE are similar in meaning

20. Example • If we burn 1 mole of fuel and constant pressure it produces 3452 kJ of heat and does 11 kJ or work. What are the values of DH and DE?

21. A. Stoichiometry involving DH • Enthalpy for chemical reactions = DHrxn For the reaction below, if you start with 13.2 kg of C3H8 what is q in kJ? C3H8 (l) + 5O2 (g)  3CO2 (g) + 4H2O (l) DHrxn = -2044 kJ

22. VI. Measuring DHrxn • Coffee cup calorimetry is a way to measure DHrxn. • Knowing the mass of solution, the heat evolved will cause a DT. Combing that with Cs of solution you can determine q.

23. Example • We want to know the DHrxn of the following reaction: Ca(s) + 2HCl (aq)  CaCl2 (aq) + H2 (g) Given 0.204 g of Ca(s) in 100.0 mL of HCl (aq), the temperature rose from 24.8 °C to 33.9 °C. The density of the solution is 1.00 g/mL and the Cs,soln is 4.18 J/g°C

24. 2H2O (s) 2H2O (l) H2O (l) H2O (g) H2O (s) H2O (l) H2O (l) H2O (s) H2O (s) H2O (l) DH = -6.01 kJ DH = 6.01 kJ DH = 2 x 6.01= 12.0 kJ DH = 6.01 kJ DH = 44.0 kJ IV. Thermochemical Equations • The stoichiometric coefficients always refer to the number of moles of a substance • If you reverse a reaction, the sign of DH changes • If you multiply both sides of the equation by a factor n, then DH must change by the same factor n. • The physical states of all reactants and products must be specified in thermochemical equations.

25. C(graphite) + O2(g) CO2(g)DH0 = -393.5 kJ rxn S(rhombic) + O2(g) SO2(g)DH0 = -296.1 kJ rxn CS2(l) + 3O2(g) CO2(g) + 2SO2(g)DH0 = -1072 kJ rxn C(graphite) + 2S(rhombic) CS2 (l) Calculate the standard enthalpy of formation of CS2 (l) given that: The final reaction for formation of CS2 is:

26. Standard enthalpy of formation (DH0) is the heat change that results when one mole of a compound is formed from its elements at a pressure of 1 atm. f DH0 (C, graphite) = 0 f V. Enthalpy of formation The standard enthalpy of formation of any element in its most stable form is zero. DHf0 (O2) = 0 DHf0 (O3) = 142 kJ/mol DH f0 (C, diamond) = 1.90 kJ/mol

27. VII. Standard Heats of Formation Example Write an equation for the formation of C6H12O6 from elements. The DHf0 for C6H12O6 is -1273.3 kJ/mole

28. B. Calculating DHf0 for a reaction • To calculate DH rxn0 subtract the heats of formations of the reactant multiplied by their stoichiometric coefficient from heats of formation of the products multiplied by their stoichiometric coefficents.

29. aA + bB cC + dD - [ + ] [ + ] = - S S = DH0 DH0 rxn rxn nDH0 (products) mDH0 (reactants) dDH0 (D) aDH0 (A) bDH0 (B) cDH0 (C) f f f f f f A. Hess’s Law The standard enthalpy of reaction (DH0rxn) is the enthalpy of a reaction carried out at 1 atm. Hess’s Law: When reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.

30. 2C6H6(l) + 15O2(g) 12CO2(g) + 6H2O (l) Example Benzene (C6H6) burns in air to produce carbon dioxide and liquid water. How much heat is released per mole of benzene combusted? The standard enthalpy of formation of benzene is 49.04 kJ/mol.