Ch. 16 Reaction Energy

1. Ch. 16 Reaction Energy

2. Thermochemistry • __________________: the study of the transfers of energy as heat that accompany chemical reactions and physical changes. • ______________: an instrument to measure the energy absorbed or released as heat in a chemical or physical change. • ______________: a measure of the average kinetic energy of the particles in a sample of matter. ________________ • ____________: the SI unit of heat as well as other forms of energy. 16-2

3. Thermochemistry • _______ = N x m = kg x m2 s2 • _______: the energy transferred between samples of matter because of a difference in their temperatures. • _____________: the amount of energy required to raise the temperature of one gram of a substance by degree (C or K) _________________________________ heat = specific x mass x change in heat temp. 16-3

4. Example q = cp x m x ΔT Example: A 4.0g sample of glass was heated from 274K to 314 K, and was found to have absorbed 32 J of energy as heat. What is the specific heat of this glass, and how much energy would be gained with a temp. change of 314k to 344K? ________________________ cp = 0.20 J/gK ________________________) q = 24 J 16-4

5. Practice q = cp x m x ΔT 1) Determine the specific heat of a material if a 35 g sample absorbed 96 J as it was heated from 293 K to 313 K. 16-5

6. Enthalpy • ________________: the amount of energy absorbed by a system as heat during a process at constant pressure. _________________________ • ________________________: the quantity of energy transferred as heat during a chemical rxn. • ________________________: an equation that includes the quantity of energy released or absorbed as heat during the reaction. 16-6

7. Enthalpy • ________________________: energy is released during the rxn. • _____________________: energy is absorbed during the rxn. • The quantity of energy ______________ is proportional to the amount of reactants. 16-7

8. Exothermic Rxns C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O + (-2043KJ) 16-8

9. Exothermic Rxns C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O + (-2043KJ) In this reaction, energy is _______. ΔH is ___________becauseproducts have a ________ value for H than the reactants. ΔH is always negative for exothermic reactions 16-9

10. Endothermic Rxns C(s) + H2O(g) + 113KJ → CO(g) + H2(g) 16-10

11. Endothermic Rxns C(s) + H2O(g) + 113KJ → CO(g) + H2(g) • In this reaction, energy is __________.ΔHis _________ because the products have a __________ value for H than the reactants. ΔH is always positive for endothermic reaction. 16-11

12. Enthalpy • _____________________: the enthalpy change that occurs when one mole of a compound is formed from its elements in their standard state at STP. (standard temp and pressure, 0oC and 1 atm.) • ______ = standard enthalpy of a rxn. • ______ = standard enthalpy of formation. (elements in their standard state have ΔHof = 0, compounds with positive values are unstable) 16-12

13. Enthalpy • ______________________: the energy change that occurs during the complete combustion of one mole of a substance. • ___________: the overall enthalpy change in rxn is equal to the sum of enthalpy changes for the individual steps in the process. 16-13

14. Hess’s Law • Rules for applying Hess’s Law: 1) If you ________________________, you must multiply the ΔHby the same coefficient: CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) ΔH = -802 kJ 2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) ΔH = -1604 kJ 2) If an equation is ___________, the sign of ΔHis also ___________. 16-14

15. Hess’s Law Ex. CalculateΔHofor NO(g) + ½O2(g) → NO2(g) from th enthalpy data found in Appendix A-14. Solve by combining known eq. Rxn1) ½N2(g) + ½O2(g) → NO(g) ΔHof= 90.29 kJ Rxn2) ½N2(g) + O2(g) → NO2(g) ΔHof= 33.2 kJ NO(g) + ½O2(g) → NO2(g) ________________________. 16-15

16. Hess’s Law Rxn1) NO(g)→½N2(g) + ½O2(g) ΔHof= - 90.29 kJ (reversed) Rxn2) ½N2(g) + O2(g) → NO2(g) ΔHof= 33.2 kJ NO(g) + ½O2(g) → NO2(g) ΔHo= (-90.29 kJ) + (33.2 kJ) = -57.1 kJ 16-16

17. Practice 2) Calculate the enthalpy of rxn for the combustion of methane, CH4, to form CO2(g) and H2O(l). 16-17

18. Practice 16-18

19. Practice Example: Calculate the enthalpy of formation of pentane, C5H12 5C(s) + 6H2(g) → C5H12(g) ΔHof = ? Rnx1) C5H12(g) + 8O2(g)→ 5CO2(g) + 6H2O(l) ΔHoc = -3535.6 kJ Rxn2) C(s) + O2(g)→ CO2(g) ΔHof = -393.5 kJ Rxn3) H2(g) + ½O2(g)→ H2O(l)ΔHof = -285.8 kJ _____________________________________________________________________________________ 16-19

20. Practice 16-20

21. Practice 3) Calculate the ΔHofof butane, C4H10 16-21

22. Practice 16-22

23. Spontaneous Reactions • _____________ a measure of the degree of randomness of the particles, such as molecules, in a system. • ____________________ a combined enthalpy-entropy function. • _______________________ the difference between the change in enthalpy and the product of the kelvin temp. and the entropy change. ________________________ ΔG = - Spontaneous ΔG = + Not spontaneous ΔG = 0 Equilibrium 16-23

24. Example Example: For the rxn NH4Cl(s)→NH3(g) + HCl(g), ΔHo = 176 kJ/mol and ΔSo = 0.285 kJ/molK. Calculate ΔGo, is the rxn spontaneous at 298.15k? ΔGo = ΔHo – TΔSo = (176 kJ/mol) - (298.15k)(0.285 kJ/molK) = 91 kJ/mol 16-24

25. Practice 4) For the rxn Br2(l)→Br2g)ΔHo = 31 kJ/mol and ΔSo = 93 J/molK. At what temp. will this rxn be spontaneous? 16-25

26. Ch. 16The End!