1. Chapter 8 Thermochemistry

2. Thermodynamics  Study of the changes in energy and transfers of energy that accompany chemical and physical processes.  address 3 fundamental questions  Will two (or more) substances react when they are mixed under specified conditions?  If they do react, what energy changes and transfers are associated with their reaction?  If a reaction occurs, to what extent does it occur?

3. The First Law of Thermodynamics  Exothermic reactions combustion of propane combustion of n-butane

4. The First Law of Thermodynamics  Exothermic reactions release specific amounts of heat as products  Potential energies of products are lower than potential energies of reactants.

5. The First Law of Thermodynamics  2 basic ideas of importance  systems tend toward a state of minimum potential energy

6. The First Law of Thermodynamics  2 basic ideas of importance  systems tend toward a state of maximum disorder

7. The First Law of Thermodynamics  Also known as Law of Conservation of Energy  The total amount of energy in the universe is constant.  Energy can be converted from one form to another but cannot be created.

8. Some Thermodynamic Terms  System - substances involved in the chemical and physical changes under investigation  for us this is what is happening inside the beaker  Surroundings - rest of the universe  outside the beaker  Universe - system plus surroundings

9. Some Thermodynamic Terms  Thermodynamic State of a System - set of conditions that describe and define the system  number of moles of each substance  physical states of each substance  temperature  pressure  State Functions - properties of a system that depend only on the state of the system  capital letters

10. Some Thermodynamic Terms  state functions are independent of pathway  climbing a mountain, taking two different paths  E 1 = energy at bottom of mountain  E 1 = mgh 1  E 2 = energy at top of mountain  E 2 = mgh 2   E 2 - E 1 = mgh 2 - mgh 1 = mg(  h)

11. Some Thermodynamic Terms  Properties that depend only on values of state functions are also state functions  examples: T P V

12. Enthalpy Change,  H  Commonly, chemistry is done at constant pressure  open beakers on a desk top are at atmospheric pressure   H - enthalpy change change in heat content at constant pressure  H = q p   H rxn - heat of reaction  H rxn = H products - H reactants  H rxn = H substances produced - H substances consumed

13. Calorimetry  coffee-cup calorimeter - used to measure the amount of heat produced (or absorbed) in a reaction at constant P  measures q P

14. Calorimetry  exothermic reaction - heat evolved by reaction is determined from the temperature rise of the solution  2 part calculation  Amount of heat gained by calorimeter is the heat capacity of the calorimeter or calorimeter constant  value determined by adding a specific amount of heat to calorimeter and measuring T rise

15. Calorimetry  When 3.425 kJ of heat is added to a calorimeter containing 50.00 g of water the temperature rises from 24.00 0 C to 36.54 0 C. Calculate the heat capacity of the calorimeter in J/ 0 C. The specific heat of water is 4.184 J/g 0 C.  Four part calculation

16. Calorimetry  Find the temperature change  Find the heat absorbed by the water in going from 24.00 0 C to 36.54 0 C.

17. Calorimetry  Find the heat absorbed by the calorimeter. total amount of heat added to calorimeter - heat absorbed by water  Find the heat capacity of the calorimeter (heat absorbed by the calorimeter)/(temperature change)

18. Calorimetry  A coffee-cup calorimeter is used to determine the heat of reaction for the acid-base neutralization When we add 25.00 mL of 0.500 M NaOH at 23.000 0 C to 25.00 mL of 0.600 M CH 3 COOH already in the calorimeter at the same temperature, the resulting temperature is observed to be 25.947 0 C.

19. Calorimetry The heat capacity of the calorimeter has previously been determined to be 27.8 J/ 0 C. Assume that the specific heat of the mixture is the same as that of water, 4.18 J/g 0 C and that the density of the mixture is 1.02 g/mL.  Two part calculation: a Calculate the amount of heat given off in the reaction.

20. Calorimetry

22.  Determine  H for the reaction under the conditions of the experiment. must determine the number of moles of reactants consumed limiting reactant calculation

23. Calorimetry

24.  finally, calculate  H based on the limiting reactant calculation

25. Thermochemical Equations  Thermochemical equations are a balanced chemical reaction plus the  H value for the reaction.  for example:  coefficients in thermochemical equations must be interpreted as numbers of moles  1 mol of C 5 H 12 reacts with 8 mol of O 2 to produce 5 mol of CO 2, 6 mol of H 2 O, and releasing 3523 kJ is referred to as one mole of reactions

26. Thermochemical Equations  Equivalent method of writing thermochemical equations   H < 0 designates an exothermic reaction   H > 0 designates an endothermic reaction

27. Thermochemical Equations  Write the thermochemical equation for the previous reaction.

28. Standard States & Standard Enthalpy Changes  Thermochemical standard state conditions T = 298.15 K P = 1.0000 atm  Thermochemical standard states pure substances in their liquid or solid phase - standard state is the pure liquid or solid gases - standard state is the gas at 1.00 atm of pressure gaseous mixtures - partial pressure must be 1.00 atm aqueous solutions - 1.00 M concentration

29. Standard Molar Enthalpies of Formation,  H f o  Standard molar enthalpy of formation  symbol is  H f o  defined as the enthalpy for the reaction in which one mole of a substance is formed from its constituent elements  for example:

30. Standard Molar Enthalpies of Formation,  H f o  Standard molar enthalpies of formation have been determined for many substances and are tabulated in Table 8.3 and Appendix 1 in the text.  Standard molar enthalpies of elements in their most stable forms at 298.15 K and 1.000 atm are zero.  Example: The standard molar enthalpy of formation for phosphoric acid is -1281 kJ/mol. Write the equation for the reaction for which  H o rxn = -1281 kJ. P in standard state is P 4 phosphoric acid in standard state is H 3 PO 4(s)

31. Standard Molar Enthalpies of Formation,  H f o

32. Hess’s Law  Hess’s Law of Heat Summation - enthalpy change for a reaction is the same whether it occurs by one step or by any (hypothetical) series of steps  true because  H is a state function  we know the following  H o ’s

33. Hess’s Law  We could calculate the  H o for [1] by properly using the  H o ’s for [2] and [3]

34. Hess’s Law  Example : Given the following equations and  H o  values calculate  H o for the reaction below.

35. Hess’s Law  Use a little algebra and Hess’s Law to get the appropriate  H o  values

36. Hess’s Law  The + sign of the  H o  value tells us that the reaction is endothermic.  The reverse reaction is exothermic, i.e.,

37. Hess’s Law  Hess’s Law in a more useful form any chemical reaction at standard conditions, the standard enthalpy change is the sum of the standard molar enthalpies of formation of the products (each multiplied by its coefficient in the balanced chemical equation) minus the corresponding sum for the reactants

38. Hess’s Law  Example: Calculate  H o 298  for the following reaction from data in Appendix 1 

39. Hess’s Law  Example: Calculate  H o 298  for the following reaction from data in Appendix 1 

40. Hess’s Law  Application of Hess’s Law and more algebra allows us to calculate the  H f o  for a substance participating in a reaction for which we know  H rxn o, if we also know  H f o  for all other substances in the reaction.

41. Hess’s Law  Example: Given the following information, calculate  H f o for H 2 S (g)

42. Hess’s Law

43. Bond Energies  Bond energy - amount of energy required to break the bond and separate the atoms in the gas phase

44. Bond Energies  Table of average bond energies  MoleculeBond Energy (kJ/mol) F 2 159 Cl 2 243 Br 2 192 O 2 (double bond)498 N 2 (triple bond)946

45. Bond Energies  Bond energies can be calculated from other  H o 298  values

46. Bond Energies  Example: Calculate the bond energy for hydrogen fluoride, HF.

47. Bond Energies  Example: Calculate the bond energy for hydrogen fluoride, HF.

48. Bond Energies  Example: Calculate the average N-H bond energy in ammonia, NH 3.

49. Bond Energies

50.  In gas phase reactions  H o values may be related to bond energies of all species in the reaction.

51. Bond Energies  Example: Use the bond energies listed in Table 8.4 to estimate the heat of reaction at 25 o C for the reaction below.

52. Bond Energies  Example: Use the bond energies listed in Table 8.4 to estimate the heat of reaction at 25 o C for the reaction below.