1. Thermodymanics

2.  Thermodynamics is a branch of science that focuses on energy changes that accompany chemical and physical changes.

3.  Objective: To calculate heat capacity.

4.  Heat (q): The energy transferred between objects that are at different temperatures.  Unit: joules (J)

5.  Molar Heat Capacity: Energy (heat) needed to increase the temperature of 1 mol of substance by 1 K. q = nC ∆ T q = heat n = # of moles C = molar heat capacity ∆ T = change in temperature

6.  The molar heat capacity of water is larger than the molar heat capacity of land. This means that water does not heat up as easily as land does. As a result, oceans can help keep coastal areas cool during the summer.  The filling of a fruit pie has a larger heat capacity than the crust. This means that fruit filling will retain heat better and the crust will cool much quicker. As a result, eating the fruit filling can cause burns (even though it may appear that the pie is cool).

7.  Determine the energy (heat) needed to increase the temperature of 10.0 mol of Hg by 7.5 K. The value of C for Hg is 27.8 J/K ۰ mol.

8. q = ? n = 10.0 mol C = 27.8 J/K ۰ mol ∆ T = 7.5 K q = nC ∆ T q = (10.0 mol)(27.8 J/K ۰ mol)(7.5 K) q= 2.1 x 10 3 J

9.  The molar heat capacity of tungsten is 24.2 J/K ۰ mol. Calculate the energy as heat needed to increase the temperature of 0.404 mol of W by 10.0 K.

10. q = ? n = 0.404 mol C = 24.2 J/K ۰ mol ∆ T = 10.0 K q = nC ∆ T q = (0.404 mol)(24.2 J/K ۰ mol)(10.0 K) q= 97.8J

11.  Suppose a sample of NaCl increased in temperature by 2.5 K when the sample absorbed 1.7 x 10 2 J energy (heat). Calculate the number of moles of NaCl if the molar heat capacity is 50.5 J/K ۰ mol.

12. q = 1.7 x 10 2 J n = ? C = 50.5 J/K ۰ mol ∆ T = 2.5 K q = nC ∆ T 1.7 x 10 2 J = n(50.5 J/K ۰ mol)(2.5 K) n= 1.3 mol

13.  Calculate the energy as heat needed to increase the temperature of 0.80 mol of nitrogen, N 2, by 9.5 K. The molar heat capacity of nitrogen is 29.1 J/K ۰ mol.

14. q = ? n = 0.80 mol C = 29.1 J/K ۰ mol ∆ T = 9.5 K q = nC ∆ T q = (0.80 mol)(29.1 J/K ۰ mol)(9.5 K) q= 2.2 x 10 2 J

15.  A 0.07 mol sample of octane, C 8 H 18, absorbed 3.5 x 10 3 J of energy. Calculate the temperature increase of octane if the molar heat capacity of octane is 254.0 J/K ۰ mol.

16. q = 3.5 x 10 3 J n = 0.07 mol C = 254.0 J/K ۰ mol ∆ T = ? q = nC ∆ T 3.5 x 10 3 J = (0.07 mol)(254.0 J/K ۰ mol) ∆ T ∆ T = 2.0 x 10 2 K

17.  Objective: To calculate the change in enthalpy.

18.  The total energy of a system is impossible to measure.  However, we can measure the change in enthalpy of a system.  Enthalpy: The total energy content of a sample.  H = Enthalpy  ∆ H = Change in Enthalpy

19.  ∆ H can be measured with a calorimeter.  A calorimeter is used to measure the heat absorbed or released in a chemical or physical change.

20.  Enthalpy changes can be used to determine if a process is endothermic or exothermic.  Exothermic Reaction: Negative Enthalpy Change  Endothermic Reaction: Positive Enthalpy Changes

21.  The standard enthalpy of formation ( ) is the enthalpy change in forming 1 mol of a substance from elements in their standard states.  Note: The value for the standard enthalpy of formation for an element is 0.  The values for the standard enthalpies of formation can be found using a table.

22. ΔH reaction = ΔH products - ΔH reactants  Step 1: Determine for each compound using enthalpy table.  Step 2: Multiply by the coefficients from the balanced equation (# of moles).  Step 3: Set up Δ H equation.  Step 4: calculate.

23. ΔH reaction = ΔH products - ΔH reactants Calculate ΔH for the following reaction and determine if the reaction is exothermic or endothermic. SO 2 (g) + NO 2 (g)  SO 3 (g) + NO(g)

24. ΔH reaction = ΔH products - ΔH reactants SO 2 (g) + NO 2 (g)  SO 3 (g) + NO(g) -296.8(1)33.1(1) -395.8(1) 90.3(1) kJ/molkJ/molkJ/molkJ/mol ΔH = [(-395.8 kJ/mol)(1) + (90.3 kJ/mol)(1)] – [(296.8 kJ/mol)(1) + (33.1 kJ/mol)(1)] Δ H = -305.5 kJ – 329.9 kJ Δ H = -635.4 kJ The reaction is exothermic.

25.  Calculate the enthalpy change for the following reaction: 2C 2 H 6 (g) + 7O 2 (g)  4 CO 2 (g) + 6H 2 O(g)

26.  Calculate Δ H for the following reaction: CaO(s) + H 2 O(l)  Ca(OH) 2 (s)

27.  Calculate the enthalpy change for the combustion of methane gas. CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l)

28.  Objective: To calculate the change in entropy.

29.  A reaction is more likely to occur if enthalpy ( Δ H) is negative.  However, some endothermic reactions can occur easily. Why? Entropy!  Entropy ( Δ S): A measure of the randomness or disorder of a system.  A process if more likely to occur if there is an increase in entropy (or if ΔS is positive).

30.  Entropy is increased by the following factors:  Diffusion (process of dispersion)  Dilution of a solution  Decreasing the pressure of a gas  Increasing temperature  The number of moles of product is greater than the number of moles of reactant  Increasing the total number of particles in a system  When a reaction produces more gas particles (opposed to liquid or solids)

31. Δ S reaction = Δ S products - Δ S reactants

32.  Find the change in entropy for the following reaction: 2Na (s) + 2HCl (g)  2NaCl (s) + H 2 (g)

33.  Find the change in entropy for the following reaction: 2Na (s) + 2H 2 O (l)  2NaOH (s) + H 2 (g)

34.  Objectives: (1) To calculate the change in Gibbs energy. (2) To determine if a reaction is spontaneous or nonspontaneous.

35.  Gibbs Energy: the energy in a system that is available to do useful work.  Gibbs Energy is also called Free Energy.  ΔG = Change in Gibbs Energy

36. ΔG reaction = ΔG products - ΔG reactants

37.  Calculate Δ G for the following water-gas reaction: C(s) + H 2 O (g)  CO (g) + H 2 (g)

38.  Calculate the Gibbs energy change that accompanies the following reaction: C(s) + O 2 (g)  CO 2 (g)

39.  Calculate the Gibbs energy change that accompanies the following reaction: CaCO 3 (s)  CaO (s) + CO 2 (g)

40.  If ΔG is negative, the forward reaction is spontaneous.  If ΔG is 0, the system is at equilibrium.  If ΔG is positive, the forward reaction is nonspontaneous.  NOTE: In this case, the reaction is spontaneous in the reverse direction.

41. ΔG = ΔH - TΔS  Step 1: Organize the information.  Step 2: Change the units.  Step 3: Step up ΔG equation.  Step 4: Calculate.

42.  Given the changes in enthalpy and entropy are -139 kJ and 277 J/K respectively for a reaction at 25 ⁰ C, calculate the change in Gibbs energy.

43. Δ H = -139 kJ Δ S = 277 J/K / (1000 J/kJ) = 0.277 kJ/K T = 25 ⁰ C + 273 = 298 K Δ G = ?

44.  Given the changes in enthalpy and entropy are -139 kJ and 277 J/K respectively for a reaction at 25 ⁰ C, calculate the change in Gibbs energy. Δ H = -139 kJ Δ S = 277 J/K / (1000 J/kJ) = 0.277 kJ/K T = 25 ⁰ C + 273 = 298 K Δ G = ? ΔG = ΔH - T ΔS ΔG = (-139 kJ) – (298K)(0.277 kJ/K) ΔG = (-139 kJ) – (82.546 kJ) ΔG = -221.55 kJ The reaction is spontaneous.

45.  A reaction has a ΔH of -76 kJ and a ΔS of -117 J/K. Calculate the ΔG at 298 K. Is the reaction spontaneous?

46.  A reaction has a ΔH of 11 kJ and a ΔS of 49 J/K. Calculate ΔG at 298 K. Is the reaction spontaneous?

47.  Objective: To calculate the melting and boiling point of a substance.

48.  T mp : melting point temperature  T bp : boiling point temperature  ΔH fus : molar enthalpy of fusion  ΔS fus : molar entropy of fusion  ΔH vap : molar enthalpy of vaporization  ΔS vap : molar entropy of vaporization

49.  Step 1: Organize the information.  Step 2: Change the units.  Step 3: Step up the equation.  Step 4: Calculate.

50.  The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar entropy of fusion is 9.79 J/mol · K. The enthalpy of vaporization at the boiling point if 59.2 kJ/mol, and the molar entropy of vaporization is 93.8 J/mol ·K. Calculate the melting point and boiling point of Hg.

51. T mp = ? ΔH fus = 2.295 kJ/mol ΔS fus = 9.79 J/mo l · K = 0.00979 kJ/ mo l · K

52.  The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar entropy of fusion is 9.79 J/mol · K. The enthalpy of vaporization at the boiling point if 59.2 kJ/mol, and the molar entropy of vaporization is 93.8 J/mol ·K. Calculate the melting point and boiling point of Hg. T mp = ? ΔH fus = 2.295 kJ/mol ΔS fus = 9.79 J/mo l · K = 0.00979 kJ/ mo l · K T mp = ( 2.295 kJ/mol) / ( 0.00979 kJ/ mo l · K ) T mp = 234 K

53.  The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar entropy of fusion is 9.79 J/mol · K. The enthalpy of vaporization at the boiling point if 59.2 kJ/mol, and the molar entropy of vaporization is 93.8 J/mol ·K. Calculate the melting point and boiling point of Hg. T mp = ? ΔH fus = 2.295 kJ/mol ΔS fus = 9.79 J/mo l · K = 0.00979 kJ/ mo l · K T mp = ( 2.295 kJ/mol) / ( 0.00979 kJ/ mo l · K ) T mp = 234 K T bp = ? ΔH vap = 59.2 kJ/mol ΔS vap = 93.8 J/mo l · K = 0.0938 kJ/ mo l · K

54.  The enthalpy of fusion of Hg is 2.295 kJ/mol, and the molar entropy of fusion is 9.79 J/mol · K. The enthalpy of vaporization at the boiling point if 59.2 kJ/mol, and the molar entropy of vaporization is 93.8 J/mol ·K. Calculate the melting point and boiling point of Hg. T mp = ? ΔH fus = 2.295 kJ/mol ΔS fus = 9.79 J/mo l · K = 0.00979 kJ/ mo l · K T mp = ( 2.295 kJ/mol) / ( 0.00979 kJ/ mo l · K ) T mp = 234 K T bp = ? ΔH vap = 59.2 kJ/mol ΔS vap = 93.8 J/mo l · K = 0.0938 kJ/ mo l · K T bp = ( 59.2 kJ/mol) / (0.0938 kJ/mo l · K ) T bp = 631 K

55.  For ethanol, the molar enthalpy of fusion is 4.931 kJ/mol and the molar entropy of fusion is 31.6 J/mol · K. The molar enthalpy of vaporization at the boiling point is 42.32 kJ/mol and the molar entropy of vaporization is 109.9 J/mol · K. Calculate the melting and boiling points for ethanol.

56.  For sulfur dioxide, the molar enthalpy of fusion is 8.62 kJ/mol and the molar entropy of fusion is 43.1 J/mol · K. The enthalpy of vaporization at the boiling point is 24.9 kJ/mol and the molar entropy of vaporization is 94.5 J/mol · K. Calculate the melting and boiling points for sulfur dioxide.

57.  For ammonia, ΔH fus is 5.66 kJ/mol and Δs fus is 29.0 J/mol · K. ΔH vap is 23.33 kJ/mol and ΔS vap is 97.2 J/mol · K. Calculate the melting and boiling points for ammonia.