1. Chapter Ten Energy Changes in Chemical Reactions

2. Section 10.1 Energy & Energy Changes

3. Energy

4. System and Surroundings
System: Part we care about
Reactants & Products
Surroundings: Everything
else in the universe

5. Types of Systems Aopen system (mass and heat pass through)
Bclosed system (heat only pass through)
Cisolated system (no heat or mass transfer)

6. Energy Flow and Reactions
For chemical reactions to happen spontaneously, the final products must be more stable than the starting reactants
Higher energetic substances are typically less stable and more reactive
Lower energetic substances are typically more stable and less reactive

7. Exothermic and Endothermic
Thermal energy flows from warmer to cooler
H2O(s)  H2O(l)
2H2(g) + O2(g)  2H2O(l)

8. Section 10.2 Introduction to Thermodynamics

9. What is Thermodynamics
Study of heat and its transformations into other energies
Thermochemistry is a part of this
Thermodynamics studies changes in the state of a system

10. State Functions
State functions are properties that are determined by the state of the system, regardless of how it was achieved
Only concerned with change, Δ
Final – Initial
Ex:
Energy
Pressure
Volume
Temperature

11. 1st Law of Thermodynamics

12. Internal Energy (U) Systems have a certain amount of internal energy
Has 2 components:
Kinetic energy: various types of molecular and electron motion
Potential energy: attractive and repulsive interactions between atoms and molecules
ΔU = U(products) – U(reactants)

13. Potential and Kinetic

14. Calculating ΔU
ΔU = q + w
q = heat (absorbed or released by the system)
w = work (done on or by the system)

15. Calculating ΔU

16. Example
Calculate the overall change in internal energy (ΔU) for a system that absorbs 188 J of heat and does 141 J of work on its surroundings.

17. Group Quiz #1 Convert 723.01 J into calories
SKETCH and LABEL what an exothermic and endothermic energy vs. time graph would look like.
Calculate the overall change in internal energy for a system that releases 43 J in heat and has 37 J of work done on it by its surroundings

18. Section 10.3 Enthalpy

19. Energy and Enthalpy Reactions can be carried out in two ways:
In a closed container (constant volume):
qv = ΔU (ΔU = q + w)
In an open container (constant pressure):
qp = ΔH (ΔH = ΔU + PΔV)
work (w)
heat (q)

20. Enthalpy
Enthalpy is the heat content absorbed or released in a system under constant pressure
Combustion of propane gas: U

21. Enthalpy of Reaction (ΔH)
ΔH = H(products) – H(reactants)
“+” = endothermic
“—” = exothermic

22. Enthalpy and Exo and Endo

23. Thermochemical Equations
H2O(s)  H2O(l) ΔH = kJ/mol CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) ΔH = kJ/mol

24. Looking at the Numbers
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) ΔH = kJ/mol
How much energy is release from g of methane being burned?
If kJ of energy was released, how many grams of water was produced?

25. Properties of Enthalpy
If you change the AMOUNTS in a balanced equation, you change the enthalpy the same way
Ex: if coefficients are doubled, so is the enthalpy
2H2(g) + O2(g)  2H2O(l) ΔH = kJ
4H2(g) + 2O2(g)  4H2O(l) ΔH = kJ
If you reverse the equation, you reverse the sign of the ΔH
Ex: 2H2(g) + O2(g)  2H2O(l) ΔH = kJ
2H2O(l)  2H2(g) + O2(g) ΔH = kJ

26. Group Quiz #2
How much energy is associated with burning 75.3 g of sulfur dioxide according to the following equation:
What would be the energy associated with decomposing 1 mole of sulfur trioxide gas into sulfur dioxide gas and oxygen gas? (Hint: use properties of enthalpy!)

27. Section 10.4 Calorimetry

28. Calorimetry Measurement or heat changes within a system
Using a calorimeter

29. Looking at Heat and Temperature
vs.
vs.

30. Specific Heat vs. Heat Capacity
Specific Heat (s): amount of heat required to raise the temperature of 1 g of a substance by 1°C (ex: liquid water is J/(g·°C)
q = (s)(m)(ΔT)
q = heat (J); m=mass (g); ΔT=change in temp (oC)
Heat Capacity (C): amount of heat required to raise the temperature of an object by 1°C (J/oC)
q = (C)(ΔT)

31. Example Specific Heats

32. Example
What is the amount of heat (in kJ) required to heat 255 g of water from 25.2 °C to 90.5 °C?

33. Group Quiz #3
8540 J of energy is released when 927 g of granite is cooled. If the original temperature was 46.8 oC, what is the final temperature?

34. Coffee-Cup Calorimetry
Can calculate changes in heat using styrofoam cups
Assuming constant pressure
Therefore…
qp = (m)(s)(ΔT) = ΔH

35. Physical Example of Coffee-Cup Calorimetry
A 30.4-g piece of unknown metal is heated up in a hot bath to a temperature of 92.4°C. The metal is then placed in a calorimeter containing 100. g of water at 25.0°C. After the calorimeter is capped, the temperature of the calorimeter raises to 27.2°C. What was the specific heat of the unknown metal?

36. Group Quiz #4
125.0-g of a metal is heated to 100.0°C. It is then placed into a calorimeter containing mL (100.0 g) of water at 25.0°C and capped. The energy is transferred and the max temperature of 34.1°C is reached. What is the specific heat of the metal?

37. Coffee Cup Calorimetry with Chemicals
System: reactants and products (the reaction)
Surroundings: water in calorimeter (resulting solution)
For an exothermic reaction:
The system loses heat
The surroundings gain (absorb) heat

38. Chemical Example
Ex: mL of 1.00 M HCl and 50.0 mL of 1.00 M NaOH are mixed in a calorimeter and capped at room temp (25.0°C). The reaction reaches a max of 31.7°C. What is the ΔH°rxn?

39. Section 10.5 Hess’s Law

40. Hess’s Law

41. Hess’s Law H2O (l)  H2O(g) ∆H = +44.0 kJ/mol
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) ∆H = kJ/mol
H2O (l)  H2O(g) ∆H = kJ/mol

42. Examples

43. Group Quiz #5 Given the following, determine the ΔH for
3H2(g) + O3(g)  3H2O(g)

44. Section 10.6 Standard Enthalpies of Formation

45. Enthalpy of Formation
Standard Enthalpy of Formation (ΔH°f): heat change that results when 1 mole of a compound is formed from its constituent elements in their standard states
“Standard State” means “stable form” at:
1 atm and 25°C
Example: O(g) (249.4), O2(g) (0), O3(g) (142.2)

46. Enthalpy of Formation

47. Standard Enthalpy of Reaction
ΔH°rxn: enthalpy of a reaction under standard conditions

48. An Example
When we know reactions go to completion or can be done in one step, we can use a direct method
Ex: Calculate ΔH°rxn for
2SO(g) + 2/3O3(g)  2SO2(g)
From Appendix 2: SO(g): (5.01), O3(g): (142.2), SO2(g): (-296.4)

49. Group Quiz #6 Calculate the ΔHrxno for the following:
CH4(g) + 2 O2(g) -> CO2(g) + 2 H2O(l)
2 H2S(g) + 3 O2(g) -> 2 H2O(l) + 2 SO2(g) 

50. More on Enthalpy of Reaction
When a reaction is too slow or side reactions occur, enthalpy of reaction can be calculated using Hess’s Law

51. Section 10.7 Bond Enthalpy and the Stability of Covalent Molecules

52. Bond Enthalpy
Recall: when bonds are made, energy is given off (exo); when bonds break, energy is needed (endo)
Bond Enthalpy: the measure of stability of a molecule
Enthalpy change associated with breaking a particular bond in 1 mole of gaseous molecules
H2(g)  H(g) + H(g) ΔH = kJ/mol

53. Bond Enthalpy The higher the bond enthalpy, the stronger the bond
The bonds in different compounds have different bond enthalpies
Ex: O—H bond in water vs. O—H bond in methanol are different
Therefore, we speak of AVERAGE bond enthalpy

54. Bond Enthalpy

55. Bond Enthalpy

56. Bond Enthalpy

57. Example

58. Section 10.8 Lattice Energy and the Stability of Ionic Solids

59. Lattice Energy
Recall: amount of energy required to convert 1 mole of ionic solid to its constituent ions in the gas phase
Ex: NaCl(s)  Na+(g) + Cl-(g)