1. Thermodynamics

2. Energy Energy- is the ability to do work or produce heat.
Potential Energy- energy due to the composition or position of an object.
Kinetic Energy- is energy of motion.

3. Thermodynamics
Law of Conservation of Energy- states that in any chemical reaction or physical process, energy can be converted from one form to another, but it is neither created nor destroyed.
Chemical Potential Energy- the energy stored in a substance because of its composition. (like food or gasoline)
Thermochemistry- is the study of heat change that accompany chemical reaction and phase changes.

4. Thermodynamics
Heat – (q) is energy that is in the process of flowing from a warmer object to a cooler object.
*Heat flows from ___hot___ to __cold__*
calorie- cal unit of heat. 1 calorie = J
Joule- J- the SI unit of heat and energy.

5. Thermodynamics
Temperature: A measure of the average kinetic energy of the atoms or molecules of substance
Units: Celsius, Kelvin and Fahrenheit
Thermal Energy: TOTAL kinetic energy of the particles of the substance
Units: Joules (J)

6. Conversions 1000J = 1 kJ 1 calorie = 4.184J
1 kcal= 1 Calorie = 1000 cal

7. Practice Problem 1
A fruit and oatmeal bar contains 142 Calories. Convert this energy to calories.

8. Practice Problem 2
If an endothermic process absorbs 256 J, how many kilocalories are absorbed?

9. Practice Problem 3
The breakfast I ate this morning contains 230 nutritional Calories. How much energy in joules will this breakfast supply?

10. Difference between Heat and Temperature
Temperature is the AVERAGE kinetic energy of all the particles (does NOT depend on the amount)
Units: Celsius, Fahrenheit, Kelvin
Heat (thermal) energy is the TOTAL kinetic energy of all of the particles (depends on the amount present)
Units: Joules, calories
Heat is the TRANSFER of thermal energy from hot to cold

11. Try it!
Compare the temperature and thermal (heat) energy of a small pot of boiling water vs. the entire Atlantic ocean!

12. Try it!
Compare the temperature and thermal (heat) energy of a small pot of boiling water vs. the entire Atlantic ocean!
The temperature would be higher in the pot of boiling water (higher average kinetic energy), but the thermal energy would be higher for the Atlantic Ocean (so many more particles to the total KE would be greater)

13. Enthalpy
Enthalpy of reaction is the amount of heat gained or lost in a chemical reaction.
Enthalpy = ΔH
ΔH = Heatproducts - Heatreactants
Δ means “change in”

14. Endothermic / Exothermic
Endothermic – heat/energy must be added for a reaction. (ΔH = +)
Exothermic – heat/energy are released from the reaction. (ΔH = - )

15. Endothermic vs. Exothermic

16. Practice!
Label the following reactions as exothermic or endothermic: H2 + ½ O2  H2O ΔH = kJ H2O  H2 + ½ O2 ΔH = 241.8kJ

17. Practice!
Rewrite these equations to include the ΔH value as a reactant or product: H2 + ½ O2  H2O ΔH = kJ H2O  H2 + ½ O2 ΔH = 241.8kJ

18. Try these!
Label as exothermic or endothermic and rewrite the equation to include the ΔH as a product or reactant.
4Fe + 3O2  2Fe2O3 ∆H = -1625kJ
NH4 + NO3-  NH4NO3 ∆H = +27 kJ

19. Answers!
Label as exothermic or endothermic and rewrite the equation to include the ΔH as a product or reactant. 4Fe + 3O2  2Fe2O3 ∆H = -1625kJ exothermic 4Fe + 3O2  2Fe2O kJ NH4 + NO3-  NH4NO3 ∆H = +27 kJ endothermic NH4 + NO kJ NH4NO3

20. Changes of State
Heat of vaporization (ΔHvap)  heat required to vaporize (boil) one gram of a liquid
Heat of fusion (ΔHfus)  heat required to melt one gram of a solid substance

21. Equations for change in state
q = mΔHf
q = mΔHv
q = heat (J)
m= mass (g)
ΔHf = heat of fusion (J/g); found on reference table
ΔHv = heat of vaporization (J/g); found on reference table
*These formulas are in your reference tables on the formula page!

22. During a phase change, temperature remains constant because kinetic energy remains constant as potential energy increases or decreases.

23. Practice Problem 1
Calculate the heat required to melt 25.7 g of solid water at its melting point.

24. Practice Problem 2
How much heat is evolved when 275 g of water vapor condenses to a liquid at its boiling point?

25. Specific Heat
amount of heat required to raise the temperature of one gram of a substance by one degree Celsius
You can find specific heat values in the reference table

26. Specific Heat-the formula!
q= mcΔt
Specific Heat J/g °C
Change in Temperature
(K or ° C)
Mass (g)
Amount of Heat (J)

27. Try it! Which has a higher specific heat? Water or sand
Metal pan or oven mitts

28. Try it! Which has a higher specific heat? Water or sand
Metal pan or oven mitts
Why? The more difficult it is to heat something up, the higher the specific heat because more heat is required to raise the temperature of that substance

29. Practice Problem 1
How much heat is lost when 4110 g of aluminum cools from 660°C to 25.0°C?

30. Practice Problem 2
How much heat is required to increase the temperature of 124 g of water from 17.5°C to 45.8°C?

31. Calorimeter qlost=-qgained
Insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process
qlost=-qgained
mc∆Tlost = -mc∆Tgained

32. Example
A 20.0 g piece of metal at a temperature of 90.0˚C is dropped into an insulated container holding 125 g of water at 20.0˚C. If the final temperature is 23.0˚C, what is the specific heat of the metal?

33. Label the parts of the heating curve above
Which numbers show an increase in kinetic energy? Potential energy?

34. Review! Exothermic or endothermic? Condensation Melting Freezing
Boiling
Sublimation (s  g)
ΔH = -13.6J
A + B J  C + D

35. Review! – Answers! Exothermic or endothermic? Condensation exothermic
Melting endothermic
Freezing exothermic
Boiling endothermic
Sublimation (s  g) endothermic
ΔH = -13.6J exothermic
A + B J  C + D endothermic

36. Quiz Topics Heat calculations (phase and temperature changes)
Energy conversions
Heat vocab
Exothermic vs. Endothermic
Heating Curves
Calorimeter (concepts only, NO calculations)

37. Entropy
Entropy – is the measure of disorder or randomness of the particles that make up a system.
Variable: S
Unit: Joules (J)
The Law of Disorder –states that spontaneous processes always proceed in such a way that the entropy of the universe increases. (The second law of Thermodynamics)

38. Entropy From each group, select the part with the highest entropy:
Increasing temperature or decreasing temperature of a substance
Reactants or products: 2NH3(g)  3H2(g) + N2(g)
ice or water or steam

39. Entropy The states of matter from lowest entropy to highest entropy:
(s)  (l)  (aq)  (g)
When S is positive that means that entropy is increasing
When S is negative that means that entropy is decreasing

40. Entropy Practice
Predict the sign of ∆Ssystem for each of the following changes.
CIF(g) + F2(g)  CIF3(g)
NH3(g)  NH3(aq)
CH3OH(l)  CH3OH(aq)
C10H8(l)  C10H8(s)

41. Kinetics and Equilibrium

42. Collision Theory
Reacting substances (atoms/ions/molecules) must collide.
Reacting substances must collide with sufficient energy to form activated complex. (Activation Energy = Ea)

43. Collision theory
3. Reacting substance must collide with the correct orientation.

44. Factors that affect reaction rates
Nature of reactants
More reactive the substance, the faster the reaction
Concentration
When increased, the rate increases (more collisions)
Surface area
More surface area  quicker reaction
Temperature
Usually increasing temperature  increases rate
Catalysts
Reaction takes a different path and lowers activation energy

45. Potential Energy Diagram

46. Endothermic vs. Exothermic

47. Endothermic vs. Exothermic
Endothermic because energy is added
Exothermic because energy is released

48. Catalyst
A substance that increases the rate of a chemical reaction without itself being consumed in the reaction.
An example of a catalyst is an enzyme

49. Equilibrium
Reversible Reaction is one that can occur in both forward and reverse directions.
Chemical equilibrium is a state in which the forward and reverse reactions balance each other because they take place at equal rates.
Rateforward reaction = Ratereverse reaction

50. Equilibrium
Equilibrium is DYNAMIC, meaning that Equilibrium is a state of action. The reaction is constantly shifting and moving to maintain equilibrium even though we can not observe this.

51. Equilibrium
Law of chemical equilibrium states that at a given temperature, a chemical system may reach a state in which a particular ratio of reactant to product concentrations has a constant value.
Equilibrium constant (Keq) is the numerical value of the ratio of products concentrations to reactant concentrations, with each concentration raised to the power of its coefficients in the balanced equation.

52. Generic reaction and formula
aA + bB ↔ cC + dD
Where A and B are the reactants, C and D are the products, and a, b, c, and d are the coefficients in a balanced equation.
Keq= [C]c[D]d
[A]a[B]b
Do not put SOLIDS and LIQUIDS in the keq equation!!!!!

53. Equilibrium Keq > 1 : More products than reactants at equilibrium.
Keq < 1 : More reactants than products at equilibrium.

54. Equilibrium Find the Equilibrium Expression for the following:
N2O4(g) ↔ 2NO2(g)
CO(g) + 3H2(g) ↔ CH4(g) + H2O(g)
Ca(OH)2(s) + H2O(l)  Ca2+(aq) + 2OH-(aq)
CaCO3(s) ↔ CaO(s) + CO2(g)
2H2(g) + O2(g)  2H2O (l)

55. Le Chatelier

56. Le Chatelier’s Principle:
A principle stating that if a stress is applied to a system in equilibrium, the equilibrium will shift to alleviate the effect of the stress.

57. A “stress” can be: Change in concentration of product or reactant
Change in temperature
Change in pressure (just gases)

58. Changing Concentration
Equilibrium will shift AWAY from an increase in concentration and TOWARD a decrease in concentration.

59. This reaction is at equilibrium
heat
NH3 N2 H2
What happens to the reaction if you increase the concentration of Nitrogen?

60. Increase nitrogen concentration
heat
NH3 N2 H2
How does the reaction need to shift to make it go back to equilibrium?
What will happen to all of the substances as a result?
H2 will decrease, NH3 will increase
Right

61. How have things changed?
heat
NH3 N2 H2
How have things changed?
heat
NH3 H2 N2

62. Example 2 2SO3(g) ↔ 2SO2(g) + O2(g) Describe what happens when:
You decrease SO2
You increase O2

64. Changing Temperature If endothermic, heat is added to the product side
If exothermic, heat is added to the reactant side
Keq (equilbrium constant): is only affected by a temperature change. Concentration and pressure do NOT change Keq
Treat heat the same way you do concentration changes. If temperature increases, heat increases!

65. heat
NH3 N2 H2
What happens when you increase the temperature?
heat
NH3 N2 H2
How does the reaction need to shift to make it go back to equilibrium?
What will happen to all of the substances as a result?
N2 and H2 increase, NH3 decreases
Left

66. How have things changed?
heat
NH3 N2 H2
How have things changed?
heat
NH3 N2 H2

67. Try it 2SO3(g) + heat ↔ 2SO2(g) + O2(g) Increase the temperature
Decrease the temperature

68. Changing Pressure
Only affects equilibrium reactions with gaseous substances
Think of your reaction in a closed container.
When you increase pressure (decrease volume) you want to have less gas in your system
When you decrease pressure (increase volume) you want to produce more gas in your system

69. Changing Pressure Rule Thumb
When you INCREASE pressure: always shift toward fewer moles of gas!
When you DECREASE pressure: always shift toward more moles of gas!

70. Decrease
Pressure
Normal
Increase
Pressure

71. Try It! Ex 1: N2(g) + 3H2(g)  2NH3(g) Increase the pressure
Decrease the pressure
Ex 2: 2SO3(g) ↔ 2SO2(g) + O2(g)

72. Stress Shift [CH4] [H2S] [CS2] [H2] Keq
CH4(g) + 2H2S(g) ↔ CS2(g) + 4H2(g) H = kJ/mol
Stress
Shift
[CH4]
[H2S]
[CS2]
[H2]
Keq
Increase CS2
Decrease H2
Increase Temp
Increase Pressure

73. Stress Shift [CH4] [H2S] [CS2] [H2] Keq
CH4(g) + 2H2S(g) ↔ CS2(g) + 4H2(g) H = kJ/mol
Stress
Shift
[CH4]
[H2S]
[CS2]
[H2]
Keq
Increase CS2
Left 
Increase 
Increase
-----
Decrease
No Change
Decrease H2
 Right
Increase Temp
Right 
Changes!
Increase Pressure
 Left

74. REVIEW QUESTIONS!

75. Practice!
How much heat is required to heat up 22g of gold from 15˚C to 79˚C?

76. Answer!
How much heat is required to heat up 22g of gold from 15˚C to 79˚C?
𝑞=𝑚𝐶∆𝑇 𝑞= −15 =𝟏𝟖𝟏.𝟔𝑱

77. Practice!
How much heat is released to freeze 179 g of water?

78. Answer! How much heat is released to freeze 179 g of water?
𝑞= 𝑚∆𝐻 𝑓𝑢𝑠 =179 −334 =−𝟓𝟗𝟕𝟖𝟔𝑱

79. Practice!
How much heat is released to decrease the temperature of 13.6g water by 13.2˚C?

80. Answer!
How much heat is released to decrease the temperature of 13.6g water by 13.2˚C?
𝑞=𝑚𝐶∆𝑇= −13.2 =−𝟕𝟓𝟎.𝟒𝑱

81. Practice!
How much heat is required to boil 58 g of water?

82. Answer! How much heat is required to boil 58 g of water?
𝑞=𝑚 ∆𝐻 𝑣𝑎𝑝 = =𝟏𝟑𝟏𝟎𝟖𝟎𝐉

83. Practice! Exothermic or Endothermic? Boiling Freezing ΔH = - 25.6 kJ
C2H5OH + 3O2  2CO2 + 3H2O kJ
2NaHCO kJ  Na2CO3 + H2O +CO2

84. Answer! Exothermic or Endothermic? Boiling Endothermic
Freezing Exothermic
ΔH = kJ Exothermic
ΔH = kJ Endothermic
C2H5OH + 3O2  2CO2 + 3H2O kJ Exothermic
2NaHCO kJ  Na2CO3 + H2O +CO2 Endothermic

85. Practice! Write the equilibrium expression for this reaction:
Ca(OH)2(s)  Ca2+(aq) + 2OH-(aq)

86. Answer! Write the equilibrium expression for this reaction:
Ca(OH)2(s)  Ca2+(aq) + 2OH-(aq)
Keq = [Ca2+][OH-]2

87. Practice! Write the equilibrium expression for this reaction:
2SO2(g) + O2(g)  2SO3(g)

88. Answer! Write the equilibrium expression for this reaction:
2SO2(g) + O2(g)  2SO3(g)
𝑲𝒆𝒒= [ 𝑺𝑶 𝟑 ] 𝟐 [ 𝑺𝑶 𝟐 ] 𝟐 𝑶 𝟐

89. Practice!
What happens when SO2 increases?
2SO2(g) + O2(g)  2SO3(g)

90. Answer! What happens when SO2 increases? 2SO2(g) + O2(g)  2SO3(g)
Shift is to the right
O2 decreases
SO3 increases

91. Practice!
What happens when SO3 increases?
2SO2(g) + O2(g)  2SO3(g)

92. Answer! What happens when SO3 increases? 2SO2(g) + O2(g)  2SO3(g)
Shift is to the left
SO2 and O2 increase

93. Practice! What happens when pressure increases?
2SO2(g) + O2(g)  2SO3(g)

94. Answer! What happens when pressure increases?
2SO2(g) + O2(g)  2SO3(g)
Shift is to the right
SO2 and O2 decrease
SO3 increase

95. Problem!
Draw an exothermic potential energy diagram and label the energy of the products, energy of the reactants, energy of the activated complex, activation energy (of the forward and reverse reactions), and enthalpy.

96. Answer!
Draw an exothermic potential energy diagram and label the energy of the products, energy of the reactants, energy of the activated complex, activation energy (of the forward and reverse reactions), and enthalpy.
1: Energy of the reactants
2: Energy of the activated complex
3: Activation energy of the forward
4: Enthalpy
5: Activation energy of the reverse
6: Energy of the products 5 6