1. Energy Thermodynamics (rev. 0910)
Chapter 5
Energy
Thermodynamics
(rev. 0910)

2. Definition
Thermodynamics-
is the study of energy transformations.

3. Chemical Reactions
Chemical reactions involve not just the conversion of reactants into products, but also involve an energy change in the form of heat—heat released as the result of a reaction, or heat absorbed as a reaction proceeds.
Energy changes accompany all chemical reactions and are due to rearranging of chemical bonding.

4. Making Bonds
Addition of energy is always a requirement for the breaking of bonds but the breaking of bonds in and of itself does not release energy.
Energy release occurs when new bonds are formed.

5. Bond Energy
If more energy is released when new bonds form than was required to break existing bonds, then the difference will result in an overall release of energy.
If, on the other hand, more energy is required to break existing bonds than is released when new bonds form, the
difference will result overall in energy being absorbed.

6. Overall Reaction Energy
Whether or not an overall reaction releases or requires energy depends upon the final balance between the breaking and forming of chemical bonds.

7. Energy is... the ability to do work or produce heat. conserved.
made of heat and work.
a state function. (Energy is a property that is determined by specifying the condition or “state” (e.g., temperature, pressure, etc.) of a system or substance.)
independent of the path, or how you get from point A to B.

8. Energy
While the total internal energy of a system (E) cannot be determined, changes in internal energy (E) can be determined.
The change in internal energy will be the amount of energy exchanged between a system and its surroundings during a physical or chemical change.
Δ E = E final - E initial

9. Definitions Work is a force acting over a distance.
Heat is energy transferred between objects because of temperature difference. (Heat is not a property of a system or substance and is not a state function. Heat is a process—the transfer of energy from a warm to a cold object.)

10. System vs Surroundings
The Universe is divided into two halves.
system and the surroundings.
In a chemistry setting, a system includes all substances undergoing a physical or chemical change.
The surroundings would include everything else that is not part of the system.

11. Heat
Most commonly, energy is exchanged between a system and its surroundings in the form of heat.
Heat will be transferred between objects at different temperatures.
Thermochemistry is the study of thermal energy changes.

12. Exo vs Endo Exothermic reactions release energy to the surroundings.
Endothermic reactions absorb energy from the surroundings.

13. Heat
Potential energy

14. Heat
Potential energy

15. Three Parts Every energy measurement has three parts:
A unit ( Joules of calories).
A number.
a sign to tell direction.
negative - exothermic
positive- endothermic

16. Surroundings
System
Energy
DE <0

17. Surroundings
System
Energy
DE >0

18. Same rules for heat and work
Heat given off is “negative”.
Heat absorbed is “positive”.
Work done by the system on the surroundings is negative.
Work done on the system by the surroundings is positive.

19. First Law of Thermodynamics
The energy of the universe is constant.
It is also called the:
Law of conservation of energy.
q = heat w = work
In a chemical system, the energy exchanged between a system and its surroundings can be
accounted for by heat (q) and work (w).
DE = q + w
Take the system’s point of view to decide signs.

20. Conservation of Energy
Energy exchanged between a system and its surroundings can be considered to off set one another.
The same amount of energy leaving a system will enter the surroundings (or vice versa), so the total amount of energy remains constant.

21. Metric Units
The SI (Metric System) unit for all forms of energy is the joule (J).

22. Heat and Work DE = q + w - q is exothermic -q = -∆H +q is endothermic
-w is done “by” the system
+w is done “on” the system
Note:
∆H stands for enthalpy which is the heat of reaction

23. Practice Problem
A gas absorbs 28.5 J of heat and then performs 15.2 J of work. The change in internal
energy of the gas is:
(a) 13.3 J
(b) J
(c) 43.7 J
(d) J
(e) none of the above

24. Answer
(b) E = q + w
28.5 J J = J

25. Practice Problem
Which of the following statements correctly describes the signs of q and w for the following exothermic process at 1 atmosphere pressure and 370 Kelvin?
H2O(g) → H2O(l)
(a) q and w are both negative
(b) q is positive and w is negative
(c) q is negative and w is positive
(d) q and w are both positive
(e) q and w are both zero

26. Answer
(c). An exothermic indicates q is negative and the gas is condensing to a liquid so it is exerting less pressure on its surroundings indicating w is positive.

27. What is work? Work is a force acting over a distance. w= F x Dd
P = F/ area
d = V/area
w= (P x area) x D (V/area)= PDV
Work can be calculated by multiplying pressure by the change in volume at constant pressure.
Use units of liter•atm or L•atm

28. Pressure and Volume Work
Work refers to a force that moves an object over a distance.
Only pressure/volume work (i.e., the expansion/contraction of a gas) is of significance in chemical systems and only when there is an increase or decrease in the amount of gas present.

29. Work needs a sign
If the volume of a gas increases, the system has done work on the surroundings.
work is negative
w = - PDV
Expanding work is negative.
Contracting, surroundings do work on the system W is positive.
1 L•atm = J

30. Example
When, in a chemical reaction, there are more moles of product gas compared to
reactant gas, the system can be thought of as performing work on its surroundings (making w < 0) because it is “pushing back,” or moving back the atmosphere to make room for the expanding gas.
When the reverse is true, w > 0.

31. Compressing and Expanding Gases
Compressing gas
Work on the system is positive
Work is going into the system
Expanding gas
Work on the surroundings is negative
Work is leaving the system

32. Clarification Info
If the reaction is performed in a rigid container, there may be a change in pressure,
but if there is no change in volume, the atmosphere outside the container didn’t “move” and without movement, no work is done by or on the system.
If there is no change in volume (V= 0), then no work is done by or on the system (w= 0) and the change in internal energy will be entirely be due to the heat involved ( ΔE = q).

33. Examples
What amount of work is done when 15 L of gas is expanded to 25 L at 2.4 atm pressure?
If J of heat are absorbed by the gas above. what is the change in energy?
How much heat would it take to change the gas without changing the internal energy of the gas?

34. Enthalpy The symbol for Enthalpy is H
H = E + PV (that’s the definition)
at constant pressure.
DH = DE + PDV
the heat at constant pressure qp can be calculated from
DE = qp + w = qp - PDV
qp = DE + P DV = DH

35. DH = DE + PDV Using DH = DE + PDV
the heat at constant pressure qp can be calculated from:
DE = qp + w (if w = - PDV then…)
DE = qp – PDV (now rearrange)
qp = DE + P DV = DH

36. Examples of Enthapy Changes
KOH(s) → K(aq) + OH-1 (aq) ΔHsolution = kJ mol1
C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l) ΔHcombustion = -2221kJ/mol
H2O(s) → H2O(l) ΔHfusion = 6.0 kJ/mol
Fe2O3(s) + 2Al(s) → Al2O3(s) + 2Fe(s) ΔHreaction kJ/mol
Ca(s) + O2(g) H2(g) → Ca(OH)2(s) ΔHformation kJ/mol1

37. 3 Methods
There are a variety of methods for calculating overall enthalpy changes that you should be familiar with.
The three most common are the:
the use of Heats of Formation
Hess’s Law
the use of Bond Energies

38. Heat of Reaction
To compare heats of reaction for different reactions, it is necessary to know the temperatures at which heats of reaction are measured and the physical states of the reactants and products.
Look in the Appendix of the textbook to find Standard Enthapy tables.

39. Standard Enthalpy of Formation
Measurements have been made and tables constructed of Standard Enthalpies of Formation with reactants in their “standard states”.
Use the symbol DHºf
Standard state is the most stable physical state of reactants at:
1 atmosphere pressure
specified temperature—usually 25 °C
1 M solutions
For solids which exist in more than one allotropic form, a specific allotrope must be specified.

40. ΔH°formation
It is important to recognize that the ΔH°formation (abbreviated as ΔH°f) is really just the heat of reaction for a chemical change involving the formation of a compound from its elements in
their standard states.

41. Standard Enthalpies of Formation
The standard heat of formation is the amount of heat needed to form 1 mole of a compound from its elements in their standard states.
See the table in the Appendix
Remember: For an element the value is 0

42. Equation Practice
You need to be able to write the equation correctly before solving the problem.
Try…
What is the equation for the formation of NO2 ? Try writing the equation.

43. Practice Answer ½N2 (g) + O2 (g) ® NO2 (g)
You must make one mole to meet the definition.

44. Since we can manipulate the equations
We can use heats of formation to figure out the heat of reaction.
Lets do it with this equation.
C2H5OH +3O2(g) ® 2CO2 + 3H2O
which leads us to this rule.

45. Hess’s Law Definition:
When a reaction may be expressed as the algebraic sum of other reactions, the enthalpy change of the reaction is the algebraic sum of the enthalpy changes for the combined reactions.

46. Hess’ Law Enthalpy is a state function. It is independent of the path.
We can add equations to come up with the desired final product, and add the DH values.
Two rules to remember:
If the reaction is reversed the sign of DH is changed
If the reaction is multiplied, so is DH

47. Enthalpy
As enthalpy is an extensive property, the magnitude of an enthalpy change for a chemical reaction depends upon the quantity of material that reacts.
This means:
if the amount of reacting material in an exothermic reaction is doubled, twice the quantity of heat energy will be released.

48. For the oxidation of sulfur dioxide gas
SO2(g) + ½O2(g) → SO3(g)
ΔH° = - 99 kJ/mol
Doubling the reaction results in:
2SO2(g) + O2(g) → 2SO3(g)
ΔH° = kJ/mol
Notice that if you double the reaction, you must double the ΔH° value.

49. Sign Change of ΔH 2SO2(g) + O2(g) → 2SO3(g) ΔH° = - 198 kJ/mol
If the reaction is written as an endothermic reaction:
2SO3(g) → 2SO2(g) + O2(g)
ΔH° = kJ/mol

50. Tips for Hess’s Law Problems
It is always a good idea to begin by looking for species that appear as reactants and products in the overall reaction.
This will provide a clue as to whether a reaction needs to be reversed or not.
Second, consider the coefficients of species that appear in the overall reaction.
This will help determine whether a reaction needs to be multiplied before the overall summation.

51. Example C(s) + O2(g) → CO2(g) ΔH°f = - 394 kJ/mol
2H2(g) + O2(g) → 2H2O(l) ΔH°f = kJ/mol
CO2(g) + 2H2O(l) → CH4(g) + 2O2(g) ΔH°f = kJ/mol
C(s) + 2H2(g) → CH4(g) ΔH°f = - 75 kJ/mol

52. Hess’s Law Example Given DHº= +77.9kJ DHº= +495 kJ DHº= +435.9kJ
Calculate DHº for this reaction

53. Example Given calculate DHº for this reaction DHº= -1300. kJ

54. Problems to Try
Try 12-3 Practice Problems

55. O2
NO2
-112 kJ
H (kJ)
180 kJ
NO2
68 kJ N2
2O2

56. Calorimetry
Calorimetry is the study of the heat released or absorbed during physical and chemical reactions.
For a certain object, the amount of heat energy lost or gained is proportional to
the temperature change. The initial temperature and the final temperature in the calorimeter
are measured and the temperature difference is used to calculate the heat of reaction.

57. Equipment: Calorimeter
There are two kinds of calorimeters:
constant pressure
bomb

58. Calorimetry Constant pressure calorimeter
(called a coffee cup calorimeter)
A coffee cup calorimeter measures DH.
The calorimeter can be an insulated cup, full of water.

59. Definitions:
Heat capacity is the amount of energy required to raise the temperature of an object 1 kelvin or 1 °C.
Specific heat capacity is the heat capacity of
1 gram of a substance.
Molar heat capacity is the heat capacity of
1 mole of a substance.

60. Heat Capacity
Heat capacity is an extensive property, meaning it depends on the amount present—a large amount of a substance would require more heat to raise the temperature 1 K than a small amount of the same substance.

61. Specific Heat Capacity
Definition:
The specific heat capacity of each substance is an intensive property which relates the heat capacity to the mass of the substance.

62. Specific Heat Capacity “c”
q = mass x c x DT
Or written as
q = mcDT
This is the main equation for calorimetry calculations
mass will be in grams
Units for “c” are J/g K or J/g °C
The specific heat of water is 1 cal/g ºC

63. Molar Heat Capacity
Molar Heat Capacity is the heat capacity of 1 mole of a substance.
molar heat capacity = c/moles
heat = molar heat x moles x DT
Remember that heat is shown as ∆H
Make the units work and you’ve done the problem right.

64. Sign of q
• If a process results in the sample losing heat energy, the loss in heat is designated as q is negative.
The temperature of the surroundings will increase during this exothermic process.
• If the sample gains heat during the process, then q is positive. The temperature of the
surroundings will decrease during an endothermic process.
The amount of heat that an object gains or loses is directly proportional to the change in
temperature.

65. Examples The specific heat of graphite is 0.71 J/gºC.
Calculate the energy needed to raise the temperature of 75 kg of graphite from 294 K to 348 K.

66. Extra Problem
A 46.2 g sample of copper is heated to 95.4ºC and then placed in a calorimeter containing 75.0 g of water at 19.6ºC. The final temperature of both the water and the copper is 21.8ºC.
What is the specific heat of copper?

67. Calorimetry Constant volume calorimeter is called a bomb calorimeter.
Material is put in a container with pure oxygen. Wires are used to start the combustion. The container is put into a container of water.
The heat capacity of the calorimeter is known and tested.
Since DV = 0, PDV = 0, DE = q

68. Bomb Calorimeter thermometer stirrer full of water ignition wire
Steel bomb
sample

69. Properties
intensive properties are not related to the amount of substance.
Examples: density, specific heat, temperature.
Extensive property - does depend on the amount of substance.
Examples: Heat capacity, mass, heat from a reaction.

70. Collegeboard. ( ). Professional Development workshop materials: Special focus thermochemistry.