1. Energy Thermodynamics
Chapter 6
Energy
Thermodynamics

2. Energy is... The ability to do work. Conserved. made of heat and work.
a state function.
independent of the path, or how you get from point A to B.
Work is a force acting over a distance.
Heat is energy transferred between objects because of temperature difference.

3. Chapter 6: Thermochemistry
Energy
Literally means “work within,” however no object contains work
Energy refers to the capacity to do work
– that is, to move or displace matter
EOS
2 basic types of energy:
– Potential (possibility of doing work because of composition or position)
– Kinetic (moving objects doing work)
Chapter 6: Thermochemistry

4. Chapter 6: Thermochemistry
Energy
Potential Energy – in a gravitational field
(= position)
Kinetic Energy – energy of motion
PE = mgh
m = mass (kg)
g = gravity constant (m s-2)
h = height (m)
KE = 1/2mv2
m = mass (kg)
v = velocity (m s–1)
v2 = (m2 s–2)
EOS
units are kg m2 s–2  J
units are kg m2 s–2 = J
Chapter 6: Thermochemistry

5. Potential Energy
Potential energy is energy an object possesses by virtue of its position or chemical composition.
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6. Kinetic Energy
Kinetic energy is energy an object possesses by virtue of its motion. 1 2
KE =  mv2
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7. Conversion of Energy Energy can be converted from one type to another.
For example, the cyclist above has potential energy as she sits on top of the hill.
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8. Conversion of Energy
As she coasts down the hill, her potential energy is converted to kinetic energy.
At the bottom, all the potential energy she had at the top of the hill is now kinetic energy.
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9. Units of Energy The SI unit of energy is the joule (J). kg m2
An older, non-SI unit is still in widespread use: the calorie (cal).
1 cal = J
1 J = 1 
kg m2 s2
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10. The universe is divided into two halves.
the system and the surroundings.
The system is the part you are concerned with.
The surroundings are the rest.
Exothermic reactions release energy to the surroundings.
Endo thermic reactions absorb energy from the surroundings.

11. Chapter 6: Thermochemistry
Thermochemistry is the study of energy changes that occur during chemical reactions
Universe
Focus is on heat and matter transfer between the system ...
System
Surroundings
and the surroundings
EOS
Chapter 6: Thermochemistry

12. Definitions: System and Surroundings
The system includes the molecules we want to study (here, the hydrogen and oxygen molecules).
The surroundings are everything else (here, the cylinder and piston).
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13. Chapter 6: Thermochemistry
Types of systems one can study:
OPEN
Matter
Energy
CLOSED
Matter
Energy
EOS
ISOLATED
Matter
Energy
Chapter 6: Thermochemistry

14. Definitions: Work
Energy used to move an object over some distance is work.
w = F  d
where w is work, F is the force, and d is the distance over which the force is exerted.
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15. Chapter 6: Thermochemistry
Internal Energy
Internal Energy (U) is the total energy contained within the system, partly as kinetic energy and partly as potential energy
EOS
Kinetic involves three types of molecular motion ...
Chapter 6: Thermochemistry

16. Chapter 6: Thermochemistry
Internal Energy
Internal Energy (U) is the total energy contained within the system, partly as kinetic energy and partly as potential energy
Potential energy involves intramolecular interactions ...
EOS
and intermolecular interactions ...
Chapter 6: Thermochemistry

17. Heat Energy can also be transferred as heat.
Heat flows from warmer objects to cooler objects.
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18. Chapter 6: Thermochemistry
Heat (q)
Heat is energy transfer resulting from thermal differences between the system and surroundings
“flows” spontaneously from higher T  lower T
EOS
“flow” ceases at thermal equilibrium
Chapter 6: Thermochemistry

19. Heat
Potential energy

20. Heat
Potential energy

21. Same rules for heat and work
Heat given off is negative.
Heat absorbed is positive.
Work done by system on surroundings is positive.
Work done on system by surroundings is negative.
Thermodynamics- The study of energy and the changes it undergoes.

22. Chapter 6: Thermochemistry
Work (w)
Work is an energy transfer between a system and its surroundings
Recall from gas laws … the product PV = energy
EOS
Pressure–volume work is the work of compression (or expansion) of a gas
Chapter 6: Thermochemistry

23. Chapter 6: Thermochemistry
Calculating Work (w)
PV work is calculated as follows:
w = –PDV
Sign conventions: think from the perspective of the system
SYSTEM
WORK
EOS
If work is done by the system, the system loses energy equal to –w
Chapter 6: Thermochemistry

24. Chapter 6: Thermochemistry
Calculating Work (w)
SYSTEM
WORK
Expansion is an example of work done by the system—the weight above the gas is lifted
EOS
compression (or expansion) of a gas
ExpansionWork
Chapter 6: Thermochemistry

25. 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.
units of liter - atm L-atm

26. 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

27. 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?

28. Surroundings
System
Energy
DE <0

29. Surroundings
System
Energy
DE >0

30. Direction Every energy measurement has three parts.
A unit ( Joules of calories).
A number how many.
and a sign to tell direction.
negative - exothermic
positive- endothermic

31. First Law of Thermodynamics
Energy is neither created nor destroyed.
In other words, the total energy of the universe is a constant; if the system loses energy, it must be gained by the surroundings, and vice versa.
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32. First Law of Thermodynamics
The energy of the universe is constant.
Law of conservation of energy.
q = heat
w = work
DE = q + w
Take the systems point of view to decide signs.

33. Internal Energy
The internal energy of a system is the sum of all kinetic and potential energies of all components of the system; we call it E.
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34. Internal Energy
By definition, the change in internal energy, E, is the final energy of the system minus the initial energy of the system:
E = Efinal − Einitial
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35. Changes in Internal Energy
When energy is exchanged between the system and the surroundings, it is exchanged as either heat (q) or work (w).
That is, E = q + w.
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36. E, q, w, and Their Signs
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37. Exchange of Heat between System and Surroundings
When heat is absorbed by the system from the surroundings, the process is endothermic.
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38. Exchange of Heat between System and Surroundings
When heat is absorbed by the system from the surroundings, the process is endothermic.
When heat is released by the system into the surroundings, the process is exothermic.
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39. Chapter 6: Thermochemistry
States of a System
The state of a system refers to its exact condition, determined by the kinds and amounts of matter present, the structure of this matter at the molecular level, and the prevailing pressure and temperature
Example: internal energy (U) is a function of the state of the system ...
EOS
Chapter 6: Thermochemistry

40. Chapter 6: Thermochemistry
State Functions
A state function is a property that has a unique value that depends only on the present state of a system and not on how the state was reached, nor on the history of the system
EOS
DU = Uf – Ui
Chapter 6: Thermochemistry

41. State Functions
However, we do know that the internal energy of a system is independent of the path by which the system achieved that state.
In the system below, the water could have reached room temperature from either direction.
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42. State Functions Therefore, internal energy is a state function.
It depends only on the present state of the system, not on the path by which the system arrived at that state.
And so, E depends only on Einitial and Efinal.
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43. State Functions However, q and w are not state functions.
Whether the battery is shorted out or is discharged by running the fan, its E is the same.
But q and w are different in the two cases.
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44. Enthalpy Symbol is H Change in enthalpy is DH delta H
If heat is released the heat content of the products is lower
DH is negative (exothermic)
If heat is absorbed the heat content of the products is higher
DH is positive (endothermic)

45. Enthalpy
If a process takes place at constant pressure (as the majority of processes we study do) and the only work done is this pressure-volume work, we can account for heat flow during the process by measuring the enthalpy of the system.
Enthalpy is the internal energy plus the product of pressure and volume:
H = E + PV
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46. Enthalpy abbreviated 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

47. Enthalpy
Since E = q + w and w = -PV, we can substitute these into the enthalpy expression:
H = E + PV
H = (q+w) − w
H = q
So, at constant pressure, the change in enthalpy is the heat gained or lost.
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48. Endothermicity and Exothermicity
A process is endothermic when H is positive.
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49. Endothermicity and Exothermicity
A process is endothermic when H is positive.
A process is exothermic when H is negative.
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50. Enthalpy of Reaction
The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the reactants:
H = Hproducts − Hreactants
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51. Enthalpy of Reaction
This quantity, H, is called the enthalpy of reaction, or the heat of reaction.
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52. The Truth about Enthalpy
Enthalpy is an extensive property.
H for a reaction in the forward direction is equal in size, but opposite in sign, to H for the reverse reaction.
H for a reaction depends on the state of the products and the state of the reactants.
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53. Chapter 6: Thermochemistry
Calorimetry
Calorimetry is a technique used to measure heat exchange in chemical reactions
A calorimeter is the device used to make heat measurements
EOS
Calorimetry is based on the law of conservation of energy
Chapter 6: Thermochemistry

54. Calorimetry Relationships
The heat capacity (C) of a system is the quantity of heat required to change the temperature of the system by 1 oC
calculated from C = q/DT units of J oC–1 or J K–1
Specific heat is the heat capacity of a one-gram sample
EOS
Specific heat = C/m = q/mDT
units of J g–1 oC–1 or J g–1 K–1
Chapter 6: Thermochemistry

55. Specific Heats
Molar heat capacity is the product of specific heat times the molar mass of a substance
units are J mol–1 K–1
A useful form of the specific heat equation is:
q = m CDT
If DT > 0, then q > 0 and heat is gained by the system
EOS
If DT < 0, then q < 0 and heat is lost by the system
Chapter 6: Thermochemistry

56. Calorimetry Measuring heat. Use a calorimeter. Two kinds
Constant pressure calorimeter (called a coffee cup calorimeter)
heat capacity for a material, C is calculated
C= heat absorbed/ DT = DH/ DT
specific heat capacity = C/mass

57. Calorimetry molar heat capacity = C/moles
heat = specific heat x m x DT
heat = molar heat x moles x DT
Make the units work and you’ve done the problem right.
A coffee cup calorimeter measures DH.
An insulated cup, full of water.
The specific heat of water is 1 cal/gºC
Heat of reaction= DH = sh x mass x DT

58. Heat Capacity and Specific Heat
The amount of energy required to raise the temperature of a substance by 1 K (1C) is its heat capacity.
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59. Heat Capacity and Specific Heat
We define specific heat capacity (or simply specific heat) as the amount of energy required to raise the temperature of 1 g of a substance by 1 K.
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60. Heat Capacity and Specific Heat
Specific heat, then, is
Specific heat =
heat transferred
mass  temperature change
s = q
m  T
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61. 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.
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?

62. 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

63. Constant Pressure Calorimetry
By carrying out a reaction in aqueous solution in a simple calorimeter such as this one, one can indirectly measure the heat change for the system by measuring the heat change for the water in the calorimeter.
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64. Constant Pressure Calorimetry
Because the specific heat for water is well known (4.184 J/g-K), we can measure H for the reaction with this equation:
q = m  s  T
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65. Bomb Calorimeter thermometer stirrer full of water ignition wire
Steel bomb
sample

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

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

68. Hess’s Law
Hess’s law states that “[i]f a reaction is carried out in a series of steps, H for the overall reaction will be equal to the sum of the enthalpy changes for the individual steps.”
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69. Hess’s Law
Because H is a state function, the total enthalpy change depends only on the initial state of the reactants and the final state of the products.
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70. O2
NO2
-112 kJ
H (kJ)
180 kJ
NO2
68 kJ N2
2O2

71. Chapter 6: Thermochemistry
Standard Enthalpies
The standard enthalpy of reaction (DHo) is the enthalpy change for a reaction in which the reactants in their standard states yield products in their standard states
The standard enthalpy of formation (DHof) of a substance is the enthalpy change that occurs in the formation of 1 mol of the substance from its elements when both products and reactants are in their standard states
EOS
Chapter 6: Thermochemistry

72. Standard Enthalpy
The enthalpy change for a reaction at standard conditions (25ºC, 1 atm , 1 M solutions)
Symbol DHº
When using Hess’s Law, work by adding the equations up to make it look like the answer.
The other parts will cancel out.

73. Standard Enthalpies of Formation
Standard enthalpies of formation, Hf°, are measured under standard conditions (25 °C and 1.00 atm pressure).
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74. Calculation of H C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l)
H = [3( kJ) + 4( kJ)] – [1( kJ) + 5(0 kJ)] = [( kJ) + ( kJ)] – [( kJ) + (0 kJ)] = ( kJ) – ( kJ) = kJ
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75. Example Given calculate DHº for this reaction DHº= -1300. kJ

76. Example Given DHº= +77.9kJ DHº= +495 kJ DHº= +435.9kJ
Calculate DHº for this reaction

77. Standard Enthalpies of Formation
Hess’s Law is much more useful if you know lots of reactions.
Made a table of standard heats of formation. The amount of heat needed to for 1 mole of a compound from its elements in their standard states.
Standard states are 1 atm, 1M and 25ºC
For an element it is 0
There is a table in Appendix 4 (pg A22)

78. Standard Enthalpies of Formation
Need to be able to write the equations.
What is the equation for the formation of NO2 ?
½N2 (g) + O2 (g) ® NO2 (g)
Have to make one mole to meet the definition.
Write the equation for the formation of methanol CH3OH.

79. 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.

80. 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.

81. Calculations Based on Standard Enthalpies of Formation
General Expression:
DHo = Snp × DHof (products) – Snr × DHof (reactants)
Each coefficient is multiplied by the standard enthalpy of formation for that substance
The sum of numbers for the reactants is subtracted from the sum of numbers for the products
EOS
With organic compounds, the measured DHof is often the standard enthalpy of combustion DHocomb
Chapter 6: Thermochemistry