1. Chapter 7. THERMODYNAMICS: THE FIRST LAW
SYSTEMS, STATES, AND ENERGY
7.1 Systems
7.2 Work and Energy
7.3 Expansion Work
7.4 Heat
7.5 The Measurement of Heat
7.6 The First Law
7.7 A Molecular Interlude: The Origin of
Internal Energy
2012 General Chemistry I

2. SYSTEMS, STATES, AND ENERGY (Sections 7.1-7.7)
Thermodynamics deals with transformation (from one form to another) and transfer (from one place to another) of energy.
- System means the region in which we are interested
- Surroundings: everything else
- Universe (the system and the surroundings

3. Types of systems - Open system: exchanging both matter and
energy with the surroundings
- Closed system: a fixed amount of matter,
but exchanging energy with the surroundings
- Isolated system: no contact with its surroundings

4. 235s
Example 7.1
Identify the following systems as open, closed, or isolated:
Coffee in a very high quality thermos bottle
Coolant in a refrigerator coil
A bomb calorimeter in which benzene is burned
Gasoline burning in an automobile engine
Mercury in a thermometer
A living plant
Solutions

5. 7.2 Work and Energy DU = Ufinal – Uinitial
Work: the process of achieving motion against an
opposing force.
- unit: joule (J), 1 J = 1 kg·m2·s-2 = 1 N·m
Energy: the capacity of a system to do work:
Internal energy (U) is the total store of energy in a
system. In thermodynamics, "changes" in energy
(ΔU) are dealt with.
DU = Ufinal – Uinitial
Symbol w: the energy transferred to a system by doing
work; in the absence of other changes, DU = w

6. 7.3 Expansion Work
Expansion work: the work arising from a change in the
volume of a system. Table 7.1 shows expansion and
nonexpansion work.

7. Irreversible expansion work
Units of work:
Units of work
Irreversible expansion work
Only when the external pressure is constant during
the expansion.
The negative sign relates work to the system.
Note if Pex = 0, a vacuum, w = 0; free expansion (expansion
against vacuum)

9. 238s
Example 7.3
Air in a bicycle pump is compressed by pushing in the handle. If the inner diameter of the pump is 3.0 cm and the pump is depressed 20 cm with a pressure of 2.00 atm,
How much work is done in the compression?
Is the work positive or negative with respect to the air in the pump?

10. Solution

11. Reversible process can be reversed by an
infinitesimal change in a variable.
E.g. reversible, isothermal expansion of
an ideal gas
PexV = nRTconst = constant
- Work done is the area beneath the ideal gas isotherm
lying between the initial and the final volumes.

12. EXAMPLE 7.2
A piston confines mol Ar(g) in a volume of 1.00 L at 25 oC. Two
experiments are performed. (a) The gas is allowed to expand through
an additional 1.00 L against a constant pressure of 1.00 atm. (b) The
gas is allowed to expand reversibly and isothermally to the same
final volume. Which process does more work?
(a) Irreversible path:
(b) Reversible path:

14. 7.4 Heat
Heat is the energy transferred as a result of a temperature difference.
- Thermal energy of a system: the sum of the potential and kinetic energies arising from the chaotic thermal motion of atoms, ions, and molecules
-The energy transferred to a system as heat is q
Internal energy of the system changed by transferring energy as heat
DU = q
- Unit: cal, the energy needed to raise T of 1 g of water by 1 oC.
1 cal = J (exactly)
1 Cal (nutritional calorie) = 1 kcal

15. releasing heat into the surroundings Endothermic process:
Exothermic process:
releasing heat into the
surroundings
Endothermic process:
absorbing heat from the
surroundings
- Thermite reaction:

16. 7.5 The Measurement of Heat
Adiabatic Process: one performed in isolated system; no
heat exchanged with surroundings
Heat capacity, C is the ratio of the heat supplied to the rise in temperature observed.
heat supplied
temperature rise produced
Heat capacity =

17. Some Specific and Molar Heat Capacities
(Table 7.2)

19. - Calorimeter: a device in which heat transfer is monitored by recording the change in temperature that it produces
- Styrofoam calorimeter: q at constant pressure (qp)
- Bomb calorimeter:
q at constant volume (qV)
Types of
Calorimeters

21. Example 7.17
246s
(a) Calculate the heat that must be supplied to a g copper kettle containing g of water to raise its temperature from 22.0 oC to the boiling point of water, oC (See Table 7.2).
(b) What percentage of the heat is used to raise the temperature of the water?
Solution

22. 7.6 The First Law The first law of thermodynamics (closed system)
states that the change in internal energy (DU) is
the sum of the work and heat changes: it is applicable
to any process that begins and ends in equilibrium
states.
All the energies received are turned into the energy of the system: this is a form of the energy conservation law.

23. U is a state function: a property that depends only on the current state of the system and is independent of how that state was prepared.
Other state functions are P, V, T, H, S, G

24. - For any ideal gas, ΔU = 0 for an isothermal process.
No changes in the kinetic energy of an ideal gas if ΔT = 0
No intermolecular forces for ideal gas molecules
No changes in potential energy during expansion
or compression
No changes in the total energy; ΔU = 0

26. EXAMPLE 7.5
Suppose that 1.00 mol of ideal gas molecules maintained at 292 K and
3.00 atm expands from 8.00 L to L and a final pressure of 1.20 atm
by two different paths. (a) Path A is an isothermal, reversible expansion.
(b) Path B has two parts. In step 1, the gas is cooled at constant volume
until its pressure has fallen to 1.20 atm. In step 2, it is heated and allowed
to expand against a constant pressure of 1.20 atm until its volume is
20.00 L and T = 292 K. Determine for each path the work done, the heat
transferred, and the change in internal energy (w, q, and DU).

27. (a) From ,

28. (b) Step 1, w = 0
Step 2, from
Less work is done in the irreversible path and
less energy has to enter the system as heat
to maintain its temperature.

30. 250s
Example 7.13
In an adiabatic process, no energy is transferred as heat. Indicate whether each of the following statements about an adiabatic process in a closed system is always true, always false, or true in certain conditions (specify the conditions):
(a) ΔU = 0
(b) Q = 0
(c) q < 0
(d) ΔU = q
(e) ΔU = w
Solution

31. 7.7 A Molecular Interlude: The Origin of Internal Energy
U = K + EP
- A system at high temperature has a greater internal energy
than the same system at a lower temperature.
(Greater kinetic energy)
Equipartition theorem: The average value of each quadratic
contribution to the energy of a molecule in a sample
at a temperature T is equal to
- kB = 1.381×10–23 J·K-1; Boltzmann’s constant
- R = kBNA; RT = 2.48 kJ·mol-1 at 25 oC
Can be used for vibrational motion only when quantum effects
can be neglected, at high temperatures. (kBT >> ΔE)

32. - Molecules of N atoms; 3N degrees of freedom
(a) Center of mass
3 translational degrees
of freedom
(b) Linear molecule
2 rotational degrees
of freedom
3N – 5 vibrational
degrees of freedom
(c) Nonlinear molecule
3 rotational degrees
of freedom
3N – 6 vibrational
degrees of freedom

33. - At room temperature, the equipartition theorem holds for
translational and rotational motions.
(a) Monatomic ideal gas:
translational kinetic energy only (3 modes)
(b) Diatomic or linear ideal gas:
translational energy (3 modes) + rotational energy (2 modes)
(c) Nonlinear ideal gas:
translational energy (3 modes) + rotational energy (3 modes)

34. 252s
Example 7.25
7.25 Which molecular substance do you expect to have the higher molar heat capacity, NO or NO2? Why?
Solution

35. Chapter 7. THERMODYNAMICS: THE FIRST LAW
ENTHALPY
7.8 Heat Transfers at Constant Pressure
7.9 Heat Capacities at Constant Volume and
Constant Pressure
7.10 A Molecular Interlude:
The Origin of the Heat Capacities of Gases
7.11 The Enthalpy of Physical Change
7.12 Heating Curves
2012 General Chemistry I

36. 7.8 Heat Transfers at Constant Pressure
ENTHALPY (Sections )
- At constant volume and no nonexpansion work:
w = -PexDV = 0; DU = w + q = q
- However, most chemical reactions take place at a
constant pressure of about 1 atm.
7.8 Heat Transfers at Constant Pressure
- At constant pressure Pex, if no work other than
pressure-volume work is done, then

37. Enthalpy is defined as:
(constant pressure, pressure-volume work only)
Since U, P, and V are state functions, H is also a state function!

38. Exothermic and endothermic reactions
reaction (ΔH < 0)
Endothermic
reaction (ΔH > 0)
DH = -208 kJ

40. 7.9 Heat Capacities at Constant Volume and Constant Pressure
for any ideal gas

41. qV < qP - If CV and CP do not change with temperature,
qV = nCV,m ΔT qP = nCP,m ΔT
qV < qP

42. 7.10 A Molecular Interlude: The Origin of the Heat Capacities of Gases
- For monatomic ideal gases (see 7.7),
- For linear molecules (see 7.7),

43. Variation of molar heat
capacity of iodine vapor
at constant volume
(Fig. 7.20)

44. 255s
Example 7.29
Predict the contribution to the heat capacity CV,m made by molecular motions for each of the following atoms and molecules:
(a) HCN; (b) C2H6; (c) Ar; (d) HBr
Solutions

45. EXAMPLE 7.6
Calculate the final temperature and the change in internal energy
when 500 J of energy is transferred as heat to mol O2(g) at
298 K and 1.00 atm at (a) constant volume; (b) constant pressure.
Treat the gas as ideal.
(a)
= K
(b)

48. 7.11 The Enthalpy of Physical Change
Enthalpy of vaporization is the
difference in molar enthalpy
between the vapor and the
liquid states (> 0, always).
- Energy needed to separate
liquid molecules
- Temperature dependence:

49. DHfreez = Hm(solid) – Hm(liquid)
Enthalpy of fusion is the molar enthalpy change
that accompanies melting (fusion).
Enthalpy of freezing is the change in molar enthalpy
change of a liquid when it solidifies.
DHfreez = Hm(solid) – Hm(liquid)
In general:

51. Enthalpy of sublimation is the molar enthalpy change when a solid sublimes

52. Standard Enthalpies of Physical Change*
Table 7.3)

54. 7.12 Heating Curves
- Heating curve is the graph showing the variation in the temperature of a sample as it is heated at a constant rate at constant pressure and therefore at a constant rate of increase in enthalpy.
The steeper the slope of a heating curve, the lower
the heat capacity.
The horizontal sections correspond to phase changes:
melting and boiling.

55. Heating curve of water (Fig. 7.26)
- In water, the slope for liquid < those for solid or vapor,
the high heat capacity of the liquid is due largely to the
extensive hydrogen bonding network.

56. 259s
Example 7.41
The following data were collected for a new compound used in
cosmetics: DHfus = 10.0 kJ·mol-1, DHvap = 20.0 kJ·mol-1; heat capacities: 30 J·mol-1 for the solid; 60 J·mol-1 for the liquid; 30 J·mol-1 for the gas. Which heating curve below best matches the data for this compound?
Solution: (c)

57. Chapter 7. THERMODYNAMICS: THE FIRST LAW
THE ENTHALPY OF CHEMICAL CHANGE
7.13 Reaction Enthalpies
7.14 The Relation Between DH and DU
7.15 Standard Reaction Enthalpies
7.16 Combining Reaction Enthalpies: Hess’s Law
7.17 The Heat Output of Reactions
7.18 Standard Reaction Enthalpies
7.19 The Born-Haber Cycle
7.20 Bond Enthalpies
7.21 The Variation of Reaction Enthalpy with Temperature
2012 General Chemistry I

58. THE ENTHALPY OF CHEMICAL CHANGE (Sections 7.13-7.21)
7.13 Reaction Enthalpies
Thermochemical equation: consisting of a chemical
equation together with a statement of the reaction
enthalpy, the corresponding enthalpy change.

60. Example 7.43
261s
7.43 Carbon disulfide can be prepared from coke (an impure form of
carbon) and elemental sulfur:
4 C(s) + S8(s) → 4 CS2(l) DHo = kJ
(a) How much heat is absorbed in the reaction of 1.25 mol S8?
(b) Calculate the heat absorbed in the reaction of 197 g of carbon with an excess of sulfur.
(c) If the heat absorbed in the reaction was 415 kJ, how much CS2 was produced?

61. 7.14 The Relation Between DH and DU
- For reactions in liquids and solids only,
- If a gas is formed in the reaction,
DH = Hfinal – Hinitial = DU + (nfinal – ninitial)RT
= DU + DngasRT

62. 261s
Example 7.51
Oxygen difluoride is a colorless, very poisonous gas that reacts
rapidly with water vapor to produce O2, HF, and heat:
OF2(g) + H2O(g) → O2(g) + 2HF(g) DH = -318 kJ
What is the change in internal energy for the reaction of 1.00 mol OF2?
Solution

64. 7.15 Standard Reaction Enthalpies
- Reaction enthalpies depend on the physical states of
the reactants and products
88 kJ = enthalpy of vaporization
of water (44 kJ·mol-1) × 2
It is therefore useful if enthalpies
(DHo) can be referred to a
standard state.

65. Standard state: refers to a pure form at exactly 1 bar
or for a solute in a liquid solution: concentration
of 1 mol·L-1.
Standard reaction enthalpy, DHo is the reaction enthalpy when reactants in their standard states change into products in their standard states.
Most thermochemical data is reported for 25 oC
( K) but the temperature is not part of standard states.

66. 7.16 Combining Reaction Enthalpies: Hess’s Law
- Enthalpy is a state function; DH is independent of the path.
Hess’s law: the overall reaction enthalpy is the sum of
the reaction enthalpies of the steps into which the reaction
can be divided.
DHo = kJ

67. Consider the synthesis of propane, C3H8,
a gas use as camping fuel:
EXAMPLE 7.9
It is difficult to measure the enthalpy change of this reaction directly.
However, standard enthalpies of combustion reactions are easy to
measure. Calculate the standard enthalpy of this reaction from the
following experimental data:

69. 266s
Example 7.67
Calculate the reaction enthalpy for the synthesis of hydrogen chloride gas, H2(g) + Cl2(g) → 2 HCl(g), from the following data:
NH3(g) + HCl(g) → NH4Cl(s) DHo = kJ
N2(g) + 3 H2(g) → 2 NH3(g) DHo = kJ
N2(g) + 4 H2(g) + Cl2(g) → 2 NH4Cl(s) DHo = kJ
Solution

70. 7.17 The Heat Output of Reactions
Combustion reactions are important throughout
chemistry - they are always exothermic.
the standard enthalpy of combustion, DHco
is defined as the change in enthalpy per mole of
a substance that is burned in a combustion reaction
under standard conditions.
There is also the specific enthalpy of combustion:
the enthalpy of combustion per gram

71. Standard Enthalpies of Combustion at 25 oC
(Table 7.4)*

73. 7.18 Standard Enthalpies of Formation
Standard enthalpy of formation, DHfo is defined as
the standard reaction enthalpy per mole of formula units for the formation of a substance from its elements in their most stable form.

74. Standard Enthalpies of Formation at 25 oC (kJ mol-1)
(Table 7.5)

75. Standard reaction enthalpy can be calculated from standard enthalpies of formation of reactants and products.

77. Use the information of standard enthalpies of formation and the
EXAMPLE 7.12
Use the information of standard enthalpies of formation and the
enthalpy of combustion of propane gas to calculate the enthalpy
of formation of propane, a gas commonly used for camping stoves
and outdoor barbecues.
= kJ
= kJ

79. 273s
Example 7.73
Calculate the standard enthalpy of formation of PCl5(s) from the
standard enthalpy of formation of PCl3(l) (DHof(PCl3,l) = kJ·mol-1)
and PCl3(l) + Cl2(g) → PCl5(s), DHo = -124 kJ.
Solution

80. 7.19 The Born-Haber Cycle Lattice enthalpy of the solid, DHL
- Lattice enthalpies at 25oC (kJ·mol-1)(Table 7.)

81. Born-Haber cycle: a closed path of steps,
one of which is the formation of a solid
lattice from the gaseous ions

82. Lattice enthalpy of KCl
EXAMPLE 7.13
Lattice enthalpy of KCl
DHf (K, atoms)
DHf (Cl, atoms)
I(K)
-Eea(Cl)
-DHf(KCl)
K+Cl-(s) K+(g) + Cl-(g)

84. 7.20 Bond Enthalpies
Bond enthalpy, DHB is the difference between the standard molar enthalpies of a molecule, X-Y, and its fragments X and Y in the gas phase.
E.g.
- Bond breaking is always endothermic and bond formation is always exothermic.

85. Mean bond enthalpy: small variations of a specific bond enthalpy in polyatomic molecules give a guide to the average bond strength.
- The enthalpy change from a liquid (or solid) sample,

86. EXAMPLE 7.14
Estimate the enthalpy of the reaction between gaseous iodoethane
and water vapor:
- Breaking the bonds (reactants)
DHBo(C-I) + DHBo(O-H) =
- Forming the bonds (products)
DHBo(C-O) + DHBo(H-I) =
- The overall enthalpy change:

88. 277s
Example 7.81
7.81 Use the bond enthalpies in the table to estimate the reaction enthalpy
for 3 C2H2(g) → C6H6(g)
Solution

89. 7.21 The Variation of Reaction Enthalpy with Temperature
- The enthalpies of both reactants and products increase with temperature.

90. Kirchhoff’s law

91. EXAMPLE 7.15
The standard enthalpy of reaction of N2(g) + 3H2(g) → 2NH3(g) is
kJ·mol-1 at 298 K. The industrial synthesis takes place at 450 oC.
What is the standard reaction enthalpy at the latter temperature?

93. 278s
Example 7.87
(a) Calculate the enthalpy of vaporization of benzene (C6H6) at K. The standard enthalpy of formation of gaseous benzene is kJ·mol-1, and that of liquid benzene is kJ·mol-1.
Solution

94. 278s
(b) Given that, for liquid benzene, CP,m = J·mol-1·K-1 and that,
for gaseous benzene, CP,m = J·mol-1·K-1, calculate the enthalpy of vaporization of benzene at its boiling point (353.2 K).

95. 278s
(c) Compare the value obtained in part (b) with the real value of 30.8 kJ·mol-1. What is the source of difference between these numbers?