1. Chapter 6 Thermochemistry
General Chemistry: An Integrated Approach Hill, Petrucci, 4th Edition
Chapter 6 Thermochemistry
Mark P. Heitz
State University of New York at Brockport
© 2005, Prentice Hall, Inc.

2. 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

3. 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
Chapter 6: Thermochemistry

4. Chapter 6: Thermochemistry
Work
Work is the product of the force in the direction of motion and the distance the object is moved
Work = force × distance  energy (J)
Collisions in the real world are not perfectly elastic
EOS
Energy transfer occurs as a result of inelastic collisions
e.g., the ball loses height
Chapter 6: Thermochemistry

5. 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

6. Chapter 6: Thermochemistry
Types of systems one can study:
OPEN
Matter
Energy
CLOSED
Matter
Energy
EOS
ISOLATED
Matter
Energy
Chapter 6: Thermochemistry

7. 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

8. 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

9. 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

10. Heat Transfer Mechanism Illustrated
Inelastic molecular collisions are responsible for heat transfer
EOS
Chapter 6: Thermochemistry

11. Heat Transfer Illustrated
Chapter 6: Thermochemistry

12. 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
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13. 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
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14. 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

15. Chapter 6: Thermochemistry
Calculating Work (w)
SYSTEM
WORK
If work is done on the system, the system gains energy equal to +w
EOS
compression (or expansion) of a gas
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16. 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
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17. 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

18. First Law of Thermodynamics
The Law of Conservation of Energy states that in a physical or chemical change, energy can be exchanged between a system and its surroundings, but no energy can be created or destroyed
The change in U is related to the energy exchanges that occur as heat (q) and work (w)
EOS
The First Law: DU = q + w
Chapter 6: Thermochemistry

19. First Law – Sign Conventions
Energy entering a system carries a positive sign: if heat is absorbed by the system, q > 0. If work is done on a system, w > 0
Energy leaving a system carries a negative sign: if heat is given off by the system, q < 0. If work is done by a system, w < 0
EOS
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20. Chapter 6: Thermochemistry
Heats of Reaction
The heat of reaction (qrxn) is the quantity of heat exchanged between the system and its surroundings
Examples – for exothermic reactions,
in isolated systems, system T 
in non-isolated systems, heat is given off to the surroundings, i.e., q < 0
EOS
– endothermic reactions,
in isolated systems, system T 
in non-isolated systems, heat is absorbed from the surroundings, i.e., q > 0
Chapter 6: Thermochemistry

21. Conceptualizing an Exothermic Reaction
Surroundings are at 25 °C
Typical situation: some heat is released to the surroundings, some heat is absorbed by the solution.
35.4 °C
32.2 °C
25 °C
Hypothetical situation: all heat is instantly released to the surroundings. Heat = qrxn
In an isolated system, all heat is absorbed by the solution. Maximum temperature rise.
Chapter 6: Thermochemistry

22. Internal Energy Changes
w = –PDV and DU = q + w
For systems where the reaction is carried out at constant volume, DV = 0 and DU = qV
All the thermal energy produced by conversion from chemical energy is released as heat
EOS
Because the reaction is exothermic, both qV and DU are negative
Chapter 6: Thermochemistry

23. Internal Energy Changes
w = –PDV and DU = q + w
For systems where the reaction is carried out at constant pressure, DU = qP – PDV or qP = DU + PDV
Most of the thermal energy is released as heat, but some is work used to expand the system against the surroundings
EOS
The quantity of heat liberated is somewhat less than in the constant-volume case
Chapter 6: Thermochemistry

24. Chapter 6: Thermochemistry
Example 6.2
The internal energy of a fixed quantity of an ideal gas depends only on its temperature. If a sample of an ideal gas is allowed to expand against a constant pressure at a constant temperature, (a) what is ∆U for the gas? (b) Does the gas do work? (c) Is any heat exchanged with the surroundings?
Analysis and Conclusions
(a) Because the expansion occurs at a constant temperature, the expanded gas (state 2) is at a lower pressure than the compressed gas (state 1) but the temperature is unchanged. Because the internal energy of the ideal gas depends only on the temperature, U2 = U1 and ∆U = U2 – U1 = 0.
(b) The gas does work in expanding against the confining pressure, P. The pressure–volume work is w = –P∆V, as was illustrated in Figure 6.8. The work is negative because it is done by the system.
(c) The work done by the gas represents energy leaving the system. If this were the only energy exchange between the system and its surroundings, the internal energy of the system would decrease, and so would the temperature. However, because the temperature remains constant, the internal energy does not change. This means that the gas must absorb enough heat from the surroundings to compensate for the work that it does in expanding: q = –w. And, according to the first law of thermodynamics, ∆U = q + w = –w + w = 0.
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25. Chapter 6: Thermochemistry
Example 6.2 continued
In an adiabatic process, a system is thermally insulated from its surroundings so that there is no exchange of heat (q = 0). If an ideal gas undergoes an adiabatic expansion against a constant pressure, (a) does the gas do work? (b) Does the internal energy of the gas increase, decrease, or remain unchanged? (c) What happens to the temperature?
Exercise 6.2A
yes w = - PΔV
ΔU = q + w = 0 + w = w = - PΔV since ΔV >0, ΔU< 0
For an ideal gas U ~ T so if U decreases, T decreases
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26. Chapter 6: Thermochemistry
Enthalpy
Most heats of reaction are measured at constant pressure … it is useful to have a function equal to qP
Enthalpy (H) is the sum of the internal energy and the pressure–volume product of a system
qP = DH = DU + PDV
Enthalpy is an extensive property (depends on how much of the substance is present)
EOS
Enthalpy is a state function. U, P, and V are all state functions, therefore H must be a state function also
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27. Chapter 6: Thermochemistry
Enthalpy Diagrams
EOS
Chapter 6: Thermochemistry

28. Same magnitude; different signs.
Reversing a Reaction
DH changes sign when a process is reversed.
Therefore, a cyclic process has the value DH = 0.
Same magnitude; different signs.
Chapter 6: Thermochemistry

29. Chapter 6: Thermochemistry
Using DH
Values are measured experimentally
Negative values indicate exothermic reactions
Positive values indicate endothermic reactions
Changes sign when a process is reversed. Therefore, a cyclic process has the value DH = 0
EOS
For problem-solving, one can view heat being absorbed in an endothermic reaction as being like a reactant and heat being evolved in an exothermic reaction as being like a product
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30. Chapter 6: Thermochemistry
Calorimetry
We measure heat flow using calorimetry.
A calorimeter is a device used to make this measurement.
A “coffee cup” calorimeter may be used for measuring heat involving solutions.
A “bomb” calorimeter is used to find heat of combustion; the “bomb” contains oxygen and a sample of the material to be burned.
Chapter 6: Thermochemistry

31. 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

32. 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
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33. Chapter 6: Thermochemistry
Specific Heats
EOS
Chapter 6: Thermochemistry

34. Hess’s Law of Constant Heat Summation
The heat of a reaction is constant, regardless of the number of steps in the process
DHoverall = S DH’s of individual reactions
When it is necessary to reverse a chemical equation, change the sign of DH for that reaction
EOS
When multiplying equation coefficients, multiply values of DH for that reaction
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35. Chapter 6: Thermochemistry
An Enthalpy Diagram
EOS
Chapter 6: Thermochemistry

36. Standard State Conditions
The standard state of a solid or liquid substance is the pure element or compound at 1 atm pressure and the temperature of interest
Gaseous standard state is the “ideal gas” at 1 atm pressure and the temperature of interest
EOS
e.g., at 1 atm, 25 oC standard state for Hg is liquid, C is solid, water is liquid, He is gas
Chapter 6: Thermochemistry

37. 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

38. Standard Enthalpies of Formation at 25 oC
EOS
Chapter 6: Thermochemistry

39. 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

40. Standard Enthalpies of Formation of Ions in Aqueous Solution
Chapter 6: Thermochemistry

41. Chapter 6: Thermochemistry
Combustion Fuels
Fossil Fuels: Coal, Natural Gas, and Petroleum
A fuel is a substance that burns with the release of heat
EOS
These fossil fuels were formed over a period of millions of years from organic matter that became buried and compressed under mud and water
Chapter 6: Thermochemistry

42. Chapter 6: Thermochemistry
Respiration Foods
Foods: Fuels for the Body
The three principal classes of foods are fats, proteins, and carbohydrates
EOS
1 Food Calorie (Cal) is equal to 1000 cal
(or 1 kcal)
Chapter 6: Thermochemistry

43. Chapter 6: Thermochemistry
Summary of Concepts
Thermochemistry concerns energy changes in physical processes or chemical reactions
Thermochemical ideas include the notion of a system and its surroundings; the concepts of kinetic energy, potential energy, and internal energy; and the distinction between two types of energy exchanges: heat (q) and work (w)
EOS
Internal energy (U) is a function of state
Chapter 6: Thermochemistry

44. Chapter 6: Thermochemistry
Summary of Concepts
Enthalpy (H) is a function based on internal energy, but modified for use with constant-pressure processes
The first law of thermodynamics relates the heat and work exchanged between a system and its surroundings to changes in the internal energy of a system
EOS
A calorimeter is used to measure quantities of heat
Chapter 6: Thermochemistry

45. Chapter 6: Thermochemistry
Summary of Concepts
The concepts of standard state, a standard enthalpy change, and a standard enthalpy of formation are important in thermochemical calculations
Some practical applications of thermochemistry deal with the heats of combustion of fossil fuels and the energy content of foods
EOS
Chapter 6: Thermochemistry