Download From the first law of thermodynamics

Thermochemistry
David P. White
University of North Carolina, Wilmington
Chapter 5
Copyright 1999, PRENTICE HALL
Chapter 5
1
The Nature of Energy
Kinetic and Potential Energy
From Physics:
• Force is a push or pull on an object.
• Work is the product of force applied to an object over
a distance:
w=Fd
• Energy is the work done to move an object against a
force.
• Kinetic energy is the energy of motion:
2
1
E k  mv
2
Copyright 1999, PRENTICE HALL
Chapter 5
2
The Nature of Energy
Kinetic and Potential Energy
• Potential energy is the energy an
object possesses by virtue of its
position.
• Potential energy can be converted
into kinetic energy. Example: a ball
of clay dropping off a building.
Copyright 1999, PRENTICE HALL
Chapter 5
3
The Nature of Energy
Energy Units
SI Unit for energy is the joule, J:
E k  1 mv 2  1 2 kg 1 m/s 2
2
2
 1 kg m 2 / s 2
1J
We sometimes use the calorie instead of the joule:
1 cal = 4.184 J (exactly)
A nutritional Calorie:
1 Cal = 1000 cal = 1 kcal
Copyright 1999, PRENTICE HALL
Chapter 5
4
The Nature of Energy
Systems and Surroundings
System: part of the universe we
are interested in.
Surroundings: the rest of the
universe.
Copyright 1999, PRENTICE HALL
Chapter 5
5
First Law of Thermodynamics
Internal Energy
• Internal Energy:
total energy of a
system.
• Cannot measure
absolute internal
energy.
• Change in internal
energy, DE = Efinal Einitial
Copyright 1999, PRENTICE HALL
Chapter 5
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First Law of Thermodynamics
Relating DE to Heat and Work
Energy cannot be created or destroyed.
Energy of (system + surroundings) is constant.
Any energy transferred from a system must be
transferred to the surroundings (and vice versa).
From the first law of thermodynamics:
when a system undergoes a physical or chemical change, the
change in internal energy is given by the heat added to or
absorbed by the system plus the work done on or by the
system:
DE = q + w
Copyright 1999, PRENTICE HALL
Chapter 5
7
First Law of Thermodynamics
Relating DE to Heat and Work
Copyright 1999, PRENTICE HALL
Chapter 5
8
First Law of Thermodynamics
Relating DE to Heat and Work
Copyright 1999, PRENTICE HALL
Chapter 5
9
First Law of Thermodynamics
Endothermic and Exothermic Processes
Endothermic: absorbs heat from the surroundings.
Exothermic: transfers heat to the surroundings.
An endothermic reaction feels cold.
An exothermic reaction feels hot.
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Chapter 5
10
First Law of Thermodynamics
State Functions
State function: depends only on the initial and final
states of system, not on how the internal energy is used.
Copyright 1999, PRENTICE HALL
Chapter 5
11
First Law of Thermodynamics
State Functions
Copyright 1999, PRENTICE HALL
Chapter 5
12
Enthalpy
Enthalpy, H: Heat transferred between the system and
surroundings carried out under constant pressure.
Can only measure the change in enthalpy:
DH = Hfinal - Hinitial = qP
Copyright 1999, PRENTICE HALL
Chapter 5
13
Enthalpies of Reaction
For a reaction
DHrxn = H(products) - H (reactants)
Enthalpy is an extensive property (magnitude DH is
directly proportional to amount):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
DH = -802 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) DH = -1604 kJ
When we reverse a reaction, we change the sign of DH:
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g)
DH = +802 kJ
Change in enthalpy depends on state:
H2O(g)  H2O(l) DH = -88 kJ
Copyright 1999, PRENTICE HALL
Chapter 5
14
Calorimetry
Heat Capacity and Specific Heat
Calorimetry = measurement of heat flow.
Calorimeter = apparatus that measures heat flow.
Heat capacity = the amount of energy required to raise
the temperature of an object (by one degree).
Molar heat capacity = heat capacity of 1 mol of a
substance.
Specific heat = specific heat capacity = heat capacity of
1 g of a substance.
q = (specific heat)  (grams of substance)  T.
Be careful of the sign of q.
Copyright 1999, PRENTICE HALL
Chapter 5
15
Calorimetry
Heat Capacity and Specific Heat
Copyright 1999, PRENTICE HALL
Chapter 5
16
Calorimetry
Constant-Pressure
Calorimetry
Atmospheric pressure is
constant!
DH = qP
qrxn = -qsoln = -(specific heat of
solution)  (grams of solution) 
DT.
Copyright 1999, PRENTICE HALL
Chapter 5
17
Calorimetry
Bomb Calorimetry (Constant-Volume
Calorimetry)
Reaction carried out under constant volume.
Use a bomb calorimeter.
Usually study combustion.
Copyright 1999, PRENTICE HALL
Chapter 5
18
Calorimetry
Bomb Calorimetry (Constant-Volume
Calorimetry)
qrxn = -CcalorimeterT.
Copyright 1999, PRENTICE HALL
Chapter 5
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Hess’s Law
• Hess’s law: if a reaction is carried out in a number of
steps, DH for the overall reaction is the sum of DH for
each individual step.
• For example:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
DH = -802 kJ
2H2O(g)  2H2O(l)
DH = -88 kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
DH = -890 kJ
Copyright 1999, PRENTICE HALL
Chapter 5
20
Hess’s Law
In the above enthalpy diagram note that
DH1 = DH2 + DH3
Copyright 1999, PRENTICE HALL
Chapter 5
21
Enthalpies of Formation
• If 1 mol of compound is formed from its constituent
elements, then the enthalpy change for the reaction is
called the enthalpy of formation, DHof .
• Standard conditions (standard state): 1 atm and 25 oC
(298 K).
• Standard enthalpy, DHo, is the enthalpy measured
when everything is in its standard state.
• Standard enthalpy of formation: 1 mol of compound
is formed from substances in their standard states.
• If there is more than one state for a substance under
standard conditions, the more stable one is used.
Copyright 1999, PRENTICE HALL
Chapter 5
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Enthalpies of Formation
• Standard enthalpy of formation of the most stable
form of an element is zero.
Copyright 1999, PRENTICE HALL
Chapter 5
23
Enthalpies of Formation
Using Enthalpies of Formation to Calculate
Enthalpies of Reaction
We use Hess’ Law to calculate enthalpies of a reaction
from enthalpies of formation.
DHrxn = DH1 + DH2 + DH3
Copyright 1999, PRENTICE HALL
Chapter 5
24
Enthalpies of Formation
Using Enthalpies of Formation to Calculate
Enthalpies of Reaction
For a reaction:
DH rxn   nDH  f products    mDH  f reactants
Copyright 1999, PRENTICE HALL
Chapter 5
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Foods and Fuels
Foods
•Fuel value = energy released when 1 g of substance is
burned.
•1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal.
•Energy in our bodies comes from carbohydrates and
fats (mostly).
•Intestines: carbohydrates converted into glucose:
C6H12O6 + 6O2  6CO2 + 6H2O, DH = -2816 kJ
•Fats break down as follows:
2C57H110O6 + 163O2  114CO2 + 110H2O, DH = -75,520 kJ
•Fats: contain more energy; are not water soluble, so are
good for energy storage.
Copyright 1999, PRENTICE HALL
Chapter 5
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Foods and Fuels
Foods
Copyright 1999, PRENTICE HALL
Chapter 5
27
Foods and Fuels
Fuels
U.S.: 1.0 x 106 kJ of fuel per day.
Most from petroleum and natural gas.
Remainder from coal, nuclear, and hydroelectric.
Fossil fuels are not renewable.
Copyright 1999, PRENTICE HALL
Chapter 5
28
Foods and Fuels
Fuels
Fuel value = energy released when 1 g of substance is
burned.
Hydrogen has great potential as a fuel with a fuel value
of 142 kJ/g.
Copyright 1999, PRENTICE HALL
Chapter 5
29
Thermochemistry
End of Chapter 5
Copyright 1999, PRENTICE HALL
Chapter 5
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