Transcript Chapter 5 Thermochemistry

Thermochemistry
Thermodynamics = study of energy and
its transformations
Thermochemistry = study of chemical
reactions involving changes in heat
energy
Thermochemistry
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Energy
Energy = the ability to do work or transfer
heat energy.
Work = energy used to cause an object
with mass to move (w = f x d)
Heat = energy used to cause the
temperature of an object to increase
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Major Types of Energy
Potential energy = energy an object
possesses by virtue of its position or
chemical composition.
Kinetic energy = energy an object
possesses by virtue of its motion.
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Kinetic Energy
1
KE =  m v
2
2
m = mass in kilograms (kg)
v = velocity in meters per second (m/s)
KE = kinetic energy in joules (J)
1 Joule = 1 kg-m2/s2
A mass of 2 kg moving at a speed of one meter per second
possesses a kinetic energy of 1 Joule.
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Potential Energy
PE = m g h
m = mass in kilograms (kg)
g = acceleration due to gravity (9.8 m/s2)
h = height (m)
PE = potential energy in joules (J)
1 Joule = 1 kg-m2/s2
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Units of Energy
• The SI unit of energy is the joule (J).
• An older, non-SI unit is still in
widespread use: the calorie (cal).
1 cal = 4.184 J
• 1000 calories = one nutritional Calorie
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First Law of Thermodynamics
• Energy is neither created nor destroyed, but
it can undergo a transformation from one
type to another. (Law of Conservation of
Energy)
• The total energy of the universe is a constant.
• The energy lost by a system must equal the
energy gained by its surroundings, and vice
Thermochemistry
versa.
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System and Surroundings
System = the molecules
we want to study (here,
the hydrogen and
oxygen molecules).
Surroundings = everything
else (here, the cylinder
and piston).
Thermochemistry
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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.
E = Efinal − Einitial
(It’s a state function)
• If E is positive, the system absorbed energy
from the surroundings.
• If E is negative, the system released energy
Thermochemistry
to the surroundings.
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E = q + w
• 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.
Thermochemistry
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q, w, and their signs
+ q = system gains or takes in heat
- q = system loses or gives off heat
+ w = work is done on the system by the
surroundings (piston pushed in)
- w = work is done by the system on its
Thermochemistry
surroundings (piston moves out)
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Example
As hydrogen and oxygen gas are ignited in a cylinder,
the system loses 550 J of heat to its surroundings.
The expanding gases move a pistion to do 240 J of
work on its surroundings. E for system = ?
Answer:
E = q + w
E = (-550 J) + (-240 J)
E = - 790 J
What does it mean?
The system gave off 790 J of energy to its surroundings
Thermochemistry
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Enthalpy & H
• The symbol for enthalpy is H.
• Enthalpy is the internal energy plus the
product of pressure and volume:
H = E + PV
• At constant pressure:
H = E = q
• So at constant pressure,
H = heat lost or gained by the system.
Thermochemistry
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Recall: Endothermic
• When heat is absorbed (taken in) by the
system from the surroundings, the process is
endothermic.
H = Hfinal − Hinitial
H = Hproducts − Hreactants
H = positive value
for endothermic
Thermochemistry
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Recall: Exothermic
When heat is released (given off) by the
system into the surroundings, the process is
exothermic.
H = Hfinal − Hinitial
H = Hproducts − Hreactants
H = negative value
for exothermic
Thermochemistry
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