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Thermodynamics
Internal
Energy
Study of energy and its
transformations
Thermochemistry
Study of energy changes
that accompany chemical
and physical changes.
Energy
• the capacity to do work or transfer heat
Internal energy
• total kinetic and potential energy in a
system
System:
the part of the universe
undergoing a change or being
studied
Surroundings:
region of space around system
(the rest of the universe)
In any system, energy is either
being stored or put in motion.
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).
Kinetic Energy
Kinetic energy is energy an object
possesses by virtue of its motion:
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Ek =  mv2
2
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K.E. (EK)= Energy of Motion
•Falling Body –
energy released as an object travels
through space.
rotate
vibrate
translate
Kinetic energy is
a function of
mass and
velocity:
1
2
Ek =  mv2
1
2
Ek =  mv2
•Thermal Energy –
energy associated with movement of
compounds, atoms within compounds,
atoms and subatomic particles.
P.E. (EP)= Stored Energy
•Positional Energy –
Energy stored relative to a
position under the influence of
gravity.
EP  mgh
•Chemical Energy –
Energy stored in bonds.
Potential Energy
• Potential energy is
energy an object
possesses by virtue
of its position or
chemical
composition.
• The most important
form of potential
energy in molecules
is electrostatic
potential energy, Eel:
Eel =
KQ1Q2
d
Conversion of Energy
• Add in Eel calculation
• Energy can be converted from one type to
another.
• For example, the cyclist in Figure 5.2 has
potential energy as she sits on top of the
hill.
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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.
Conversions of chemical
energy occur by: Heat
• Energy can also be
transferred as heat.
• q = Cs x m x ∆T
where q is heat, Cs is
specific heat, m is
mass and ∆T is the
change in temperature.
• Heat flows from
warmer objects to
cooler objects.
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
Conversions of chemical
energy occur by: 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.
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.
State Functions
Usually we have no way of knowing the
internal energy of a system; finding that
value is simply too complex a problem.
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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 depicted in Figure 5.9, the water
could have reached room temperature from
either direction.
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.
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.
First Law of Thermodynamics
“energy is neither created nor destroyed,
but changes form”
All processes in universe are accompanied
by a change in energy and since energy
must remain constant:
Esystem + Esurroundings= 0
Esystem = -Esurroundings
Changes in Internal Energy
• If E > 0, Efinal > Einitial
Therefore, the system absorbed energy
from the surroundings.
This energy change is called
endergonic.
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Changes in Internal Energy
• If E < 0, Efinal < Einitial
Therefore, the system released energy
to the surroundings.
This energy change is called exergonic.
There are two pathways by which
energy enters or exits a system:
1. Heat transferred (q) = heat absorbed
or emitted by system.
2. Work (w) = work done by or on
surroundings by system.
E → state function
q & w → path functions
E = q + w
Work is defined as the energy
required to move an object a
distance against a force
Descriptive Factors of A Chemical
System:
• (P) Pressure
w=Fd
• (V) Volume
However, in
chemical systems,
forces are difficult to
measure, so we must
put work into terms
we can measure
Work
We can measure the work done by the
gas if the reaction is done in a vessel that
has been fitted with a piston.
w = −PV
• (T) Temperature
• (n) amount
• Reactivity
notice no F or d
Using pressure:
P=F/A
Rearrange for Force:
F=PxA
By substitution:
w=PxAxd
Now, to measure most chemical systems a
cylinder can be used for the piston;
Therefore…
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…for a cylinder:
Therefore, by substitution:
V=Axh
w = -Pexternal ΔV
Therefore:
(1L . atm = 101.3 J )
ΔV = A x Δh
and:
d = Δh
So:
w=PxAxΔh
w > 0 (+w)
• surrounding does work
on the system (cylinder
contracts)
w < 0 (-w)
• system does work on the
surrounding (cylinder
expands)
Factors Determining Amount of
Heat Transferred
q = Cs x m x ∆T
1.
2.
3.
4.
m Amount of material
Grams of material
ΔT Temperature change
T = Tf - Ti
Cm Molar Heat Capacity (J/moleK)
Heat capacity of one mole of a substance
Cs Specific heat (J/gK)
Heat capacity of one gram of a substance
The negative sign has been added
to show that work is negative (or, that
energy leaves the system)
So, what does this mean?
By measuring P and ΔV of a
chemical system, we can determine the
work done on or by the system.
Remember, there are two
pathways energy enters or leaves
the system. One was work (w)
the second is the heat transferred
(q)
Heat transferred is defined as
the amount of heat energy lost or
gained by a system
q = Cs x m x ∆T
Specific Heat
Amount of heat needed to raise the
temperature of 1 gram of material 1
degree centigrade.
* What is the specific
 J 
heat of water?

 * You should memorize
 g  deg 
this number
(Handout)
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Note:
q = Cs x m x ∆T
Remember, a chemical system
is described by P, V, n, T, and
reactivity based on the properties
of species reacting. Notice P and V
are the variables of work. Also,
notice that T, n (in terms of mass)
and reactivity (in terms of specific
heat) are the variables of heat
transferred.
SI unit


From the equation of KE:



exothermic
q > 0 (+) endothermic
calorie
Commonly, energies are described in
the units of calories. A calorie is the
amount of heat needed to raise the
temperature of 1 gram H20, 1 degree
centigrade
1 cal = 4.184 J
EK = ½ mv2
2
Cp is the Heat Capacity of an object
Cp = Cs x m
q < 0 (-)
Joule(J)
the kinetic energy required to
move a 2Kg mass a distance of one
meter in one sec.

kg  m
 J 
s2

One common departure is the
expression
q = Cp T
where:
nutritional calorie
“Calorie”
1 Cal = 1 Kcal
James Joule
1818-1889
So:
• From Cs, m and ΔT, we can calculate q:
+ q = heat gained (feels cold)
- q = heat lost (feels Hot)
• From P and change in V, we can find w:
+ w = work done on system
- w = work done by system
Remember overall:
ΔE = q + w
Therefore:
 If Σ q + w > 0
Δ E = (+) endergonic

If Σ q + w < 0
Δ E = (-) exergonic
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