Transcript Chem 1202 - LSU Department of Chemistry

Watkins
Chapter 5
Thermochemistry
Chapter 5
(sections 1 thru 4 only)
Force, Energy, Heat, Work
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Chapter 5
The Nature of Force
Mechanical Force (f ) is a push or pull applied to a
material object
• A material object has mass (m, kg)
• Mechanical force changes the velocity (m/s) of
the object: acceleration ( a, m/s2 or ms-2)
• Newton's 2nd law of motion:
f =  ma (kgm/s2 or kgms-2 = Newton)
• +f is a push, -f is a pull)
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Chapter 5
The Nature of Force
Gravitational Force ( f ) is a pull between two
masses, m and M
• If the masses are separated by distance d ...
f = -GmM /d 2
• If M is earth mass and d is distance to earth
center:
f = -m(GM /d 2) = -mg
• g = 9.80 m/s2 at mean sea-level (“gravitational
acceleration”)
• "weight" = -f = mg
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Chapter 5
The Nature of Force
Coulombic Force ( f ) is a push or pull between
two charged particles (e.g., ions)
• If the ions, separated by distance d, have charges
±z1 and ±z2 ...
f = k z1 z2 / d 2
• Two kinds of Coulombic force:
Repulsion pushes if (+z1, +z2) or (-z1, -z2)
( f > 0)
Attraction pulls if (-z1, +z2) or (+z1, -z2)
( f < 0)
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Chapter 5
The Nature of Force
Some of the different kinds of force:
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Mechanical
Gravitational
Coulombic
Torque
Nuclear weak/strong
Magnetic
Electromagnetic
Frictional/Drag
Centrifugal/Centripetal
Coriolis
Static
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Force f is a push
or pull on a mass,
and is usually
associated with
motion of that
mass.
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Chapter 5
The Nature of Energy
Force f moves an object a distance d
•
w = f d (Nm = kgm2/s2 = kgm2 s-2 = Joule)
Work w is also called Mechanical Energy
When the object moves at velocity v
• The energy of motion is called Kinetic Energy
Ek = ½ m v2 (kgm2/s2 = Joule)
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Chapter 5
The Nature of Energy
Stored Energy is called Potential
Energy; for example, the energy of
gravitational attraction between an
object and the center of the earth.
• Example: a ball on top of a wall.
Gravitational PE = m g h
• Potential energy is converted into
kinetic energy when the ball falls.
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Chapter 5
The Nature of Energy
Potential energy can also be stored
• in a battery or a capacitor
(electrical potential energy);
Each of these
forms of energy
• in a flywheel
can be converted
(kinetic potential energy);
to other forms by
• in the bonds between atoms
various processes
(chemical potential energy);
and/or machines.
• in an atomic nucleus
(nuclear potential energy).
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The Nature of Energy
Chapter 5
Some of the different forms of energy:
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Mechanical Energy (E = fd, E = PDV )
Gravitational Energy (E = -GmM/d)
Kinetic Energy (E = 1/2mv2)
Coulombic Energy (E = kz1z2/d)
Radiant (light) Energy (E = hn) (chapter 6)
Nuclear Energy (E = mc2)
Electrical Energy (E = VQ = VIt )(chapter 20)
Heat flow (q)
Chemical Heat (DHrxn)
Chemical Potential Energy (DGrxn) (chapter 19)
Waste Energy (TDS) (chapter 19)
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Chapter 5
The Nature of Energy
Different units of energy:
• Joule, the SI Unit of energy: f × d
1 J = 1 Nm = 1 kg m2 s-2
• Calorie (used in the past for heat flow)
1 cal = 4.184 J (exactly)
• Nutritional Calorie
1 Cal = 1000 cal = 1 kcal = 4.184 kJ
• Electron Volt (used in atomic physics)
1 ev = 1.602×10-19 J
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Chapter 5
The Nature of Energy
We perceive the form of energy
called "heat" as subjectively different
from other forms of energy.
All other forms of energy are
distinquished from heat, and they are
collectively called "work".
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Chapter 5
Thermodynamics
Literally: the study of heat movement
Generally: the study of the movement and
conversion of all forms of energy
The First Law of Thermodynamics
Energy cannot be created or destroyed; it can be
moved from one place to another, and it can
be converted from one form to another.
Thermodynamics is the set of rules that govern
• the movement of energy from one place to
another;
• the conversion of energy from one form to
another.
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Chapter 5
Thermodynamics
DEFINITIONS
System (sys)
that part of the
universe in which we
are interested.
Surroundings (surr)
the rest of the
universe.
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Chapter 5
Thermodynamics
DEFINITIONS
Internal Energy
Total energy of system
Esys
We cannot measure Esys
We can measure the change
in Esys during a process:
DEsys
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Chapter 5
Thermodynamics
Initial State
H2(g) & O2(g)
DEFINITIONS
Internal Energy
Total energy of system
Esys
We cannot measure Esys
We can measure the change
in Esys during a process:
DEsys
For example, during a
chemical reaction
DEsys = Efinal - Einitial
Esys
H2(g) + ½ O2(g) → H2O(g)
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DEsys < 0
H2O(g)
Final State
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Chapter 5
Thermodynamics
Final State
H2(g) & O2(g)
DEFINITIONS
Internal Energy
Total energy of system
Esys
We cannot measure Esys
We can measure the change
in Esys during a process:
DEsys
For example, during a
chemical reaction
DEsys = Efinal - Einitial
Esys
H2O(g) → H2(g) + ½ O2(g)
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DEsys > 0
H2O(g)
Initial State
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Chapter 5
First Law of Thermodynamics
Relating DE to Heat and Work
When a system undergoes a physical or
chemical change, the change in its internal
energy is equal to the sum of
q, the heat flow into or out of the system
w, the work done on or by the system:
DEsys = q + w
Work w is all forms of energy other than q.
Both q and w are signed.
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Chapter 5
First Law of Thermodynamics
Example I
"Endothermic"
q>0
DEsys = q + w
Heat
Surroundings
w>0
DEsys > 0
System
Work
work is done on
the system by the
surroundings
Process I occurs in the system
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Chapter 5
First Law of Thermodynamics
Example II
"Exothermic"
q<0
DEsys = q + w
Heat
Surroundings
w<0
DEsys < 0
System
Work
work is done by
the system on the
surroundings.
Process II occurs in the system
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Chapter 5
First Law of Thermodynamics
Example III
"Exothermic"
q<0
DEsys = q + w
Heat
Surroundings
w>0
DEsys ?
System
Work
work is done on
the system by the
surroundings.
Process III occurs in the system
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Chapter 5
First Law of Thermodynamics
Example IV
“Endothermic"
q>0
DEsys = q + w
Heat
Surroundings
w<0
DEsys ?
System
Work
work is done by
the system on the
surroundings.
Process IV occurs in the system
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DE = q + w
Energy E
Chapter 5
One form of work is expansion or “PV” work:
w = -PDV, Pressure (atm) × DVolume (L)
1 L.atm = 101.325 J
Often, PV work is the only form of work in a
chemical process, and if DV = 0, then
wV = -PDV = 0 and
subscript means “this
variable is held constant
during the process”.
DE = qv
Heat flow qV in a constant volume process is a
direct measure of DE.
DV = 0 for a rigid container.
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DE = qV
Energy E
Chapter 5
However, most chemical reactions are
not carried out at constant volume.
They are usually carried out at constant
pressure (in the open atmosphere).
In some reactions, both heat and work
are involved (e.g., explosions).
So a new energy function was invented
for contant pressure processes.
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DE = qV
Energy E
Chapter 5
For a contant pressure process (DP = 0),
if the system expands (DV > 0) it does
work on the surroundings:
wp = -PDV
DE = qp + wp = qP - PDV
qP = DE + PDV = DH
Energy H is called Enthalpy
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DE = qV DH = qP
Chapter 5
Enthalpy H
At contant pressure, H is the “total heat energy
in the system”, also called the “internal heat”
or “heat content”
When qp Joules of heat flow into or out of the
system during the process (reaction) the heat
content changes by an amount qp:
DH = Hfinal - Hinitial = qP
We cannot measure either H, but we can
measure DH = qp, the flow of heat between sys.
and surr. during a constant pressure process.
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DE = qV DH = qP
Chapter 5
Enthalpy of Reaction
Consider an open
beaker in which a
reaction takes place at
constant atmospheric
pressure:
Reactants → Products
aA + bB → cC + dD
HReact
HProd
qp = HProd - HReact
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Enthalpy of Reaction
There are two kinds
of chemical reaction
H ("heat content")
DE = qV DH = qP
Chapter 5
reactants H
products H
qP = DH < 0
heat out
exo
qP = DH > 0
heat in
endo
Consider an open
beaker in which a
reaction takes place at
constant atmospheric
pressure:
Reactants → Products
HReact
or
products H
HProd
qp = HProd - HReact
reactants H
positive DH = endothermic; negative DH = exothermic
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DE = qV DH = qP
Chapter 5
Enthalpy of Reaction
For a chemical reaction, qp is called the
reaction heat: qP = qrxn = DHrxn
Reaction heat is an extensive property - the
magnitude of DH is directly proportional to
amounts of reactants and products:
1CH4(g) + 2O2(g)  1CO2(g) + 2H2O(g) DH1 = -803 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) DH2 = -1606 kJ
Multiply a reaction, multiply the heat.
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DE = qV DH = qP
Chapter 5
Enthalpy of Reaction
When a reaction is reversed, the sign of DH changes:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) DH = -803 kJ
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g) DH = +803 kJ
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DE = qV DH = qP
Chapter 5
Enthalpy of Reaction
Enthalpy depends on state: at 100o C
H2O(l)  H2O(g)
Hl
Hg
DH = Hg - Hl
DHvap = +88 kJ
DH = Hg – Hl > 0
H[steam] > H[water]
You have to add heat to water to produce steam!
H[vapor] > H[liquid]
The process called “vaporization” is endothermic for all
substances at all temperatures.
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DE = qV DH = qP
Chapter 5
Enthalpy of Reaction
Four things to remember
1. DHrxn = qp = Hprod -Hreact
2. When you multiply a reaction, multiply DH by
the same factor.
3. When you reverse a reaction, DH changes sign.
4. DH depends on state; in general, for a specific
substance: H(s) < H(l) < H(g)
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DE = qV DH = qP
Chapter 5
Enthalpy of Reaction
You can measure reaction heat (calorimetry)
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You can calculate reaction heat in two ways:
• Chemical bond enthalpies (Chapter 8)
• Hess's Law (Chapter 5, Chem 1422)
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Chapter 5
Chapter 5, Part 1 Problems
• Energy Calculations
• ½mv2, mgh, PDV, fd
• First Law; DE = q + w; DE = qV
• DHrxn, DH = qP
• Multiply reaction heats
• Reverse reaction heats
• States: H(s) < H(l) < H(g)
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