Download Lecture01.f

Lecture 1: Energy and Enthalpy
• Reading: Zumdahl 9.1 and 9.2
• Outline
–
–
–
–
Energy: Kinetic and Potential
System vs. Surroundings
Heat, Work, and Energy
Enthalpy
• Energy is the capacity to do work or to
produce heat
• Energy is conserved, it can neither be
created nor destroyed, different forms of
energy interconvert
• However, the capacity to utilize energy to
do work is limited (entropy)
Energy: Kinetic vs. Potential
• Potential Energy (PE)
m
– Energy due to position or
chemical composition
– Equals (mgh) in example.
h
v
• Kinetic Energy (KE)
– Energy due to motion.
– Equals mv2/2 in example.
Mechanical Energy = KE + PE
• Energy is the sum of kinetic energy and
potential energy.
• Energy is readily interconverted between
these two forms.
• If the system of interest is isolated (no
exchange with surroundings), then total
energy is constant.
Example: Mass on a Spring
• Initial PE = 1/2 kx2
• At x = 0:
0
– PE = 0
– KE = 1/2mv2=1/2kx2
• Units of Energy
Joule = kg.m2/s2
• Example:
– Init. PE = 10 J
– M = 10 kg
– Vmax = [2(PE)/M]1/2 = 1.4m/s
Energy: Kinetic vs. Potential
• Potential Energy (PE)
m
– Energy due to position or
chemical composition
– Equals (mgh) in example.
h
v
• Kinetic Energy (KE)
– Energy due to motion.
– Equals mv2/2 in example.
First Law of Thermodynamics
First Law: Energy of the Universe is Constant
E=q+w
q = heat. Transferred between two bodies of
differing temperature. Note: q ≠ Temp!
w = work. Force acting over a distance (F x d)
Applying the First Law
• Need to differentiate
between the system and
surroundings.
Surroundings
System
• System: That part of the
universe you are interested in
(i.e., you define it).
• Surroundings: The rest of
the universe.
q transfer
w transfer
Conservation of Energy
• Total energy is conserved.
P = 1atm
• Energy gained by the
system must be lost by the
surroundings.
Initial
P = 1atm
• Energy exchange can be in
the form of q, w, or both.
Final
Heat Exchange: Exothermic
• Exothermic Reaction.
Chemical process in which
system evolves resulting in
heat transfer to the
surroundings
Energy
Water @ 80° C
Einitial
Water @ 20° C
q
Efinal
• Heat flows out of the system
Efinal < Einitial
• q < 0 (heat is lost)
Another Example of Exothermic
Heat Exchange: Endothermic
Energy
Water @ 80° C
Efinal
Water @ 20° C
• Endothermic Reaction:
Chemical process in
which system evolves
resulting in heat transfer
to the system
q
Einitial
• Heat flows to the system
Efinal > Einitial
• q > 0 (heat is gained)
Another Example of Endothermic
• In exothermic reactions, the potential
energy stored in chemical bonds is
converted into thermal energy (random
kinetic energy), i.e. heat
• Once we have done that, we have lost the
ability to utilize the same potential energy
to do work or generate heat again
(dissipation)
Energy and Sign Convention

Energy
Einitial
Eout
Efinal
DE < 0
Energy
Efinal
Einitial
DE > 0
Ein
• If system loses energy:
Efinal < Einitial
Efinal-Einitial = DE < 0.
• If system gains energy:
Efinal > Einitial
Efinal-Einitial = DE > 0.
Heat and Work Sign Convention
• If system gives heat
q < 0 (q is negative)
•If system does work
w < 0 (w is negative)
• If system gets heat
q > 0 (q is positive)
•If work done on system
w > 0 (w is positive)
Example: Piston
• Figure 9.4, expansion
against a constant
external pressure
• No heat exchange:
q = 0 (adiabatic)
• System does work:
w<0
Example (cont.)
• How much work does the
system do?
• Pext = force/area
• |w| = force x distance
= Pext x A x Dh
= Pext DV
• w = - Pext DV (note sign)
• When it is compressed, work is done to a
gas
• When it is expanded, work is done by the
gas (e.g. your car’s engine)
Example 9.1
• A balloon is inflated from 4 x 106 l to 4.5 x
106 l by the addition of 1.3 x 108 J of heat.
If the balloon expands against an external
pressure of 1 atm, what is DE for this
process?
• Ans: First, define the system: the balloon.
Example 9.1 (cont.)
DE = q + w
= (1.3 x 108 J) + (-PDV)
= (1.3 x 108 J) + (-1 atm (Vfinal - Vinit))
= (1.3 x 108 J) + (-0.5 x 106 l.atm)
• Conversion: 101.3 J per l x atm
(-0.5 x 106 l.atm) x (101.3 J/l.atm) = -5.1 x 107 J
Example 9.1 (cont.)
DE = (1.3 x 108 J) + (-5.1 x 107 J)
= 8 x 107 J (Ans.)
The system gained more energy through heat
than it lost doing work. Therefore, the
overall energy of the system has increased.
Definition of Enthalpy
• Thermodynamic Definition of Enthalpy (H):
H = E + PV
E = energy of the system
P = pressure of the system
V = volume of the system
Why we need Enthalpy?
• Consider a process carried out at constant
pressure.
• If work is of the form D(PV), then:
DE = qp + w
= qp - PDV
DE + PDV = qp
qp is heat transferred at constant pressure.
Definition of Enthalpy (cont.)
• Recall: H = E + PV
DH = DE + D(PV)
= DE + PDV (P is constant)
= qp
• Or DH = qp
• The change in enthalpy is equal to the heat
transferred at constant pressure.
Changes in Enthalpy
• Consider the following expression for a chemical
process:
DH = Hproducts - Hreactants
If DH >0, then qp >0. The reaction is endothermic
If DH <0, then qp <0. The reaction is exothermic
Enthalpy Changes Pictorially
Enthalpy

Hinitial
q out
• Similar to previous discussion
for Energy.
Hfinal
• Heat comes out of system,
enthalpy decreases (ex. Cooling
water).
Enthalpy
DH < 0
Hfinal
Hinitial
DH > 0
q in
• Heat goes in, enthalpy increases
(ex. Heating water)