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Lecture 1: Energy and Enthalpy

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Title: Lecture 1: Energy and Enthalpy

1
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

2
• 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)

3
Energy Kinetic vs. Potential
• Potential Energy (PE)
• Energy due to position or chemical composition
• Equals (mgh) in example.
• Kinetic Energy (KE)
• Energy due to motion.
• Equals mv2/2 in example.

4
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.

5
Example Mass on a Spring
• Initial PE 1/2 kx2
• At x 0
• PE 0
• KE 1/2mv21/2kx2
• Units of Energy
• Joule kg.m2/s2
• Example
• Init. PE 10 J
• M 10 kg
• Vmax 2(PE)/M1/2 1.4m/s

0
6
Energy Kinetic vs. Potential
• Potential Energy (PE)
• Energy due to position or chemical composition
• Equals (mgh) in example.
• Kinetic Energy (KE)
• Energy due to motion.
• Equals mv2/2 in example.

7
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)

8
Applying the First Law
• Need to differentiate between the system and
surroundings.
• System That part of the universe you are
interested in (i.e., you define it).
• Surroundings The rest of the universe.

9
Conservation of Energy
• Total energy is conserved.
• Energy gained by the system must be lost by the
surroundings.
• Energy exchange can be in the form of q, w, or
both.

10
Heat Exchange Exothermic
• Exothermic Reaction. Chemical process in which
system evolves resulting in heat transfer to the
surroundings
• Heat flows out of the system
• q lt 0 (heat is lost)

11
Another Example of Exothermic
12
Heat Exchange Endothermic
• Endothermic Reaction Chemical process in which
system evolves resulting in heat transfer to the
system
• Heat flows to the system
• q gt 0 (heat is gained)

13
Another Example of Endothermic
14
• 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)

15
Energy and Sign Convention
• If system loses energy
• Efinal lt Einitial
• Efinal-Einitial DE lt 0.
• If system gains energy
• Efinal gt Einitial
• Efinal-Einitial DE gt 0.

16
Heat and Work Sign Convention
• If system gives heat
• q lt 0 (q is negative)
• If system gets heat
• q gt 0 (q is positive)
• If system does work
• w lt 0 (w is negative)
• If work done on system
• w gt 0 (w is positive)

17
Example Piston
• Figure 9.4, expansion against a constant external
pressure
• No heat exchange
• q 0
• System does work
• w lt 0

18
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)

19
• When it is compressed, work is done to a gas
• When it is expanded, work is done by the gas

20
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.

21
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

22
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.

23
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

24
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.

25
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.

26
Changes in Enthalpy
• Consider the following expression for a chemical
process
• DH Hproducts -
Hreactants
• If DH gt0, then qp gt0. The reaction is
endothermic
• If DH lt0, then qp lt0. The reaction is
exothermic

27
Enthalpy Changes Pictorially
• Similar to previous discussion for Energy.
• Heat comes out of system, enthalpy decreases (ex.
Cooling water).
• Heat goes in, enthalpy increases (ex. Heating
water)