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.

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)– Energy due to position or

chemical composition

– Equals (mgh) in example.

• Kinetic Energy (KE)– Energy due to motion.

– Equals mv2/2 in example.

h

m

v

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:– PE = 0– KE = 1/2mv2=1/2kx2

• Units of EnergyJoule = kg.m2/s2

• Example:– Init. PE = 10 J– M = 10 kg– Vmax = [2(PE)/M]1/2 = 1.4m/s

0

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.

h

m

v

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.

• System: That part of the universe you are interested in (i.e., you define it).

• Surroundings: The rest of the universe.

System

Surroundings

q transfer

w transfer

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.

P = 1atm

P = 1atm

Initial

Final

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 < 0 (heat is lost)

Ene

rgy

Ene

rgy Einitial

Efinal

Efinal < Einitial

Water @ 80° C

Water @ 20° C

q

Another Example of Exothermic

Heat Exchange: Endothermic

• Endothermic Reaction: Chemical process in which system evolves resulting in heat transfer to the system

• Heat flows to the system

• q > 0 (heat is gained)

Ene

rgy

Ene

rgy

Einitial

Efinal

Efinal > Einitial

Water @ 80° C

Water @ 20° C

q

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

• If system loses energy:

Efinal < Einitial

Efinal-Einitial = E < 0.

• If system gains energy:

Efinal > Einitial

Efinal-Einitial = E > 0.

Ene

rgy

Ene

rgy

Ene

rgy

Einitial

Efinal

E > 0

Ein

Ene

rgy Einitial

Efinal

E < 0

Eout

Heat and Work Sign Convention

• If system gives heat

q < 0 (q is negative)

• If system gets heat

q > 0 (q is positive)

•If system does work

w < 0 (w is negative)

•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• System does work: w < 0

(adiabatic)

Example (cont.)• How much work does the

system do?

• Pext = force/area

• |w| = force x distance

= Pext x A x h

= Pext V

• w = - Pext V (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 E for this process?

• Ans: First, define the system: the balloon.

Example 9.1 (cont.)

E = q + w

= (1.3 x 108 J) + (-PV)

= (1.3 x 108 J) + (-1 atm (VfinalVinit))

= (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.)

E = (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 PV), then:

E = qp + w

= qp - PV

E + PV = qp

qp is heat transferred at constant pressure.

Definition of Enthalpy (cont.)

• Recall: H = E + PV

H = E + PV) = E + PV (P is constant)

= qp

• Or H = 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:

H = Hproducts - Hreactants

If H >0, then qp >0. The reaction is endothermic

If H <0, then qp <0. The reaction is exothermic

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)

Ent

halp

yE

ntha

lpy

Ent

halp

y

Hinitial

Hfinal

H > 0

qin

Ent

halp

y Hinitial

Hfinal

H < 0

qout