Lecture 23 Highlights…
ÆStart Thermodynamics
Æ Heat Flow and Work for Systems /
Surroundings
Æ Internal Energy
Æ Enthalpy
Thermodynamics
9Definition: The study of energy and
its transformations.
9To Begin : Review Dr. Teece’s
discussion of Energy, Spontaneity,
Free Energy, Enthalpy and Entropy!!
9Energy of phase transitions
9Reaction energies
• Calorimetry
Chemistry
The Science in Context
Chapter 11 and 13
1
A Thermodynamic System…..
The “system”:
gas + oxygen
→ CO2, water +
heat
The “surroundings”:
Food, pot, kitchen,
etc…..
First Law of Thermodynamics: Energy gained or lost by
the “system” must equal the energy gained or lost by its
“surroundings”…… Energy is “conserved” in a process
Some Definitions…..
: Exothermic: If a chemical reaction or physical
change results in heat flow from a system to its
surroundings, its exothermic
: Endothermic: Reactions and physical changes that
absorb heat from their surroundings are
endothermic.
: Example: H2O(l) →H2O(g) ?
: q is heat flow: By definition, if q < 0, reaction /
change is exothermic, and if q> 0, its endothermic.
Require input of heat
11_06.jpg
q is positive
2
Mechanical Energy Model
Potential vs. Kinetic Energy
Potential Energy (PE)
9It’s a “State Function.”
9PE does not depend on path, just initial
and final state.
9The work to climb the hill can depend
on pathway (straight up on chair-lift) or
a winding path.
9Mechanical work is a function of gravity
and distance.
Click on Image
9PE is “stored up” work.
Energy Stored in Chemical Bonds
11_07.jpg
3
Internal Energy
9 Let’s call a gas trapped in a vessel a “Chemical
System”.
9 Internal energy of a system = Sum of kinetic
and potential energies = E
9 Change in Internal Energy going from state A to
state B is ΔE = q + w, where q = heat flow into
or out of the system and w = work done on or
by the system. (Unit of work is the joule (J)).
This is the First Law of Thermodynamics:
Click on Image
Isothermal Expansion of Ideal Gas
P = 1 atm
An ideal gas in a sealed piston is
allowed to expand against a
pressure of 1 atm. In what
direction, if at all, does heat flow
for this process?
A) into the system
B) out of the system
C) heat does not flow
How Did Internal Energy Change ?
(How do we get ΔE ?)
• ΔE = q + w
• Heat is absorbed by the system from the
surroundings (q)
• The ideal gas expands (its volume increases) against
a constant atmospheric pressure so w = PΔV (this is
the definition of work)
• ΔE = heat absorbed by system + work done by the
system.
• ΔE gained or lost by system = ΔE gained or lost by
the surroundings
4
Here’s the Confusing Part ……
• If work is done by a system through expansion (like
the previous position example) then….
• ΔE = q + w = q - PΔV
• Why is there a negative sign in front of PΔV ?
• Answer: If the system expands (ΔV is positive),
then it loses internal energy as it does work on its
surroundings.
• q is positive since the system absorbs energy from
its surroundings to do work…(the system is
endothermic).
Confusing Part II ……
• What if work is done on a system, say compressing a
gas in a piston, then ΔE = q + w = q + PΔV
• So, the sign of the work term (w = PΔV) is
dependent on what’s happening with the system.
• Sorry, this issue has plagued scientists and
engineers since the dawn of time……. You are not the
first group of students to say “Huh ???”
• If you are solving a problem like Chapter 11, Problem
29, apply ΔE = q + w, and use whatever signs of q
and w that are given in the problem. (I think they
used the wrong signs in Problem 31 !!)
Summarizing the Confusion ……
• q > 0 if heat is transferred from the surroundings
to the system (the system is endothermic)
• q < 0 if heat is transferred from the system to the
surroundings (the system is exothermic)
• w > 0 if work is done by the surroundings on the
system . (Compression)
• w < 0 if work is done by the system on the
surroundings. (Expansion)
Remember
These !
• Therefore: The value of ΔE depends on the signs
and magnitude of q and w.
5
Internal Energy and Enthalpy
9 Chemists like to do experiments where they can
make measurements easily (We’re lazy and have
little money for complicated equipment !)
9 Measuring heat flow (q) is easy to do because all you
need to measure is temperature change (ΔT) if you
know something about the composition of the system
(Heat Capacity …More on this in future lectures)
9 For reactions or processes where there is no volume
change (PΔV = 0), the internal energy change (ΔE) is
equal to heat flow into or out of the system (qv).
9 So, at constant volume, heat flow is a direct measure
of internal energy change (qv).
Internal Energy and Enthalpy II
9 Chemists like to do experiments where they can
make measurements easily (I repeat, “We’re lazy and
have little money for complicated equipment !”)
9 Lazy chemists prefer to work out in the open at
constant pressure (1 atm) rather than design a
vessel that has constant volume. This gave birth to a
new function: H = Enthalpy
Internal Energy and Enthalpy
9 For a chemical system at constant pressure,
a new “State” function is defined where: H
= E +PV, where H = Enthalpy
9 If this system experiences a change (say
work is done on it or by it), then ΔH = ΔE
+ PΔV.
9 However, since ΔE = q +w = q - PΔV.
9 Then ΔH = q -PΔV + PΔV.
9 Then, at constant pressure, ΔH = qp, where
qp = heat flow at constant pressure.
6
Internal Energy / Enthalpy
Fixed
Volume
Heat Flow
Open to Heat Flow
Atm Pressure
Internal
Energy
∆E = qv
Enthaply
∆H = qp
Changes in Enthalpy (∆E)
9 Chemical reactions and physical changes that
occur at constant pressure (say, atmospheric
pressure) absorb heat from their
surroundings and they are endothermic.
9 ΔH (= qp >0) is positive for endothermic
reactions.
9 Examples:
• Ice cubes melting on a hot day
• Perspiration evaporating from skin
• Dissolving certain compounds in liquids
Exothermic Reactions
9Chemical reactions and physical changes
that release heat (energy) to their
surroundings at constant pressure are
exothermic.
9ΔH is negative (= qp <0)
9Examples:
• Fireworks
• Combustion in a car engine.
7
Require input of heat
ΔH positive11_06.jpg
ΔH negative
Chemical Thermodynamics
9Will a chemical reaction proceed?
9Does a reaction occur spontaneously
in the forward direction? Such as:
H2O(l) → H2O(g)
9This depends both on enthalpy and
entropy.
8