Download Energy: Conservation and Interconversion Demonstration:

1
Energy: Conservation and Interconversion
The energy of the Universe is constant, but can be converted from one form to another
We will be concerned with these three forms of energy:
•
•
•
Chemical energy
Electrical energy
Nuclear energy
The energy liberated in these processes can be in the form of light, charge or current, and heat.
Nuclear energy brings into focus the fact that energy and mass can be interconverted. Although
the interconversion of energy and mass is manifested most clearly in nuclear processes, it is true
for all processes.
This idea places mass determination at the center of energy conversion. Interestingly, concepts
of mass measurement were used in the 18th century by Dalton, Avogadro, and Cannizzaro to
formulate the Atomic Theory of Matter. Even now, mass and its origin are at the center of the
“big” unanswered questions about the Universe that is sometimes called the “theory of
everything”.
Demonstration:
The demonstration on the oxides of nitrogen illustrates a number of important concepts:
At room T:
NO + ½ O2 → NO2 Color change denotes spontaneous reaction. Flask
gets warm. The reaction evolves heat. The origin of color is something we’ll discuss
when we talk about chemical bonding. The electron distribution around nitrogen
changes.
At dry ice/acetone T: NO2 + NO → N2O3 (blue liquid)
N2O3 + ½ O2 → N2O4 (colorless gas)
Observe temperature dependence:
2NO2 = N2O4
At higher temperatures, the gas mixture attains a strong red-orange color. The color
fades at lower temperature. Qualitatively, we know that the reaction as written liberates
heat. LeChatelier’s Principle tells us that to compensate for increased T, the reaction
should run in the direction to absorb the added heat. That is, the reaction runs from right
to left.
2
Two important postulates of thermodynamics will guide our thinking
The First Law is a statement of energy conservation, and provides guidelines for us on what
physical processes can happen
The Second Law is more restrictive, telling us that energy conservation alone is not a sufficient
condition on whether a process occurs or not. The Second Law tells us what processes actually
do happen.
Concepts of Heat and Work
Recall that 1 Joule is the amount of energy required to raise the temperature of 1.00 grams of
liquid water exactly 0.239 °C. A more commonly used unit of heat is the calorie, the amount of
heat required to raise the temperature of 1.00 grams of liquid water by 1.00 °C.
The calorie is clearly a unit of heat, while the Joule, with units of kg (m/sec)2, is obviously a
mechanical unit. Nevertheless, we can write the equivalence
1 cal = 4.184 Joule – The Mechanical Equivalent of Heat
We like to think of heat as a highly degraded form of energy, while work is a more organized
form of energy. A device that converts heat into work is therefore a valuable energy conversion
device. The internal combustion engine is such a device. The efficiency of this conversion is
clearly an important topic – we’ll return to this soon.
Heat, Work, and Energy
To see the distinctions among these terms, let’s divide the universe into the System (the part we
focus on) and the Surroundings (everything else).
We can change the energy of the system by
Transferring heat from the surroundings to the system – heat with a Bunsen burner.
Here, the chemical reaction, the combustion of methane, occurs in the surroundings,
transfers its heat to the system.
Having the surroundings do work on the system – compress the system, transfer electrical
charge (battery charger)
3
The diagrams below help us understand how energy, heat, and work are related. Note that heat
and work are processes or interactions between the system and the surroundings. Heat and work
depend on the way the process is carried out. They “depend on the path”. Note that positive
heat absorbed by the system results in an increase in the internal energy of the system. Similarly,
the surroundings can do work on the system and increase the energy of the system
.
∆E
WORK
HEAT
Important distinction: Internal energy, E, is an intrinsic property of the system. Heat and work
are processes, not intrinsic quantities.
Similarly, when negative heat is absorbed by the system, or negative work is done on the system,
the internal energy decreases.
∆E
WORK
HEAT
These diagrams help us focus on the fact that two processes that involve interactions between
system and surrounding determine a physical property of the system.
The First Law of Thermodynamics says that
∆E = q + w
Significance: q and w are processes that depend on path, but their sum is independent of path
E is a thermodynamic function of state (of the system). So are variables like T, P, V.
4
Functions of state vs. processes
Drive from Rochester to Denver:
Path 1: Take I-90 west to Chicago, through Wisconsin, Minnesota, South Dakota,
Wyoming, then I-25 south to Denver.
Path 2: Take I-90 East to NYC, take I-95 south through New Jersey, Delmarva Peninsula
to Jacksonville, Florida. Take I-10 west along the Gulf Coast (wave to my Uncle in
Gulfport, who is still living in a FEMA trailer), catch I-25 North in Arizona.
Heat absorbed: measured by gasoline consumption
Work done: measured by tire wear
These will be different for the two paths. What is independent of path is the
Change in elevation: ~ 5000 ft
Change in temperature: ??
Change in barometric pressure: ??
These are functions of state – they just depend on the characteristics of the environments of
Rochester and Denver.
5
Pressure-Volume Work
Let’s consider a simple example of the work done on the system (considered to be a gas) when a
constant external pressure is applied to it.
Work done on the system = -(external force × displacement)
External force = external pressure × surface area
W = -P × A × ∆h (note the minus sign)
A × ∆h = ∆V = Vfinal - Vinitial
W = -P∆V = -P(Vfinal - Vinitial)
Let’s do a couple of sample problems: Apply an external pressure of 1.0 atmospheres as a gas
expands from 0.5 L to 1.0 L.
W = -1.0 atm (1.0 L – 0.5 L) = -0.5 L-atm. Not a particularly useful unit. Since pressure is a
force per unit area, the appropriate mechanical SI unit for pressure is N/m2. Turns out that
atmospheric pressure is 1.013 × 105 N/m2. Remember that 1 liter = 1000 cm3 = 10-3 m3. (Do
you know why?). Therefore 1 L-atm = 1.013 × 105 N/m2 × 103 m3 = 101.3 N-m = 101.3 J
This is an important unit conversion.
101.3J
=-50.6 J. Note that negative work is done on the
1L − atm
system. Note also for a compression, the final volume is smaller than the initial volume. So, the
work of compression on the system is positive.
For the expansion W = -0.5 L-atm ×