5.4 First Law of Thermodynamics
First Law of Thermodynamics
Heat & Work
Heat and work are both measured in energy units, so they must both represent energy. How do they differ from each other, and from just plain “energy” itself?
Heat and work are processes and cannot be stored
In our daily language, we often say that “this object contains a lot of heat”, but this kind of talk is a no-no in thermodynamics! It’s ok to say that the object is “hot”, meaning that its temperature is high.
The term “heat” has a special meaning in thermodynamics: it is a process in which a body (the contents of a tea kettle, for example) acquires or loses energy as a direct consequence of its having a different temperature than its surroundings (the rest of the world).
Figure 5.20 Transfer of Heat
Thermal energy can only flow from a higher temperature to a lower temperature. It is this flow that constitutes “heat”.
Use of the term “flow” of heat recalls the 18th-century notion that heat is an actual substance called “caloric” that could flow like a liquid.
Heat is the transfer of energy by conduction or radiation
Transfer of thermal energy can be accomplished by bringing two bodies into physical contact (the kettle on top of the stove, or through an electric heating element inside the kettle). Another mechanism of thermal energy transfer is by radiation; a hot object will convey energy to any body in sight of it via electromagnetic radiation in the infrared part of the spectrum. In many cases, a combination of modes will be active:
Thus when you place a can of beer in the refrigerator, both processes are operative: the can radiates heat to the cold surfaces around it, and absorbs it by direct conduction from the ambient air.
Figure 5.21 Different Methods of Heat Transfer
So what is work?
Work refers to the transfer of energy some means that does not depend on temperature difference.
Figure 5.22 Work Done by system
Work, like energy, can take various forms, the most familiar being mechanical and electrical. Mechanical work arises when an object moves a distance Δx against an opposing force f: w = f Δx N-m; 1 N-m = 1 J.
[image from Ben Wiens Energy site]
Figure 5.23 Work Done by Force
Electrical work is done when a body having a charge q moves through a potential difference ΔV.
Work, like heat, exists only when energy is being transferred.
When two bodies are placed in thermal contact and energy flows from the warmer body to the cooler one,we call the process “heat”. A transfer of energy to or from a system by any means other than heat is called “work”.
Heat can only partially be converted into work
Work can be completely converted into heat (by friction, for example), but heat can only be partially converted to work. Conversion of heat into work is accomplished by means of a heat engine, the most common example of which is an ordinary gasoline engine. The science of thermodynamics developed out of the need to understand the limitations of steam-driven heat engines at the beginning of the Industrial Age. A fundamental law of Nature, the Second Law of Thermodynamics, states that the complete conversion of heat into work is impossible. Something to think about when you purchase fuel for your car!
Figure 5.24 Piston
Thermodynamic View of the World
In thermodynamics, we must be very precise in our use of certain words. The two most important of these are system and surroundings. A thermodynamic system is that part of the world to which we are directing our attention. Everything that is not a part of the system constitutes the surroundings. The system and surroundings are separated by a boundary. If our system is one mole of a gas in a container, then the boundary is simply the inner wall of the container itself. The boundary need not be a physical barrier; for example, if our system is a factory or a forest, then the boundary can be wherever we wish to define it. We can even focus our attention on the dissolved ions in an aqueous solution of a salt, leaving the water molecules as part of the surroundings. The single property that the boundary must have is that it be clearly defined, so we can unambiguously say whether a given part of the world is in our system or in the surroundings.
If matter is not able to pass across the boundary, then the system is said to be closed; otherwise, it is open. A closed system may still exchange energy with the surroundings unless the system is an isolated one, in which case neither matter nor energy can pass across the boundary. The tea in a closed Thermos bottle approximates a closed system over a short time interval.
Properties and the state of a system
The properties of a system are those quantities such as the pressure, volume, temperature, and its composition, which are in principle measurable and capable of assuming definite values. There are of course many properties other than those mentioned above; the density and thermal conductivity are two examples. However, the pressure, volume, and temperature have special significance because they determine the values of all the other properties; they are therefore known as state properties because if their values are known then the system is in a definite state.
Change of state: the meaning of Δ
In dealing with thermodynamics, we must be able to unambiguously define the change in the state of a system when it undergoes some process. This is done by specifying changes in the values of the different state properties using the symbol Δ (delta) as illustrated here for a change in the volume:
ΔV = Vfinal – Vinitial(1-1)
We can compute similar delta-values for changes in P, V, ni (the number of moles of component i), and the other state properties we will meet later.
2 Internal energy and the First Law
Internal energy is simply the totality of all forms of kinetic and potential energy of the system. Thermodynamics makes no distinction between these two forms of energy and it does not assume the existence of atoms and molecules. But since we are studying thermodynamics in the context of chemistry, we can allow ourselves to depart from “pure” thermodynamics enough to point out that the internal energy is the sum of the kinetic energy of motion of the molecules, and the potential energy represented by the chemical bonds between the atoms and any other intermolecular forces that may be operative.
Khan Academy : 1st law of Thermo, 11:26 min
https://www.youtube.com/watch?v=NyOYWFirst Law of Thermodynamics, Basic Introduction – Internal Energy, Heat and Work – Chemistry
How can we know how much internal energy a system possesses? The answer is that we cannot, at least not on an absolute basis; all scales of energy are arbitrary. The best we can do is measure changes in energy. However, we are perfectly free to define zero energy as the energy of the system in some arbitrary reference state, and then say that the internal energy of the system in any other state is the difference between the energies of the system in these two different states.
The first law of thermodynamics states that the energy cannot be created or destroyed under any conditions. Furthermore, the energy is converted from one form to another, i.e. the conversion between heat q, work w and the internal energy U.
The first law of thermodynamics can be summarized in the formula below:
∆U = q + w
Where:
∆U = total internal energy change of the system
q = heat exchanged between the system and its surrounding
w = the work done by or on the system
Furthermore, work w is equal = negative value of the external pressure P on the system X change in volume ∆ V:
w = – p ∆ V
replacing w in the equation of the first law of thermodynamic on obtains:
∆U = q – p ∆ V
The internal energy of the system will:
Decrease: in case the system gives off heat or does work on the surrounding
Increase: in case the system taking in heat from the surrounding or surrounding does work on the system
Also,
If Tsystem > Tsurroundings then the heat flows from the system to the surroundings
q < 0.
This is an exothermic system: Heat is given to the surroundings by the system.
If Tsystem < Tsurroundings then the heat flows from the surroundings to the system
q > 0.
This is an endothermic system: Heat is taken by the system from the surroundings
Since the first law of thermodynamics states that the energy cannot be created of destroyed the, then the change of the internal energy is always equaling zero:
Thus, if the energy is lost by the system, the surrounding then absorbs this lost energy.
If the energy in absorbed onto the system (energy is gained), then the energy is given from the surrounding to the system.
ΔUsystem + ΔUsurroundings = 0
ΔUsystem = – ΔUsurroundings
The table below illustrates some examples of the interaction between the q and w as well as the surrounding:
Example:
NH3 gas is placed in a system with constant pressure. If the surrounding around the system loses 75 J of heat and does 555 J of work onto the system, what is the internal energy of the system?
Using the first law of thermodynamics formula:
∆U = q + w
Since the surrounding loses heat which gained by the system à q is positive
Since the surrounding does work on the system à w is positive
Thus, both q and w are positive:
∆U = 75 J + 555 J = 630 J
Example:
A system has constant volume (ΔV = 0) and the heat around the system increases by 86 J.
a. What is the sign for heat (q) for the system?
b. What is ΔU equal to?
c. What is the value of internal energy of the system in Joules?
Since the system is under constant volume (ΔV = 0), then the work will equal zero:
w = – p ∆ V
since ΔV = 0
w = – p x (0) = 0
and since the surrounding heat is increased by 86 J, then
∆U = q + w = q + 0 = q
a. What is the sign for heat (q) for the system?
q will be negative: Heat is lost by the system and gained by the surrounding. q is less than zero.
b.What is ΔU equal to?
∆U = q + w = q + 0 = q
c. What is the value of internal energy of the system in Joules?
∆U = q + w = q + 0 = q = – 86 J
This video below illustrates some examples as well:
First Law of Thermodynamics, Basic Introduction, Physics Problems