Energy, Heat and Work

Energy, Heat & Work

• What is Energy?

• Forms of Energy

• Energy Changes

• Energy Transfer (Heat and Work)

• Heat Transfer

• Forms of Work

• The First Law of Thermodynamics (Formal statement/Energy Balance)

 What is Energy?
It is extremely difficult to give a proper definition of “energy”
The term “energy” was first coined by Thomas Young in 1805, and
its use in Thermodynamics was proposed by Lord Kelvin in 1852.
Energy can be stored within a system and it can be transferred.
We are all familiar with the “conservation of energy” principle. This
principle is an expression of the first law of thermodynamics. In
simple terms the principle sates that “energy cannot be created or
destroyed during a process”; it can only change from one form to
another.
Let us test our understanding of this principle. Consider a room
whose door and windows are tightly shut and whose walls are well
insulated (no heat loss). Let us place refrigerator in the middle of
the room with its door open and plug it into an electric socket.
 What is Energy?

The system is the entire room

Well sealed – Well insulated
Is it a closed system?
What is the energy interaction?

What will happen to the average temperature of

the room? Will it increase? Will it decrease?
Will it remain the same?

In which direction is energy moving?

What does the conservation law tell us?
What happens to this energy when it enters the room?
Is it an Isothermal process?
 What is Energy?
Energy is conserved in our example of a refrigerator operating in a
closed room:
The system is an adiabatic closed system
The electric energy is the only form of energy crossing into the room
This electric energy is converted into an equivalent amount of thermal energy
This thermal energy is now stored in the room air
This results in a rise in the air temperature.
When we talk about conservation of energy we usually refer to
the “quality” of energy not quantity:
Electric energy (highest quality of energy) can be converted to an equal
amount of thermal energy (heat).
Only a small fraction of thermal energy (lowest form of energy) can be
converted back to electricity (more of this later).

What do you think would happen if we placed a fan instead of a

refrigerator?
What is the SI unit for energy? Can we derive it from first
principles?
 Forms of Energy
Energy can exist in numerous forms: thermal, mechanical,
potential, kinetic, electric, magnetic, chemical and nuclear.

The sum of these “energies” constitutes the “total energy”, E, for

a system.

It is often helpful to consider the various forms of energy that

make up the total energy of a system in two groups: macroscopic
and microscopic.
Macroscopic forms of energy are those a system possesses as a
whole with respect to some outside reference frame: Kinetic and
potential energies are examples.
Microscopic forms of energy are those related to the molecular
structure of a system and the degree of molecular activity and
are independent of outside reference frames.
 Forms of Energy
The macroscopic energy of a system is related to motion and the
influence of some external effects such as gravity, magnetism,
electricity and surface tension.
The macroscopic energy a system possesses as a result of its
motion is called Kinetic Energy (KE).
The macroscopic energy a system possesses as a result of its
elevation in a gravitational field is called Potential Energy (PE).
The magnetic, electrical and surface tension forms of macroscopic
energy are significant in specialized cases only and are usually
ignored.
The sum of all Microscopic forms of energy is called the “internal
energy” of a system and is denoted by U.
In the absence of magnetic, electrical and surface tension effects,
the total energy of a system, E, at a given state, consists of the
sum of the kinetic, potential and internal energies, i.e.:

E = U + KE + PE
 Forms of Energy: Internal Energy
Internal energy is defined earlier as the sum of all the microscopic
forms of energy of a system. It is related to the molecular structure
and the degree of molecular activity and can be viewed as the sum
of the kinetic and potential energies of the molecules.
Microscopic forms of Energy.
Sensible Energy (kinetic energy of molecules and electrons such as spin,
translational and rotational energies). Note that average velocity and degree
of activity increase with temperature.
Latent Energy (associated with phase change, energy is needed to break the
binding forces between molecules). Note that gases have higher internal
energies than liquids which in turn have higher energies than solids (you
need energy to break bonds between solids)
Chemical Energy (associated with atomic bonds in a molecule where during
a chemical reaction some bonds are destroyed and others are formed)
Nuclear energy (associated with the strong bonds within the nucleus of
the atom)
Forms of Energy: Internal Energy

Latent and sensible Energy Chemical Energy

Nuclear Energy
 Forms of Energy
Macroscopic kinetic energy of an object as
a whole is different from the microscopic
kinetic energies of its molecules.
The kinetic energy of an object is an
organized form of energy associated with
the orderly motion of all molecules in one
direction in a straight path.
Kinetic energies of the molecules are
completely random and highly
disorganized.

Organized energy can be converted to

disorganized energy completely, but only a
fraction of disorganized energy can be
converted to organized energy by specially
built devices called heat engines (car
engines and power plants).
 Energy Changes
The forms of energy already discussed, which constitute the total
energy of a system, can be contained or stored in a system and
can be viewed as the static forms of energy.
Thermodynamics provides no information about the “absolute”
value of the total stored or static energy. It deals only with the
“change” of the total energy, ∆E.
The evaluation of the energy change of a system during a
process involves the evaluation of the energy of the system at
the beginning and at the end of a process and taking their
difference. That is:
Energy Change (∆E) = Energy at final state ̶ Energy at initial state
= Efinal ̶ Einitial = E2 ̶ E1
Using our earlier definition of total energy the change in total
energy will be given by:

E  KE  PE  U
 Energy Transfer

The forms of energy not stored in a system

can be viewed as the dynamic forms of energy
or as energy interactions or transfers.

The dynamic forms of energy are recognized

at the system boundary as they cross it, and
they represent the energy gained or lost by a
system during a process.

The only two forms of energy interactions/

transfers from and to a closed system are heat
transfer and work.

An energy interaction/transfer is considered as

heat transfer if its driving force is a
temperature difference.

Otherwise it is work.
 Energy Transfer

• Work: W (in J)
– W > 0 : Work done by the system
– W < 0 : Work done on the system

– Time rate of work is Power: W
(in J/s or Watts, W)

• Heat Transfer: Q (in J)

– Q > 0 : Heat transfer into the system
– Q < 0 : Heat transfer out of the system

– Rate of heat transfer Q
(in J/s or Watts, W):
 Energy Transfer
Heat and Work are directional
quantities and thus the complete
description of heat or work requires
the specification of both magnitude
and direction
Heat and Work are recognized at the
boundaries of a system as they cross the
boundaries; i.e. they are boundary
phenomena.

Systems posses energy, but not heat and

work.

Heat and Work are associated with a

process, not a state and are not properties.
Heat and Work are path functions, i.e. their
magnitudes depend on the path followed
during a process as well as the end states.
 Energy Transfer
Path functions have inexact differentials designated by the symbol δ. A
differential amount of heat or work is represented by δQ or δW, respectively,
instead of dQ or dW.

Properties, however, are point functions (i.e., they depend on the state only,
and not on how a system reaches that state), and they have exact differentials
designated by the symbol d.
A small change in volume, for example, is
represented by dV, and the total volume change
during a process between states 1 and 2 is:
That is the volume change during process 1–2 is always the volume at state 2
minus the volume at state 1, regardless of the path followed.

The total work done during process 1-2,

however, is

The total work is obtained by following the process path and adding the
differential amounts of work (δW) done along the way. Since work is not a
property, the integral is not equal to W2- W1
 Heat Transfer
Heat is defined as the form of energy that is transferred between
two systems (or a system and its surroundings) by virtue of a
temperature difference.
Heat is energy in transition where it is recognized only as it crosses the
boundary of a system.

A process during which there is no heat

transfer is called an adiabatic process.
The word adiabatic comes from the
Greek word adiabatos, which means not
to be passed.
There are two ways a process can be
adiabatic:
The system is well insulated so that only a
negligible amount of heat can pass
through the boundary.
Both the system and the surroundings are
at the same temperature, thus there is no
driving force (temperature difference) for
heat transfer.
 Heat Transfer
An adiabatic process should not be confused with an isothermal process.
Even though there is no heat transfer during an adiabatic process, the
energy content and thus the temperature of a system can still be changed by
other means such as work.
It was only in the middle of the nineteenth century that we had a true
physical understanding of the nature of heat, thanks to the development at
that time of the kinetic theory, which treats molecules as tiny balls that are
in motion and thus possess kinetic energy. Heat is then defined as the
energy associated with the random motion of atoms and molecules.
The prevailing view of heat until the middle of the
nineteenth century was based on the caloric theory
proposed by the French chemist Antoine Lavoisier in
1789. The caloric theory asserts that heat is a fluid-
like substance called the caloric that is a massless,
colorless, odorless, and tasteless substance that can be
poured from one body into another. When caloric was
added to a body, its temperature increased; and when
caloric was removed from a body, its temperature
decreased. When a body could not contain any more
caloric, much the same way as when a glass of water
could not dissolve any more salt or sugar, the body
was said to be saturated with caloric.
 Heat Transfer
The caloric theory came under attack soon after its introduction. It
maintained that heat is a substance that could not be created or destroyed.
Yet it was known that heat can be generated indefinitely by rubbing one’s
hands together or rubbing two pieces of wood together.
In 1798, the American Benjamin Thompson (Count Rumford) showed in
his papers that heat can be generated continuously through friction. The
validity of the caloric theory was also challenged by several others. But it
was the careful experiments of the Englishman James P. Joule, published
in 1843, that finally convinced the skeptics that heat was not a substance
after all, and thus put the caloric theory to rest.
Although the caloric theory was totally abandoned in the middle of the
nineteenth century, it contributed greatly to the development of
thermodynamics and heat transfer (the calorie, as a unit for energy and
saturated liquid and saturated vapour are terms still in use today).
Heat is transferred by three mechanisms: conduction, convection, and
radiation. Conduction is the transfer of energy from the more energetic
particles of a substance to the adjacent less energetic ones as a result of
interaction between particles. Convection is the transfer of energy between a
solid surface and the adjacent fluid that is in motion, and it involves the
combined effects of conduction and fluid motion. Radiation is the transfer of
energy due to the emission of electromagnetic waves (or photons).
 Heat Transfer
 dT
Conduction Q x   A
dx

Convection Q c  hA(Tb  T f )


Radiation Q e   ATb4

Photo courtesy of Mike Benson

 Heat Transfer
• Conduction is the transfer of energy from the more energetic particles of a
substance to the adjacent less energetic ones as a result of interactions
between the particles.
• Conduction can take place in solids, liquids, or gases.
• In gases and liquids, conduction is due to the collisions of the molecules
during their random motion.
• In solids, it is due to the combination of vibrations of molecules in a lattice
and the energy transport by free electrons.
• Heat transfer by conduction is described by the following equation:

• This is known as Fourier’s law of heat conduction.

• It indicates that the rate of heat conduction in a direction is proportional to
the temperature gradient in that direction.
• Heat is conducted in the direction of decreasing temperature, and the
temperature gradient becomes negative when temperature decreases with
increasing x. Therefore, a negative sign is added in the equation above to
make heat transfer in the positive x direction a positive quantity.
• The constant of proportionality kf is the thermal conductivity of the
material, which is a measure of the ability of a material to conduct heat.
Heat Transfer
 Heat Transfer
• Convection is the mode of energy transfer between a solid surface and the
adjacent liquid or gas that is in motion, and it involves the combined effects
of conduction and fluid motion. The faster the fluid motion, the greater the
convection heat transfer.
• In the absence of any bulk fluid motion, heat transfer between a solid
surface and the adjacent fluid is by pure conduction. The presence of bulk
motion of the fluid enhances the heat transfer between the solid surface and
the fluid, but it also complicates the determination of heat transfer rates.
• Consider the cooling of a hot block by blowing of cool air over its top
surface. Energy is first transferred to the air layer adjacent to the surface of
the block by conduction. This energy is then carried away from the surface
by convection; that is, by the combined effects of conduction within the air,
which is due to random motion of air molecules, and the bulk or
macroscopic motion of the air, which removes the heated air near the
surface and replaces it by the cooler air.
• Convection is called forced convection if the fluid is forced to flow in a
tube or over a surface by external means such as a fan, pump, or the wind.
In contrast, convection is called free (or natural) convection if the fluid
motion is caused by buoyancy forces induced by density differences due to
the variation of temperature in the fluid.
• The rate of heat transfer by convection is determined from Newton’s law
of cooling, expressed as
 Heat Transfer

• where h is the convection heat transfer coefficient, A is the surface area

through which heat transfer takes place, Ts is the surface temperature, and
Tf is bulk fluid temperature away from the surface. (At the surface, the
fluid temperature equals the surface temperature of the solid.)
• The convection heat transfer coefficient h is not a property of the fluid. It is
an experimentally determined parameter whose value depends on all the
variables that influence convection such as the surface geometry, the nature
of fluid motion, the properties of the fluid, and the bulk fluid velocity.
• Heat transfer processes that involve change of phase of a fluid are also
considered to be convection because of the fluid motion induced during the
process such as the rise of the vapor bubbles during boiling or the fall of
the liquid droplets during condensation.
• Typical values of h, in W/m2 · K, are in the range of:
– 2–25 for the free convection of gases
– 50–1000 for the free convection of liquids,
– 25–250 for the forced convection of gases,
– 50–20,000 for the forced convection of liquids,
– 2,500–100,000 for convection in boiling and condensation processes.
 Heat Transfer
• Radiation is the energy emitted by matter in the form of electromagnetic waves
(or photons) as a result of the changes in the electronic configurations of the
atoms or molecules. Unlike conduction and convection, the transfer of energy
by radiation does not require the presence of an intervening medium.
• Energy transfer by radiation is fastest (at the speed of light) and it suffers no
attenuation in a vacuum. We are interested in thermal radiation, which is the
form of radiation emitted by bodies because of their temperature. It differs from
other forms of electromagnetic radiation such as X-rays, gamma rays, micro-
waves, etc. that are not related to temperature. All bodies at a temperature above
absolute zero emit thermal radiation.
• Radiation is a volumetric phenomenon, and all solids, liquids, and gases emit,
absorb, or transmit radiation of varying degrees. However, radiation is usually
considered to be a surface phenomenon for solids that are opaque to thermal
radiation such as metals, wood, and rocks since the radiation emitted by the
interior regions of such material can never reach the surface, and the radiation
incident on such bodies is usually absorbed within a few microns from the
surface.
• The radiation emitted by all real surfaces is less than the radiation emitted by a
blackbody at the same temperatures and is expressed by the Stefan–Boltzmann
law as:
 Heat Transfer

where A is the surface area

and ζ = 5.67 10-8 W/m2 · K4 is
the Stefan–Boltzmann
constant and ε is the
emissivity of the surface. The
property emissivity, whose
value is in the range 0≤ε≤ 1, is
a measure of how closely a
surface approximates a
blackbody for which ε=1
 Forms of Work
As mentioned earlier, energy can cross the boundary of a
closed system in the form of heat or work. Therefore, if the
energy crossing the boundary of a closed system is not heat, it
must be work.
The term “work” was first used in the scientific scene by
Coriolis in 1829.
A more formal, thermodynamic definition of work states: Work
is done by a system on its surroundings if the sole effect on
everything external to the system could have been the raising
of a weight.
In other words, work is the energy transfer (across a boundary)
associated with a force acting through a distance resulting
(usually) in a movement of the boundary.
Can we derive the units for work from primary units?
There are several forms of work.
 Forms of Work
V2
• Expansion/Compression Work
(Moving Boundary Work)
 p dV
V1

• Elongation of a solid bar

• Stretching of a Liquid Film
• Rotating Shaft
• Electric
• Others:
•Polarization
•Magnetization
•Surface tension
•Spring work
 The 1st Law: Conservation of Energy
Most of the credit for discovery and confirmation of the first law as a
generalization of the conservation of energy goes to two German doctors
(physicians): Hermann von Helmholtz and Julius Mayer and to an English
scientist: James Prescott Joule.
Statements of the first law of thermodynamics:
Energy may be changed from one form to another but is neither
created nor destroyed.
For all adiabatic processes between two specified states of a closed
system, the net work done is the same regardless of the nature of the
closed system and the details of the process.
Let us return to our earlier definition of the energy change of a system. Now
let us suppose that some thermodynamic system undergoes a process in
which it is changed in several ways but in the end returns to its original state
or condition. The first law tells us that the total energy of the system must be
conserved, meaning that the energy of the system is the same at the end of
the cyclical process as it was at the beginning. This can be expressed as:
E2 = E1 or E2-E1 = ∆E = 0
From the italicized statement of the first law given above we may state that the energy
change undergone by a system at the end of a non-cyclical process is equal to the
energy gained in the form of heat minus the energy lost as a result of work done on
the surroundings. This statement is expressed concisely by:
 The 1st Law: Conservation of Energy

∆E = ∆U+∆KE +∆PE = Q - W
Change in Net amount of Net amount of
amount of energy energy
energy contained transferred in transferred out
within the = across the -
across the
system during system boundary system boundary
some time by heat transfer by work during
interval during the time the time interval
interval
As discussed previously, the changes in potential and kinetic energies for a closed
stationary system are zero, i.e. both ∆KE and ∆PE are = 0. Thus, the above equation
reduces to:

∆E = ∆U = Q - W
 The 1st Law: Conservation of Energy
• Note that Q is the heat absorbed by the system and W is the
work done on the surroundings by the system.
• Because Q represents the heat absorbed by the system, +Q
means heat absorbed and -Q means heat evolved or lost by the
system. If we have a flame under a boiler or other body being
heated, Q is negative for the flame system which is losing heat
to its surroundings and Q is positive for the boiler system
which is absorbing heat from its surroundings.
• Because W represents work done by the system on the
surroundings, +W means that the system accomplishes work
on its surroundings while -W means that work is done on the
system by the surroundings.
• The first law, as summarized by ∆E = 0 for cyclical processes
and ∆E = Q - W for all processes, gives us a sound basis for
useful application of accounting procedures to all phenomena
involving heat and work.
 The 1st Law: Conservation of Energy
In the following paragraph we give a simple illustration of this
aspect of the first law, which might well be called the
"Bookkeeper's Delight" because the accounts always balance.
Suppose that a closed system absorbs 120 kJ from one part of its
surroundings and does 150 kJ of work on another part of its
surroundings. What is ∆E for the system? Since heat absorbed
corresponds to positive Q and work done on the surroundings
also corresponds to positive W, we have

∆E = Q - W = (+120) - (+150) = -30 kJ

Remembering that ∆E is shorthand for Efinal – Einitial or E2 - Е1

we also have E2 - E1 = -30 kJ. The energy of our system has
diminished by 30 KJ and has lost some of its capability for
doing work. This decreased capability for doing work might
appear in the form of a lower temperature for the system.