Thermal Physics – Phase Change, Heat Transfer, Introduction to Thermodynamics
Changes of Phase
There are four states of matter in the universe: plasma, gas, liquid and solid. But, matter on
Earth exists mostly in three distinct phases: gas, liquid and solid. A phase is a distinctive form of
a substance, and matter can change among the phases. It may take extreme temperature, pressure
or energy, but all matter can be changed.
There are six distinct changes of phase which happens to different substances at different
temperatures. The six changes are:
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Freezing: the substance changes from a liquid to a solid.
Melting: the substance changes back from the solid to the liquid.
Condensation: the substance changes from a gas to a liquid.
Vaporization: the substance changes from a liquid to a gas.
Sublimation: the substance changes directly from a solid to a gas without going through
the liquid phase.
Deposition: the substance changes directly from a gas to a solid without going through the
liquid phase.
Energy is required to melt a solid because the cohesive bonds between the molecules in the solid
must be broken apart such that, in the liquid, the molecules can move around at comparable kinetic
energies; thus, there is no rise in temperature. Similarly, energy is needed to vaporize a liquid,
because molecules in a liquid interact with each other via attractive forces. There is no temperature
change until a phase change is complete. The temperature of a cup of soda initially at 0 stays at 0
until all the ice has melted. Conversely, energy is released during freezing and condensation,
usually in the form of thermal energy. Work is done by cohesive forces when molecules are
brought together. The corresponding energy must be given off (dissipated) to allow them to stay
together.
The energy involved in a phase change depends on two major factors: the number and strength of
bonds or force pairs. The number of bonds is proportional to the number of molecules and thus to
the mass of the sample. The strength of forces depends on the type of molecules. The heat required
to change the phase of a sample of mass is given by
where the latent heat of fusion, and latent heat of vaporization, are material constants that are
determined experimentally.
(a) Energy is required to partially overcome the attractive forces between molecules in a solid
to form a liquid. That same energy must be removed for freezing to take place.
(b) Molecules are separated by large distances when going from liquid to vapor, requiring
significant energy to overcome molecular attraction. The same energy must be removed
for condensation to take place. There is no temperature change until a phase change is
complete.
Latent heat is measured in units of J/kg. Both Lf and Lv depend on the substance, particularly
on the strength of its molecular forces as noted earlier. Lf and Lv are collectively called latent
heat coefficients. They are latent, or hidden, because in phase changes, energy enters or leaves
a system without causing a temperature change in the system; so, in effect, the energy is hidden.
Phase Change Diagram
Example:
1. The specific heat for water is 4186 J/kg oC. The specific heat for ice is 2100 J/kg oC and
the specific heat for steam (water vapor) is 2010 J/kg oC. So how much heat is required
to raise a 1 kg block of ice from -40 oC to 120 oC?
2. Three ice cubes are used to chill a soda at 20C with mass of 0.25 kg. The ice is at 0C and
each ice cube has a mass of 6.0 g. Assume that the soda is kept in a foam container so
that heat loss can be ignored. Assume the soda has the same heat capacity as water. Find
the final temperature when all ice has melted.
3. What mass of steam at 130 °C must be condensed onto a 0.100 kg glass cup to warm the
cup and the 0.200 kg of water it contains from 20.0 °C to 50.0 °C?
Heat Transfer
Equally as interesting as the effects of heat transfer on a system are the methods by which
this occurs. Whenever there is a temperature difference, heat transfer occurs. Heat transfer may
occur rapidly, such as through a cooking pan, or slowly, such as through the walls of a picnic ice
chest. We can control rates of heat transfer by choosing materials (such as thick wool clothing for
the winter), controlling air movement (such as the use of weather stripping around doors), or by
choice of color (such as a white roof to reflect summer sunlight). So many processes involve heat
transfer, so that it is hard to imagine a situation where no heat transfer occurs. Yet every process
involving heat transfer takes place by only three methods:
I.
Conduction is heat transfer through stationary matter by physical contact. (The matter
is stationary on a macroscopic scale—we know there is thermal motion of the atoms
and molecules at any temperature above absolute zero.) Heat transferred between the
electric burner of a stove and the bottom of a pan is transferred by conduction.
II.
Convection is the heat transfer by the macroscopic movement of a fluid. This type of
transfer takes place in a forced-air furnace and in weather systems, for example.
III.
Heat transfer by radiation occurs when microwaves, infrared radiation, visible light, or
another form of electromagnetic radiation is emitted or absorbed. An obvious example
is the warming of the Earth by the Sun. A less obvious example is thermal radiation
from the human body.
Introduction to Thermodynamics
Thermodynamics, science of the relationship between heat, work, temperature, and energy.
In broad terms, thermodynamics deals with the transfer of energy from one place to another and
from one form to another. The key concept is that heat is a form of energy corresponding to a
definite amount of mechanical work.
Heat was not formally recognized as a form of energy until about 1798, when Count
Rumford (Sir Benjamin Thompson), a British military engineer, noticed that limitless amounts of
heat could be generated in the boring of cannon barrels and that the amount of heat generated is
proportional to the work done in turning a blunt boring tool. Rumford’s observation of the
proportionality between heat generated and work done lies at the foundation of thermodynamics.
Another pioneer was the French military engineer Sadi Carnot, who introduced the concept of the
heat-engine cycle and the principle of reversibility in 1824. Carnot’s work concerned the
limitations on the maximum amount of work that can be obtained from a steam engine operating
with a high-temperature heat transfer as its driving force. Later that century, these ideas were
developed by Rudolf Clausius, a German mathematician and physicist, into the first and second
laws of thermodynamics, respectively.
I.
The zeroth law of thermodynamics.
When two systems are each in thermal equilibrium with a third system, the first two
systems are in thermal equilibrium with each other. This property makes it meaningful to use
thermometers as the “third system” and to define a temperature scale.
II.
The first law of thermodynamics, or the law of conservation of energy.
The change in a system’s internal energy is equal to the difference between heat added to
the system from its surroundings and work done by the system on its surroundings.
III.
The second law of thermodynamics.
Heat does not flow spontaneously from a colder region to a hotter region, or, equivalently,
heat at a given temperature cannot be converted entirely into work. Consequently, the entropy of
a closed system, or heat energy per unit temperature, increases over time toward some maximum
value. Thus, all closed systems tend toward an equilibrium state in which entropy is at a maximum
and no energy is available to do useful work. This asymmetry between forward and backward
processes gives rise to what is known as the “arrow of time.”
The second law of thermodynamics can be precisely stated in the following two forms, as
originally formulated in the 19th century by the Scottish physicist William Thomson (Lord Kelvin)
and the German physicist Rudolf Clausius, respectively:
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A cyclic transformation whose only final result is to transform heat extracted from
a source which is at the same temperature throughout into work is impossible.
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IV.
A cyclic transformation whose only final result is to transfer heat from a body at a
given temperature to a body at a higher temperature is impossible.
The third law of thermodynamics.
The entropy of a perfect crystal of an element in its most stable form tends to zero as the
temperature approaches absolute zero. This allows an absolute scale for entropy to be established
that, from a statistical point of view, determines the degree of randomness or disorder in a system.