Professor Dave here, let's discuss the
first law of thermodynamics.
The first law of thermodynamics, popularly
known as the law of energy conservation,
when examined more rigorously actually
outlines the relationship between
internal energy, work, and heat, which from
now on will be represented by the letter Q.
The law can be stated as an equation
where Delta U equals Q minus W.
This means that the change in the internal
energy of a system will be equal to the
energy transferred to or from the system
as heat minus the energy transferred to
or from the system as work, and all of
these quantities will be measured in
joules. Because of this law, we can
outline a few different types of
processes that can occur. If a process
occurs where there is no change in
volume for the system, that means that no
pressure-volume work can be done on or
by the system, so work is zero in such a
case. Delta U equals Q and any change in
internal energy must be the result of
heat transfer in or out. This will be
called an isovolumetric process, meaning
no change in volume. An example would be
a bomb calorimeter where a combustion
reaction produces a change in
temperature but the rigid walls results
in no change in volume. If there is no
change in the temperature of the system
there cannot have been any change in the
internal energy of the system, since
these two values are proportional. Delta
U will be zero which makes Q equal to W.
This means that any heat transferred
into the system is used by the system to
do work rather than increasing the
internal energy of a system. This is
called an isothermal process, meaning no
change in temperature. An ideal version
of a car engine would be an example of
this as the pistons ought to convert all
of the heat energy from the combustion
reaction directly into expansion
work that moves the car. If there is no
heat transferred, Q will be 0 and Delta U
will equal negative W. This means that
the internal energy of a system changes
as a result of doing work on its
surroundings or the surroundings doing
work upon the system. Such a process will
be called an adiabatic process, meaning
no heat transfer. We can see this in
certain processes in Earth's atmosphere
as masses of air change position due to
pressure differences. And if Q and W are
both zero, meaning there is no heat
transfer and no work done, there can be
no change in internal energy and this
must be an isolated system. Hopefully
these scenarios make some intuitive
sense because we will frequently use
this equation to do calculations. We will
also need to define the signs of these
quantities in order to use this equation
properly so let's note that when heat is
absorbed by the system, Q will be
positive. If heat is lost by the system, Q
will be negative. If work is done by the
system, like an expanding gas, W will be
positive. If work is done on the system
by the surroundings, like gas compression,
W will be negative. If there is no
transfer of heat or no work done these
values can also equal 0 as we have
previously discussed. When doing
calculations make sure that you use the
correct signs for these values or the
math will be incorrect. For example, if
100 joules of compression work is done
on a system and as a result the internal
energy of the system increases by 74
joules, how much of the energy is
transferred as heat and in which
direction? Let's take our equation and
rearrange to solve for Q, which will be
Delta U plus W, then we can plug in
positive 74 joules for Delta U, since
internal energy increases, and negative
100 joules for work since work is being
done on the system, and we should get
negative 26 joules for heat. This means
that as 100 joules of work is applied to
the system,
only 74 go towards increasing the
internal energy of the system while 26
joules are lost as heat dissipates out
of the system. These kinds of
calculations will happen a lot in
thermodynamics so let's check comprehension.
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