From Thermodynamics to Economics and Ecology - Thanatia: The Destiny of the Earth's Mineral Resources - page 64

In order to break a system's equilibrium work needs to be done. In Physics, Work
is defined as force multiplied by displacement, which is expressed as W = Fdx.
Force and the displacement caused can be generalised. In this way, a changing
volume is associated to mechanical pressure, charge with electric force and mag-
netisation to the magnetic field. Pressure, electric force and the magnetic field are
all intensive properties. Their conjugated variables volume, charge and magnetisa-
tion are extensive properties, i.e. they are dependent on the size of the system.
When systems are surrounded by a thermal insulating (adiabatic) cover, the
quantity of adiabatic work required to move them from one equilibrium state to
another is independent of how and which type of work has been (or is being) done.
This is an experimental fact. The initial and final states are what really matter.
This is a transcendental fact in Physics and constitutes the First Law of Thermody-
namics, since if the work done is path independent, there is a property that depends
only on the state of the system. This property is what is known as energy, denoted
U and shares the same units as work. So what happens if the work done to a system
is not adiabatical? Then the work is no longer equal to the energy variation and
Eq. (3.1) applies. Let's call Q the difference:
Q = U i U f W (3.1)
Q is a type of energy, just like work. It only appears as long as the system is
evolving, since at the initial and final states, only U i and U f exist. Moreover it is
only present when the adiabatic cover is removed. The intensive property associated
with the adiabatic wall is temperature. Therefore, Q is closely related to a change
in temperature and coincides with the intuitive concept of heat.
Given that work and heat are both manifestations of energy, arguably there is
no need to assign different units for heat, say calories, and for work, say Joules.
But equivalence in units does not mean equivalence in size. In fact, assimilating
that work and heat are manifestations of energy in movement may be easy to
accept but di cult to understand. Whilst a calorie is a really small amount of heat,
namely that needed to warm a gram of water by one degree centigrade, it is also the
kinetic energy of that same mass at 327 km/h! Hence, small fluctuations in system
temperatures result in considerable energy expenditures. A Joule is often perceived
to be something larger: a kilogram mass at an acceleration of one metre per squared
second, performed along one metre. And yet this energy is equivalent to only
0.24 calories, which is to many almost imperceptible. This issue of perception has
important consequences when one tries to understand the true nature of energy: this
is because the senses, which are relatively sensitive to minor variations in mechanical
energy, through touching or hearing (sound) for example, serve to mislead.
In a closed system, heat's most impressive aspect is that it is the only mani-
festation of energy in movement apart from work. As there are no other types of
energy that share such characteristics, should the walls of a system breakdown, the
system will either do work or it will absorb or release heat or both simultaneously,
until it reinstates itself at equilibrium with a subsequently new set of surrounding