- Physical Chemistry I
Spring 2012, Unique 52135
Lecture Summary, 20 January 2012
to Thermodynamics: We have spent most of our first
two days defining the subject of thermodynamics, and discussion the
most important principles of the field. Thermodynamics is the
science that studies the effect on matter from the transfer of
energy through heat or work. Thermodynamics does not require
any assumptions or understanding of the physical nature of the
underlying system. Because of that, this science can be used
to understand the properties of systems in which the physical
composition is overwhelmingly complicated or completely
state: a system containing a set of defined properties that do not depend on history.
property: a variable that can be measured, implied, or manipulated.
path: the process by which a state moves from an initial to a final set of properties. Paths do depend on history and circumstances. Much of our study of thermodynamics will be spent trying to define paths in an intelligent way.
equilibrium: the point at which all forces are equal; the state at which the system is at a minimum potential energy.
reversibility: a path that is in instantaneous equilibrium at every point. There are very few macroscopic examples of reversible paths (frictionless pendulum). However, this is an incredibly important concept of microscopic paths.
Ideal Gases: The first system that we explore in detail is the ideal gas. An ideal gas obeys the following requirements:
1) The particles are non-interacting point spheres.
2) The size of the particles are much smaller than the distance between them.
3) Collisions between the particles or with the walls of the vessel are rare, and when they do occur, they are perfectly elastic.
We performed a few thought experiments to arrive at the following empirical (i.e. derived through experience) conclusions:
Boyle's Law: PV = constant
Charles' Law: P = constant x T
V = constant x T
When combined, these observations lead to the ideal gas law:
PV = nRT
Our first state function.
A corollary: if we don't care about the molecular nature of the underlying system, then it doesn't matter if our ideal gas is composed of one thing or many. This leads to Dalton's Law, which says that there is no distinction between homogeneous and heterogeneous mixtures of ideal gases:
P(total) = sum(Pi)
where Pi is the partial pressure of each individual component of the ideal gas.