Thermodynamics-notes

Thermodynamics-notes - © M S Shell 2010 1/20 last modified...

Info iconThis preview shows pages 1–3. Sign up to view the full content.

View Full Document Right Arrow Icon

Info iconThis preview has intentionally blurred sections. Sign up to view the full version.

View Full DocumentRight Arrow Icon
This is the end of the preview. Sign up to access the rest of the document.

Unformatted text preview: © M. S. Shell 2010 1/20 last modified 10/4/2010 Energy, thermodynamics, and molecular processes Overview Leading question: do cells use energy? Cells consume negative entropy, i.e., generate positive entropy Life = negative entropy generation Cells are bound by the laws of thermodynamics? Yes The central question is: how can complex, seemingly highly ordered processes and biological structures emerge “spontaneously” in living things Review of thermodynamics What do we need to do here? Cells are very small, so we need to understand thermodynamics at a molecular level. This is somewhat different than thermodynamics for macroscopic systems. In macroscopic, bulk systems, everything appears to be constant – the density, temperature, pressure, heat capacity, etc. At the molecular level, however, there is a constant molecular dance and things fluctuate . We don’t notice these fluctuations at the macroscopic level because they are so small. Example: the density in a fluid. The first law Thermodynamics defines a system and an environment, separated by a system boundary The first law is a statement of the conservation of energy for a process that involves a change to the system Here we consider closed systems for simplicity, although living systems are certainly open Closed systems = those with no mass exchange with the environment, but can exchange energy Differential form (small changes) for closed systems: gG ¡ ¢£ ¤ ¢¥ Integrated form (state changes) for closed systems: © M. S. Shell 2010 2/20 last modified 10/4/2010 Δg G ¡ ¢ £ Here, • g is the total internal energy in the system. In other words, it is the total potential energy due to molecular interactions as well as kinetic energy due to the velocities of the molecules • ¡ is the heat exchanged with the environment; it comes from energies stored in random molecular motions b positive for heat added to system • £ is the work done on the system; it comes from energies due to concerted molecular motions b positive for work done on system We can define pressure-volume work as £ G ¤¥¦§¨ Notice that the internal energy is a property of the system, whereas ¡ and £ describe flows between the system and environment g is a state function. Thermodynamics defines s tate functions as quantities that depend only on the current state of a system (e.g., T and P), and not the path by which they got there. Therefore, Δg does not depend on the path the process takes (e.g., the rate at which it happens), but only on the states at the beginning and end of the process On the other hand, Q and W are not state functions and do depend on the path Interestingly, the sum of Q and W is a state function. That means that any path-dependence of these quantities exactly cancels out. There can be many processes that take a system between two states 1 and 2 with very different Q and W, but their sum must be the same....
View Full Document

This note was uploaded on 12/29/2011 for the course CHE 170 taught by Professor Ceweb during the Fall '10 term at UCSB.

Page1 / 20

Thermodynamics-notes - © M S Shell 2010 1/20 last modified...

This preview shows document pages 1 - 3. Sign up to view the full document.

View Full Document Right Arrow Icon
Ask a homework question - tutors are online