Tuesday, March 27, 2007
Tues-Day 1
AP: Thus far, I am very impressed with the effort and quality on your Solubility Eq. exam, the hardest exam of the year. I know objectively that this was the hardest test, by far, this year yet you brought your A game and are doing very well.
Today, we talked more about entropy and we derived the second law of thermodynamics, one of the most important equations in life: dS (univ) = ds (sys) + dS (surr) > 0 for any spontaneous process. "You can't break even" i.e. any exothermic reaction that "happens" (spontaneously) will always lose/release some energy as dissipated heat that cannot be used for work. This "waste"-heat will cause a net increase in the entropy of the universe.
We then rearranged this equation into the Gibbs Free Energy equation. Noting that dS and dH are remarkably constant over large temperature ranges, dG can be determined over a range of temperatures by using dG = dH - TdS. This way, we can use the relatively constant values of dH and dS (so we can just use the standard temperature values of dH and dS), we can vary T and see at what temperature a reaction becomes SPONTANEOUS (i.e. no continuously applied BATTERY/energy source needed for the reaction to continue), which is when dG just becomes negative.
Honors: discussed potential energy/enthalpy diagrams and related them quantitatively to formation reactions and other chemical reactions or physical processes. We related the reactions in Tables G and I to the energy diagrams; we also discussed the meaning of "mole of reaction" and how to treat the energy term in an equation proportionally to the stoichiometry of the reaction.
Regents: we reviewed the factors that affect rates of chemical reactions; then we began to discuss exothermic and endothermic reactions as related to enthalpy diagrams.
Today, we talked more about entropy and we derived the second law of thermodynamics, one of the most important equations in life: dS (univ) = ds (sys) + dS (surr) > 0 for any spontaneous process. "You can't break even" i.e. any exothermic reaction that "happens" (spontaneously) will always lose/release some energy as dissipated heat that cannot be used for work. This "waste"-heat will cause a net increase in the entropy of the universe.
We then rearranged this equation into the Gibbs Free Energy equation. Noting that dS and dH are remarkably constant over large temperature ranges, dG can be determined over a range of temperatures by using dG = dH - TdS. This way, we can use the relatively constant values of dH and dS (so we can just use the standard temperature values of dH and dS), we can vary T and see at what temperature a reaction becomes SPONTANEOUS (i.e. no continuously applied BATTERY/energy source needed for the reaction to continue), which is when dG just becomes negative.
Honors: discussed potential energy/enthalpy diagrams and related them quantitatively to formation reactions and other chemical reactions or physical processes. We related the reactions in Tables G and I to the energy diagrams; we also discussed the meaning of "mole of reaction" and how to treat the energy term in an equation proportionally to the stoichiometry of the reaction.
Regents: we reviewed the factors that affect rates of chemical reactions; then we began to discuss exothermic and endothermic reactions as related to enthalpy diagrams.
# posted by C @ 2:14 PM 
