ln5s08 - Lecture 5: Introduction to Chemical Equilibrium...

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

View Full Document Right Arrow Icon
Lecture 5: Introduction to Chemical Equilibrium This lecture is outside the confines of the textbook development of chemical equilibrium and is intended as a sweeping introduction to chemical equilibrium (a topic that will engage us for much the rest of the semester.) I hope you will see just how easily equilibrium concepts can be applied to other areas of endeavor in and outside of the sciences. In class today, for example, we will have a house party and see that the same stresses that drive chemical reactions one direction or another, quite logically drive the actions of young adults engaged in an activity I have heard young people cal “partying.” Having never “partied” myself, I take it at your word this is an enjoyable exercise. Some of the big picture concepts that you should appreciate by the end of the lecture: Just like thermodynamics, equilibrium theory can tell you about whether a reaction occurs, not by evaluating the flow of energy but rather by following the flow of chemical compounds The equilibrium constant, K, is a very simple ratio of stuff on one side of an equation to stuff on the other—it is the mathematical structure used in equilibrium theory to tell you at a glance whether a reaction happened or not. Stress is a concept that both chemical reactions and students hate, and both chemical reactions and you respond to it in exactly the same way—you run from it. Following chemical concentrations to determine it a reaction happens. Most of you have a prejudiced and incorrect view that a chemical reaction starts with a bunch of reactants you pull off of a shelf that are allowed to react to completion—at completion only product remains. For example, consider this reaction below for the combustion of a hydrogen balloon: For example: 2 H 2 + O 2 Q 2 H 2 O Initial amounts 100 50 0 Final amounts 0 0 100 This little bit of fiction even uses the exact ratio of reactants so that when completed, every last bit of reactant is used up (there is no “limiting reagent”. But in fact, this fiction never happens. Not only do we always have stuff “left over” but the reaction never goes to completion . Instead it goes to an equilibrium concentration in which that amount of H 2 , O 2 and H 2 O are together
Background image of page 1

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

View Full DocumentRight Arrow Icon
in amounts such that the rate of forward and reverse reactions is equal so that there is no overall change in the system concentrations. So here is the truth of a reaction:
Background image of page 2
Image of page 3
This is the end of the preview. Sign up to access the rest of the document.

Page1 / 10

ln5s08 - Lecture 5: Introduction to Chemical Equilibrium...

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