Today we went through a few more equilibrium calculations and looked at a couple ways equilibria can be simplified. First, the "concentrations" of pure solids and liquids do not appear in equilibrium constants because their concentrations are constant (essentially) throughout the reaction.
Next we looked at assumptions that can simplify the math involved in equilibrium calculations. In many problems, the numerical value of "x" {the change in concentration when a reaction goes to equilibrium} is so small that it can be ignored in places where it's added to or subtracted from another term. For example, if the term "1.328 - 3x" appears in your equilibrium calculation, you should examine just how big you expect "x" to be. If the value of "x" turns out to be 0.00000000168, then it really doesn't affect the term "1.328 - 3x", and that term can be simplified to "1.328". Knowing when to make these approximations takes a little practice, so always check your assumptions after you finish working through a problem.
FInally, we looked at the difference between using a "full molecular" chemical equation and a "net ionic" chemical equation to determine the equilibrium constant. In many ways, the net ionic equation will always give you a more correct picture of the actual chemistry that takes place, but as long as the equilibrium constant you calculate is consistent with the chemical equation you use, I don't mind which you choose. But I strongly encourage you to think about reactions as net ionic reactions whenever possible...
There's a new Mastering Chemistry assignment, due Friday.
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