1st post = 10 points
2nd post (reply to someone else's) = 5 points
total = 15 points
When we explored stoichiometry earlier in the course, we worked with reactions in which the limiting reagent was entirely consumed. This is characteristic of reactions that go to completion. The products are formed until one of the reactants has been completely consumed.
As we've just seen, many reactions don't go to completion. They reach a state in which there is a mixture of reactants and products. We say that the reaction is reversible because when equilibrium is reached, the forward and reverse directions of the reaction are occurring at the same rate.
In theory, all reactions are reversible, as long as the reactants and products remain in contact with one another. But what does this mean in practice? Did we make some incorrect assumptions about the reactions we studied earlier in the course?
Consider the combustion of methane:
CH4 (g) + 2 O2 (g) 
CO2 + 2 H2O (g)
Notice that we've shown this as a reversible reaction. However, when we studied the ideal gas law, we used the reaction stoichiometry to calculate the volume of carbon dioxide that would be produced from a certain mass of methane. In this case, we assumed that the reaction went to completion.
Now that you know about dynamic equilibrium and the equilibrium constant, consider the following question: When is it safe to assume that a reaction goes to completion, and when must you treat it as a reversible reaction? What information would you need in order to make a decision, and how would you use that information to decide whether or not a reaction is reversible?
