The first two dot points, regarding enthalpy diagrams and collision theory, are explained in my post on Reaction Rates: http://year11misadventures.blogspot.com.au/2012/10/reaction-rates.html.
Now for the new stuff! Half this stuff probably needs diagrams but I can't be bothered doing them just yet, so I might get back to this post later (key word here is "might." Great blog writer I am).
Describe and explain the characteristics of a system in dynamic chemical and physical equilibrium
A lot of the reactions that we have dealt with thus far go to completion: that is, all of the reactants are fully consumed to form products (unless there is a limiting reagent, in which case only one of the reactants is fully consumed). However, there are a lot of reactions that do not go to completion. One such reaction which you would have probably encountered is the ionisation of acetic acid. Acetic acid only partially ionises into acetate/ethanoate ions and hydrogen ions: the rest remains as acetic acid molecules. Well, actually, that isn't totally true, because at any one time some acetic acid is dissociating into ions and some acetate and hydrogen ions are combining to form acetic acid molecules. The catch is, both of these processes are happening at the same rate, so the system is therefore at equilibrium.
Since the two reactions are proceeding at the same rate, the decrease in concentration of any one particular substance is counterbalanced by an equal increase due to the opposite reaction, and vice versa. Therefore, when a system is at equilibrium, the concentrations of the substances in the system are constant. This also means that other macroscopic properties ("macroscopic" being a word I only learned just then by glancing in my textbook), such as colour, also remain constant.
Write equilibrium law expressions for homogeneous and heterogeneous systems
Writing equilibrium law expressions is the easiest part of learning about equilibrium. Basically, on the top line, you have the concentrations of each of the products raised to the power of their coefficients, and on the bottom line you do the same thing for the reactants. For example, in the Haber process of producing ammonia, N2 + 3H2 ßà 2NH3, the equilibrium law expression is:
K = ([N2] + [H2]^3)/ ([NH3]^2)
(K is the symbol used for equilibrium constants. Don't ask.)
Remember that solids and liquids do not get included in this expression, as they don't have concentrations. Only aqueous solutions and gases get included.
If your products are all solids and/or liquids, put 1 on the top line. One way of thinking about that is that the concentration of a solid or a liquid is equal to 1 (though obviously not strictly true as, as I just said, solids and liquids don't have concentrations).
Use K and equilibrium law expressions to explain the relative proportions of products and reactants in a system of dynamic chemical equilibrium
Since the products are the numerators and the reactants are the denominators, when there's a relatively large quantity of products formed, the equilibrium constant is relatively big. Similarly, a small equilibrium constant is indicative of the formation of a relatively small amount of products.
Explain, using the collision theory, the effect on the position of equilibrium when the following changes are made to a system initially at chemical equilibrium.
Changes in solution concentration: When the concentration of one of the aqueous solutions in the system gets changed, the system is no longer at equilibrium because the equilibrium law expression no longer equals whatever equilibrium constant it was equal to before (and the only time that an equilibrium constant changes for a solution is when the temperature is changed). Therefore, equilibrium shifts in favour of the reaction that will help to bring the concentration of said solution back to what it was before: if the concentration was increased, then the reaction that consumes that solution will be favoured, and if the concentration was decreased, then the reaction that forms that solution will be favoured.
Oh, wait, I was meant to use the collision theory. Oops. Well, increasing the concentration of an aqueous solution means that there are more particles of that solution, so that the chances of them colliding with other reactant particles are increased and, therefore, that reaction proceeds faster. If the concentration is decreased, then they collide less often and, therefore, that reaction proceeds slower.
Changes in partial pressures of gases: Partial pressures are kind of the same as solution concentrations (as far as I know, anyway). So just see what I've written for solution concentrations above. However, if the overall system pressure is changed, other stuff happens too, which I'll talk about later.
Addition of a catalyst: Catalysts do not affect equilibrium- they only affect the rate at which a system reaches equilibrium.
Predict, using Le Châtelier's principle, the impact of certain changes to a system initially at chemical equilibrium
Before I dive into all the prediction stuff, I first need to explain what Le Châtelier's principle is. Basically, Le Châtelier's principle states that if a change is made to a system at equilibrium, for example the concentration of one of the substances has been altered, or the temperature has been changed, then the system will do something to at least partially counteract this change. For example, if the pressure is increased, the rate of the reaction that produces fewer particles and thus partially counteracts the change in pressure will be favoured. And so on, and so forth. Let's take a closer look:
Addition of a catalyst- I'm starting with this one, because it's easy. As I said before, catalysts do not affect equilibrium- only the rate at which a system reaches equilibrium.
Changes in solution concentration/ partial pressure of gases: I accidentally touched on this in the previous section (the section where I was meant to talk about collision theory).
Change in pressure: When pressure is raised, all gases have a higher concentration, so both forward and reverse reactions proceed faster. However, one reaction will initially proceed faster than the other. The reaction favoured in this case is the one that produces fewer particles. For example, in the Haber process, N2 + 3H2 ßà 2NH3, the production of ammonia would be increased if pressure was to increase as the forward reaction produces fewer particles than the reverse reaction (2 particles of ammonia for the forward reaction and 1 particle of nitrogen gas and 3 of hydrogen gas for the reverse reaction). The inverse is also true- if pressure is decreased, the reaction producing more particles is favoured.
Change in temperature: When temperature is raised, both reactions proceed more quickly (see my previous post on Reaction Rates). However, one will proceed more quickly than the other. To cool down the system, the endothermic reaction is favoured. The inverse is also true- if temperature is decreased, the exothermic reaction is favoured in order to heat up the system.
By the way, changing the temperature is the only way to change the equilibrium constant (K).
Interpret changes, such as colour changes, of physical and chemical systems at equilibrium
For changes like these, simply work out what colour (or whatever) each "side" of the reaction is meant to have- for example, the reactant ions might be yellow while the product ions are orange. As you do stuff to the system, work out whether the changes move more in favour of the product side or the reactant side, for example, is the solution becoming more yellow or more orange?
Conditions of Industrial Processes
Sometimes you get given some random industrial process and you have to predict the ideal conditions under which to perform these processes.
First of all, you should pretty much always suggest using a catalyst, as catalysts increase reaction rate without having any effect on the yield.
Secondly, use Le Châtelier's principle to work out whether an increase or decrease in temperature will favour the products. Then do the same for pressure. If an increase is required, that's great, since increasing temperature and pressure also increases reaction rate. If a decrease is required, then say that a "medium" temperature or pressure is required, since having a really low temperature or pressure will just make the reaction proceed at a snail's pace.
However, high temperatures and pressures cannot be indefinitely high, as there's economic and safety limits on stuff. Plus our technology is not infinitely good either.
Describe and explain the conjugate nature of buffer solutions
Buffer solutions? Never heard of these before... Oh, wait, my textbook has them in the next chapter. So maybe that's why I haven't heard of them before. Maybe this dot point got put in the wrong place or something. Or maybe we're meant to apply Le Châtelier's principle to buffer solutions once we learn about them. Stay tuned...
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