[ikjadoon]

Hi. I'm in General Chemistry as a freshman in college. I'm having some trouble transitioning between kinetics and equilibria.

So, when we were doing rate laws:

rate = k[A]^x

I was told about a thousand times that when you have a chemical reaction, you DO NOT KNOW the rate-order just by looking at the reaction. You have to deduce a reaction mechanism and use the rate-limiting step's rate law for the overall rate-law. Or you can run experiments (change concentration, measure initial rate, find ratio, etc.) to get the raw law. Got that part.

Now we're doing equilibrium problems. Now the law of mass action (products/reactants and each concentration IS RAISED to it's stoichiometric coefficient) seems to throw all that out the window. You can look at a chemical reaction and get this equilibrium equation, which uses the the coefficients as the powers. Forget about rate-limiting steps, reaction mechanisms, whatever. 

I understand how K_c is a ratio between rates (forward-rate/reverse-rate) and that an equilibrium reaction just tells me which is favored (product or reactant), but it DOES NOT tell me anything about how fast the reaction will actually go right?

rate=k[A]^x

In the equation above, k is DIRECTLY proportional to rate. Higher k, higher rate. But now a ratio between rates isn't a rate?

I'm pretty confused. For some reason, my head can't wrap around this new "law of mass action".  My Chemistry book (Chang, 10th ed.) actually has a whole section which explains it. It boils down to one sentence:

"Regardless of whether a reaction occurs via single-step or a multistep mechanism, we can write the equilibrium constant expression according to the law of mass action..."

Why are we now allowed to raise concentrations to the coefficients? Overall: what does this mean in a equilibrium constant expression? What does raising the powers actually do? In rate-laws, the powers told how many of each molecule/atom/whatever was needed in the rate-limiting step. What does it mean here?

Thanks...

~Ibrahim~

[Yggdrasil]

The main difference here is that equilibrium constants are independent of the reaction mechanism whereas kinetics are dependent on the reaction mechanism.

The equilibrium constant for any given reaction is dependent only on the free energies of the reactants and products.  Because free energy is a thermodynamic potential it is a state function and the value of ΔG of a reaction is independent of the path taken from reactant to product.  The kinetics of a reaction, however, is dependent on the path taken from reactant to product as different paths will involve different reaction mechanisms and different activation energies. 

A simple chemical reaction written as A + B --> C + D tells you only the reactants and products of the reaction, and does not give you any information about the path taken between reactant and product.  Because you know the start and end points, you have all the information you need to determine the equilibrium constant.  However, because the reaction does not tell you the path, you cannot know anything about the kinetics without more information.

As an analogy, think back to working with ideal gasses.  If you know the starting pressure, temperature and volume as well as the ending P, T and V, you can calculate the ΔE, ΔH, etc. for the transformation.  However, this information is not enough to tell you about the heat transfer and work involved in the transformation.

[ikjadoon]

Ohhh, like path functions versus state functions. That makes sense!

Thanks a jillion; I have a test Wednesday and this has cleared up so much for me!! Thanks,

~Ibrahim~