- #1
Hunt_
- 26
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One thing that always bothered me in Chemistry is how the equilibrium constant is written. It never made sense to me.
If I take a simple bimolecular reaction approaching equilibrium :
[tex] aA + bB \mathop{\rightleftharpoons}^{k_1}_{k_{-1}} cC + dD [/tex]
From ART ,
[tex] r_1 = k_1 [A]^ \alpha ^ \beta [/tex]
[tex] r_{-1} = k_{-1} [C]^ \gamma [D]^ \delta [/tex]
Then if we consider the rate to be equal at equilibrium and the expression of the equilibrium constant ,
[tex] K = \frac{k_1}{k_{-1}} = \frac{[C]^ \gamma [D]^ \delta }{[A]^ \alpha ^ \beta } \neq \frac{[C]^c [D]^d }{[A]^a ^b} [/tex]
The only way for both expressions to be equal is that the reaction is elementary , which doesn't hold for most chemical reactions.
So what does that mean ? There can be approximations here , molecularity is not equal to stoichiometry. Am I missing something here or are all the claculations I made in equilibrium chemistry just wrong ?
If I take a simple bimolecular reaction approaching equilibrium :
[tex] aA + bB \mathop{\rightleftharpoons}^{k_1}_{k_{-1}} cC + dD [/tex]
From ART ,
[tex] r_1 = k_1 [A]^ \alpha ^ \beta [/tex]
[tex] r_{-1} = k_{-1} [C]^ \gamma [D]^ \delta [/tex]
Then if we consider the rate to be equal at equilibrium and the expression of the equilibrium constant ,
[tex] K = \frac{k_1}{k_{-1}} = \frac{[C]^ \gamma [D]^ \delta }{[A]^ \alpha ^ \beta } \neq \frac{[C]^c [D]^d }{[A]^a ^b} [/tex]
The only way for both expressions to be equal is that the reaction is elementary , which doesn't hold for most chemical reactions.
So what does that mean ? There can be approximations here , molecularity is not equal to stoichiometry. Am I missing something here or are all the claculations I made in equilibrium chemistry just wrong ?