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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

[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 doesnt 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 ?The only way for both expressions to be equal is that the reaction is elementary , which doesnt 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 ?