For example, consider the single-displacement reaction $$\ce{AB + C <=> BC + A}$$ My question is then if writing this reaction as the system \begin{gather} \ce{AB <=> A + B} \\ \ce{B + C <=> CB} \end{gather} is equivalent (surely, adding the reactions up gives the correct stoichiometric coefficients).

The thought behind this question is confusion regarding the notion of a "displacement" reaction. I am confused why in textbooks the displacement reaction is presented as if C is reacting with AB directly; wouldn't C instead be reacting with the free B present in the solution (i.e., the reaction is written more fundamentally as the two reaction system)?

The only time I can think of when this might not be true is if B being bound to A induces some electronic structure change localized on B that makes C able to react with B (and free B is unable to react with C).

  • 2
    $\begingroup$ Either is possible, but if AB is a normally stable molecule (in absence of C) then your second scheme seems unlikely, but only experimental data can sort this out. $\endgroup$
    – porphyrin
    Sep 12, 2022 at 7:40
  • 1
    $\begingroup$ Just read about SN1 and SN2. $\endgroup$
    – Mithoron
    Sep 12, 2022 at 18:51
  • $\begingroup$ As @porphyrin suggested, you can not make a general guess in absence of experimental data or chemical knowledge for a concrete reaction. Just a simple example of the first case, an elementary isotope exchange gas phase reaction $$\ce{H_2 + D <=> HD + H}$$ $\endgroup$
    – PAEP
    Oct 2, 2022 at 21:17

1 Answer 1


Upon thinking about this most displacement reactions follow your scenario.

If the two reactions are independent then A is not necessary to the production of BC and is not in the equation for its production. This is a common fallacy in describing consecutive or concurrent reactions especially redox reactions that are often multistep and only accidently stoichiometric because the kinetics are too fast to see what happens. This is what is behind the idea of the net ionic equation for reactions. In solution NaCl +AgNO3 = AgCl + NaNO3 This happens regardless of the Na+ ion [but possibly not regardless of the nitrate because there are few soluble silver salts]. It is correct to write Ag+ + Cl- = AgCl not NaCl + Ag+ = AgCl + Na+ unless it were done in a solid-state reaction. This example meets your requirements because it is two separate independent reactions: the dissolution of the two salts and the forming of the precipitate. Let's dissolve some HCl gas in water HCl + H2O = Cl- + H3O+ Now the H and the Cl cannot be separated in a twostep reaction. Hopefully you woke some of us up a bit. Equations are a synopsis of a process. They might relate somewhat to the actual mechanism but mechanisms are usually more complicated than the equation. Keep that in mind and look behind the scene.


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