Chemistry - Is every chemical reaction in equilibrium?

Solution 1:

  1. Yes, every chemical reaction can theoretically be in equilibrium. Every reaction is reversible. See my answer to chem.SE question 43258 for more details.

  2. This includes even precipitation reactions and reactions that release gases. Equilibrium isn't just for liquids! Multiphase equilibria exist.

  3. The only thing that stops chemical reactions from being "in equilibrium" is the lack of the proper number of molecules. For a reaction to be in equilibrium, the concentrations of reactants and products must be related by the equilibrium constant.
    $$ \ce{ A <=> B} $$ $$ K = \frac{[B]}{[A]} $$

    When equilibrium constants are extremely large or small, then extremely large numbers of molecules are required to satisfy this equation. If $K = 10^{30}$, then at equilibrium there will be $10^{30}$ molecules of B for every molecule of A. Another way to look at this is that for equilibrium to happen, there need to be at least $10^{30}$ molecules of B, i.e. more than one million moles of B, in the system for there to be "enough" B to guarantee an equilibrium, i.e. to guarantee that there will be a well-defined "equilibrium" concentration of A.

    When this many molecules are not present, then there is no meaningful equilibrium. For very large (or very small) equilibrium constants, it will be very difficult to obtain an equilibrium. In addition to needing a megamole-sized system (or bigger), the system will have to be well-mixed, isothermal, and isobaric. That's not easy to achieve on such large scales!

Update Commenters suggest that "irreversible" reactions do not have an equilibrium. This is true, but tautological. In the real world, all reactions are reversible, at least to a (perhaps vanishingly small) degree. To say otherwise would violate microscopic reversibility. A reaction that was 100% irreverible would have an equilibrium constant of infinity. But if $K= \infty$, then $\Delta G^{\circ} = -RT \ln{K}$ would turn into $\Delta G^{\circ} = -\infty$. So to get infinite energy we would just have to use 100% irreversible reactions! Hopefully the problems with the idea of "irreversible" reactions are becoming apparent.

Solution 2:

Equilibrium can only apply to a closed system.

Reactions which form insoluble precipitates or gases which escape do not exhibit the behavior of a closed system. Therefore, these reactions may not be in equilibrium. However, these claims are pragmatic rather than real.

As it turns out, in the above answer, barium sulphate has a $\ce{K_{sp}}$ of $\ce{1.1 x 10^{-10}}$, so formally there is some small equilibrium related to the amount of barium sulphate in solution $\ce{1.05 x 10^{-5}}$

As gasses are escaping in solution, they may be readsorbed, an thus there would be some small equilibrium for processes like that.

But pragmatically, these reactions are not at equilibrium.

Solution 3:

Yes every reaction is an equilibrium. A complete reaction is a equilibrium with high equlibrium constant. If you write the expression for equilibrium constant, you will find that high equilibrium constant implies that the conc. of the products is very high, i.e. the reaction has reached completion.