In teaching secondary science equilibrium, how can we actually show there is a backwards reaction occurring during equilibrium rather than no reaction. A lot of equilibrium experiments involve external changes such as temperature or change in pressure. This to me shows that a reaction can occur backwards, i.e reversible but it doesn't show that there is a forward and backwards reaction occurring at the same time at equilibrium. What's a simple science experiment or everyday life experience that equilibrium can explain where the idea that an reaction just stops taking place doesn't?

Every video just claims that there is a forward AND backwards reaction but doesn't really explain why and is also something I've taken for granted


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    $\begingroup$ The only way that I can think of to show that there is a dynamic equilibrium would be to use isotopes. Way beyond a secondary science experiment. // To show that reactions are dynamic one could use the iodine clock reaction, but that really doesn't show a dynamic reaction at equilibrium.// Lots of ways to show that a reaction can go backwards and forward, but how to show that happening dynamically is an interesting question. $\endgroup$ – MaxW Apr 16 '20 at 7:30
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    $\begingroup$ @MaxW either radioisotopes, either using enrichment by a minor stable isotope. Then the isotope redistribution can be traced by Mass Spectroscopy. $\endgroup$ – Poutnik Apr 16 '20 at 7:35
  • $\begingroup$ @Poutnik - That was my first thought. However I doubt that any middle schools or high schools in the US has a mass spec. My other thought was Ostwald ripening. Show that a centrifuge can't clear the colloidal size particles, digest for an hour, then show that the centrifuge does clear the particles. $\endgroup$ – MaxW Apr 16 '20 at 7:59
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    $\begingroup$ A colour change might be a good way of testing this in the classroom, although not ideal. Perhaps phenolphthalein or methyl orange that have a different colour in acid /base. Adding acid /base should swap the colour back and forth about equilibrium. You would need to find a reaction with a pK v. close to the indicator colour change pH. $\endgroup$ – porphyrin Apr 16 '20 at 8:03
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    $\begingroup$ @MaxW Yes, I agree. Not sure about radioisotopes though, how much better would be available. I may have misread the question, considering just theoretical ways. But how about a precise gravimetry, if e,g. heavy water is inlvolved ? Like $\ce{HX + D2O <-> DX + HDO}$ Or using other deuterize compounds. $\endgroup$ – Poutnik Apr 16 '20 at 8:47

By definition, the macroscopic observables do not change once equilibrium is reached, and experiments that probe the microscopic domain are usually expensive. I use an analogy to show the microscopic changes:

The model

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Take two beakers A and B, and fill one or both with some water (and add some food color for better visibility), and use two straws F and R to continuously transfer from beaker A to beaker B (with straw F) and from beaker B to beaker A (with straw R). The transfer is by having the straw touch the bottom, and then covering the top with your index finger or thumb.

At that point I ask students what they expect will happen if I continue for the rest of the class period. At one point, the level of liquid does not change anymore even though I am still transferring liquid. This is the state of dynamic equilibrium.

To make it more interesting, take straws of unequal diameter (or two vs one straw, like in the looped video).

The analogy explained

The liquid level in the two beakers represent the concentration of reactant A and product B. The diameter of the straws represent the rate constant of the forward reaction (straw F) and the reverse reaction (straw R).

$$\ce{A <=>[k_f][k_r] B}$$


There is only one reactant and one product. The reaction will always be first order.

Avoiding misconceptions

If you take a single straw to transfer back and forth, students might think that equilibrium is like a pendulum. If you take straws of equal diameter, they might think equilibrium means that concentrations are equal (rather than correctly, rates are equal).

Every video just claims that there is a forward AND backwards reaction but doesn't really explain why and is also something I've taken for granted

Starting with just A, you learn that the A molecules have the capability of turning into B. Starting with just B, you learn that the B molecules have the capability of turning into A. If equilibrium were static, someone would "have to tell" the molecules to stop reacting. With the dynamic equilibrium model (which has support from expensive molecular-level experiments), molecules keep on doing what they always did - react every now and then. There is no requirement for some kind of switch in behavior, it just is an emergent property of having both forward and reverse reactions going on and influencing concentrations of species.


A reversible equilibrium can be shown in the high school level with cobalt chloride $\ce{CoCl_2·6H_2O}$, which can be written as $\ce{ [Co(H_2O)_6]Cl_2}$

If a couple of milligrams of usual cobalt chloride $\ce{CoCl_2·6H_2O}$ is dissolved in ethanol in a test tube, a blue solution is obtained. This is due to the reaction :$$\ce{2 [Co(H_2O)_6]Cl_2 -> [Co(H_2O)_6]^{2+} + CoCl_4^{2-}}$$ The two ions produced in this reaction are $\ce{[Co(H_2O)_6]^{2+}}$ which is pale pink, and $\ce{CoCl_4^{2-}}$ which is dark blue. As a result, the mixture looks blue.

Now if you add one or two drops water in this blue solution, the color changes to pale pink, due to the exothermic reaction :$$\ce{CoCl_4^{2-} + 6 H_2O -> [Co(H_2O)_6]^{2+} + 4 Cl^-}$$

Now this equation is an equilibrium that can be easily reversed, according to the equation : $$\ce{CoCl_4^{2-} + 6 H_2O <=> [Co(H_2O)_6]^{2+} + 4 Cl^-}$$

To show that the equation can be reversed, you heat the pink solution. At about $60°$ C or $70°$C, the solution turns blue again : the reaction is reversed, and goes to the endothermic side. By the way, it also favorises the side with the maximum entropy.

This reaction can be reversed by cooling it to room temperature, where it goes back to pink. The same mixture can be reheated and re-cooled, to show that is is reversible.

The same reaction can be obtained with copper chloride $\ce{CuCl_2·2H_2O}$, where the color goes from green to blue.

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    $\begingroup$ I think this does not asnwer the question, that is not about reversible equilibrium. It is about to show equilibrium is a dynamic, not static process, where there are ongoing the both opposite reactions, but with the zero net result. $\endgroup$ – Poutnik Apr 16 '20 at 10:48
  • $\begingroup$ But otherwise nice experiment ! $\endgroup$ – Poutnik Apr 16 '20 at 11:37
  • $\begingroup$ I thought also any simple phase change such as freezing/melting water or solubilization/crystallization of a salt might do. But this is much neater. $\endgroup$ – Buck Thorn Apr 16 '20 at 15:40

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