2
$\begingroup$

In chemistry textbooks there is generally written that ionic equilibrium can’t be achieved in the case of strong electrolytes (because their degree of dissociation is almost equal to 1) but it can be achieved in the case of weak electrolytes(because their degree of dissociation is much much less than 1) Now my question is that what does degree of dissociation has to do with whether the reverse reaction is possible or not. Suppose we have a strong electrolyte which almost dissociates completely now why don’t the reaction occurs in the backward direction(please don’t say that it is so because the dissociation is almost 100%. If so then what, it doesn’t mean that reaction can’t occur in the backward reaction there is possibility that the almost 100% dissociated ions may again recombine to make the reaction go backwards and reach a state where the rate of the forward equals the rate of backward). But this doesn’t happen(acc. To textbooks) why?

According to me(my understanding of Arrhenius’ theory of electrolytic dissociation) the electrolytes dissociate in water to furnish ions and these ions are continuously recombining to form the electrolyte molecules and thus are in equilibrium and therefore the strong electrolytes must also follow the same principle but we can say that in this case particularly the equilibrium is very highly established towards the product side(but it will be still said to be equilibrium however product sided it may be)

One more fact in favour of my argument i.e. the values of equilibrium constants are present on the web for strong acids so if the equilibrium never exists in the case of strong electrolytes then how come we are able to calculate these constants?

NOTE- I’m a high school student.

$\endgroup$
3
  • 4
    $\begingroup$ Well there can be an equilibrium. If you dump an excess of NaCl into water, then the NaCl in solution is in equilibrium with the solid NaCl that didn't dissolve. $$\ce{NaCl(s) <=>[excess NaCl] Na+(aq) + Cl-(aq)}$$If you don't add enough salt to make a saturated solution, then the system isn't at equilibrium since more NaCl could dissolve into the solution. $\endgroup$
    – MaxW
    Apr 5, 2020 at 23:12
  • $\begingroup$ Tangible comment $\endgroup$ Apr 6, 2020 at 4:33
  • $\begingroup$ @MaxW, why not make it an answer? $\endgroup$ Nov 19, 2023 at 18:18

3 Answers 3

1
$\begingroup$

You seem to be fretting over the definition of strong versus weak electrolytes, the notion of chemical equilibria, and what I'll call exactness.

Let's discuss exactness first. Mathematicians have calculated $\pi$ to trillions of decimal places. Chemistry doesn't work like that. Models in chemistry that are good to 3 or 4 significant figures are gospel. So if an electrolyte dissociates to 99.99%, a chemist would say that it is "completely" dissociated.

The whole science of chemistry is based on the fact that reactants combine to form products driving towards some equilibrium. So there is always some equilibrium however lopsided that might be.

Let's now consider hydrochloric acid. Concentrated hydrochloric acid is made by saturating water with hydrogen chloride gas. So there are two different equilibria. First an equilibrium of hydrogen chloride molecules between the gaseous phase and those which are dissolved in water, and secondly an equilibrium between hydronium ions and chloride ions with hydrogen chloride molecules in the liquid phases. Concentrated hydrochloric acid is 12.4 molar, meaning that one liter of solution reacted with $\ce{7.5 x 10^{24}}$ molecules of hydrogen chloride gas. But the concentration of hydrogen chloride molecules in the liquid phase is effectively zero.

If your pour some concentrated hydrochloric acid into a beaker you can smell the acid in the air. Hydronium and chloride ions don't jump out of solution into the atmosphere to form a hydrogen chloride molecule, so there must be some hydrogen chloride molecules in solution. So the solution is trying to establish an equilibrium between hydrogen chloride molecules in the gas phase and hydrogen chloride molecules in the liquid phase. There are essentially no hydrogen chloride molecules in the bulk atmosphere. So ignoring the water evaporation (or replenishing it daily), this means that eventually all the hydrogen chloride in solution would dissipate into the atmosphere. Eventually in this case being a very very very long time.

$\endgroup$
1
  • $\begingroup$ Water evaporation cannot be ignored; dilute strong acids will concentrate because the vapor pressure of water is higher, concentrated acids will evaporate acid until an azeotrope is usually formed. This is why acid spills should be cleaned up and neutralized to prevent hidden drops of conc. acid forming. $\endgroup$
    – jimchmst
    Nov 20, 2023 at 1:43
1
$\begingroup$

Certain compounds that are considered ionic as solids can form significant amounts of uncharged (but obviously polar) associated species in water solution, and we can measure equilibria associated with such associations.

One of the best documented examples is magnesium sulfate, $\ce{MgSO4}$. With ions having $+2$ and $-2$ charges, this can readily form uncharged pairs in water solution.

Sebastiani et al.[1] study this solution. Through terahertz spectroscopy ("terahertz" means frequencies between infrared and microwaves), two types of ion pairs are identified: solvent-separated pairs in which the solvated magnesium and sulfate ions are joined by electrostatic attraction but retain their separate solvation spheres, and solvent-shared pairs where some overlap between the solvation spheres occurs. In neither case are the ions directly bound to each other. But the same holds true in the solid salt equilibrated with water ($\ce{MgSO4•7H2O}$), in which the magnesium is coordinated to water molecules and the sulfate ions are hydrogen-bonded to water molecules. Thus we may regard the associated species as hydrated magnesium sulfate molecules in equilibrium with the dissociated ions.

Enter image description here Illustration from Ref. 1. Upper left: solvent-shared ion pair; lower left: solvent-separated ion pair.

Reference

  1. Federico Sebastiani, Ana Vila Verde, Matthias Heyden, Gerhard Schwaab, and Martina Havenith (2020). "Cooperativity and ion pairing in magnesium sulfate aqueous solutions from the dilute regime to the solubility limit". Phys. Chem. Chem. Phys., 22, 12140-12153.
$\endgroup$
0
$\begingroup$

Your thinking is correct. What happens at the extremes of conditions is not understood [or just oversimplified]. Here is a made up example using easy math. Lets give the Ka for a strong acid such as HCl a value of 100 [they probably are closer to 5-20].

HCl = H+ + Cl-; 100 = [H+][Cl-]/[HCl]. Lets make a 1M solution of NaCl

100 = [10^-7 * 1]/[HCl]; [HCl] =10^-9M, effectively zero and difficult to measure so Cl- is a weak base and HCl a strong acid.

Lets take a glance at a 10M solution of HCl; 100 = 10*10/[HCl]; [HCl]= 1M. It looks like an eleven M solution was needed to make a 10M acid level. If you sniff a concentrated HCl solution you would think that was close to reality. That something is going on can be seen with a good pH electrode.

A comment is made that there can not be an equilibrium with a solid until the solution is saturated. This ignores interface behavior and intermediate equilibria.

$\endgroup$

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service and acknowledge you have read our privacy policy.

Not the answer you're looking for? Browse other questions tagged or ask your own question.