$\ce{NaClO_{(s)} + H2O <=> NaOH_{(aq)} + HClO_{(aq)}}$

This reaction is often used in swimming pools as an alternative to directly chlorinating the water.

I read that the water must be kept slightly acidic. Otherwise, $\ce{OH-}$ ions in the alkaline solution would react with the $\ce{H+}$ ions in the chloric (I) acid to form water. Consequently, $\ce{ClO-}$ ions are removed and the concentration of $\ce{H+}$ ions decreases.

So far this all seems fairly reasonable.

However, my textbook then said this:

In alkaline solution, the equilibrium moves to the left.

This contradicts what I have read online and elsewhere. Surely if the concentration of hydrogen ions decreases, the equilibrium would shift to the right to increase the concentration of the products and thus the hydrogen ions?

Thanks in advance.

  • $\begingroup$ Why did you introduce hydrogen ions into this equilibrium? They're not present. Therefore, if you increase the concentration of hydroxide, the equilibrium shifts left. $\endgroup$ – Zhe Jun 1 '18 at 15:18
  • $\begingroup$ Reading NaOH(aq) makes me cringe. That compound disassociates completely. $\endgroup$ – Stian Yttervik Jun 1 '18 at 20:22
  • $\begingroup$ @StianYttervik That's the whole point of using the aqueous phase notation... $\endgroup$ – Zhe Jun 3 '18 at 0:36
  • $\begingroup$ @Zhe it is redundant for those kinds of substances. Use it for weak acids or substances that doesn't dissociate completely - then the notation carries meaning. $\endgroup$ – Stian Yttervik Jun 3 '18 at 12:09
  • $\begingroup$ It is preciously with weak acids where the meaning is ambiguous... @StianYttervik $\endgroup$ – Zhe Jun 3 '18 at 14:11

This question is similar to a problem in nuclear reactor chemistry, I will get onto that soon.

If we were to combine chlorine and water then we will have an equilibrium.

$\ce{Cl2 + H2O -> HCl + HOCl}$

The HCl is a strong acid while the HOCl (hypochlorous acid) is a weak acid. By careful adjustment of the pH it is possible to determine where the equilibrium is. If we were to make the pool too acidic then the oxidant in the water would be lost and the chlorine would enter the air as Cl2. This is a waste of chlorine and also is harmful to swimmers and even sometimes the strucuture of the building.

I assume that there is some chloride present in commerial sodium hypochlorite solution as the classic way to make this chemical is to combine sodium hydroxide and chlorine. The compound is formed by the reaction below.

$\ce{2NaOH + Cl2 -> NaCl + NaOCl}$

Even if we were to have a close to chloride free solution of sodium hypochlorite then I am sure that the first equilbrium I mentioned would be important. It is important to keep in mind that hypochlorite can disproprotionate. It can be converted into chloride and chlorate. If you combine chlorine and cold dilute sodium hydroxide then you make sodium hypochlorite while if you were to use hot concentrated sodium hydroxide then you will make sodium chlorate.

$\ce{3HOCl -> 2Cl^- + ClO3^- + 3H^+}$

As hydrochloric and chloric acids are stronger than HOCl as Bronsted acids, I think that by making the system very alkaline you would favour the conversion by the disproprotionation. What I know about explosives / prolellants and pesticide chemistry is that while going along the series

$\ce{OCl- , ClO2^-, ClO3^3}$ and $\ce{ClO4^-}$

That while in terms of thermodynamics the chlorine species become better oxidants, the rate at which they can act as oxidants often becomes slower. Thus in a pool if we change from sodium hydrochloite to sodium chlorate then it would be less able to kill the germs in the pool.

Also if we made the pool too alkaline then the water would be too corrosive.

In a reactor containment with a water on the bottom such as many BWR designs the iodine chemistry is similar. When the water is acidic then the iodine will form volatile I2. But when the water is alkaline then the elemental iodine will be converted into iodide and iodate by similar reactions.


The reaction as shown goes only a little to the right. The $\ce{HClO}$ shown is poorly ionized; on the other hand, the $\ce{NaOH}$ is completely ionized. How far the reaction goes to the right will be indicated by the pH, which is probably about 10, considering that sodium carbonate is added to commercial bleach to raise the pH to 11 for greater shelf life.

When chlorinating a swimming pool, long shelf life (in the pool) is not desired. You want some more activity, even if at the expense of having some of the $\ce{NaClO}$ divert to other products.

There are two variables in the swimming pool:

  1. the chlorination amount
  2. the activity, which will be controlled by the pH.

So you put in the chlorine, then dial up the best pH. As far as moving the equilibrium, perhaps it is better to show the reaction as, $$\ce{ClO- + H2O -> HClO + HO-}$$

The original sodium hypochlorite will be totally ionized; the hydrolysis reaction is the $\ce{ClO-}$ ion attracting a proton from water to form somewhat nonionized HClO and some $\ce{OH-}$ ions. Then the prospect of adding more $\ce{OH-}$ ions is easily seen to move the equilibrium back toward the left, generating more $\ce{ClO-}$ ions. Likewise, adding acid generates more $\ce{HClO}$ in the solution.


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