# Why is water not part of the equilibrium constant?

In the acid base reaction $$\ce{NH4+ + H2O <-> NH3 + H3O+}$$

the acidity constant, which is a relation between concentrations is given by $$K_{\mathrm{a}}=\frac {\ce{[NH3]} \cdot \ce{[H3O+]}}{\ce{[NH4+]} \cdot \ce{[H2O]}}$$

Since $\ce{[H2O]}$ is constant, it can be included inside the constant to obtain a new $K_\mathrm{a}$ which is $6.3 \times 10^{-10}$.

(Being a constant allows us to relate the molar concentration to the density of water.)

And so, even though some quantity of water reacts with $\ce{NH4+}$, but the volume of water varies in a way that $\ce{[H2O]}$ would still be constant.

What I don't understand is why the molar concentration of water is constant. Why isn't it handled the same way as the other reactant? I understand that if $\ce{H2O}$ was the only liquid and the rest of the reactants and products were of different state (gaseous or solid), then water would exist alone in the beaker and its concentration would therefore remain constant, but this is not the case.

That was about heterogeneousness. Trying another guess; in the reaction of esterification $$\ce{{alcohol} + {carboxylic acid} <-> {ester} + H2O}$$ $\ce{H2O}$ is considered as a product which concentration is not constant, so it is counted explicitly in the equilibrium constant.

Could the answer be that $\ce{NH4+}$, $\ce{NH3}$ and $\ce{H3O+}$ are dissolved in water, while the alcohol, the carboxylic acid, and the ester are not?

I'm looking for an answer that clearly shows which volume is considered (water or the solution), because I'm confused as to why something would or wouldn't be included in the equilibrium constant.

• To get your chemistry to appear pretty, you need to enclose it in $_$. So $\ce{NH4+}$ becomes $\ce{NH4+}$. Do not use MathJax/LaTeX in question titles as it messes with searching. Dec 10, 2015 at 23:46
• In general the concentration of water is assumed to be a constant in aqueous solution. That isn't absolutely true of course since for example the density of water changes with temperature.
– MaxW
Dec 10, 2015 at 23:55
• related, but not duplicate: chemistry.stackexchange.com/questions/9847/… Dec 11, 2015 at 14:42

There is no IUPAC definition of $K_a$, however, many reputable texts like Levine's Physical Chemistry define $K_a$ as:

$$\frac {a(\ce{A-})a(\ce{H3O+})}{ a(\ce{HA})a(\ce{H2O})}$$

where "a(X)" is the activity of species "X".

Then various approximations can be made, such as approximating $a(\ce{H2O}) = 1$ and approximating the activity of the solutes as the concentration of the solutes.

These approximations are only reasonable in dilute solutions.

The concentration of water in water (sounds weird but this is just liquid water) is found as follows:$$[\ce{H2O}]= \mathrm{\frac{1.0\frac{g}{cm^3}}{18.0148\frac{g}{mol}}}=55.5\mathrm{\frac{mol}{L}}$$

Compare this with the usual concentrations of reactants and products which is much smaller than that. Thus, the change in concentration of water throughout reactions which take place in water is considered to be constant.

The above is a very good approximation. While it is true that water's density varies some with temperature, at $98~\mathrm{^\circ C}$ the density of water is still $0.96\mathrm{\frac{g}{cm^3}}$, so the concentration of water will still be much much more than that of any of the reactants.

So, as you speculated, whether or not water is considered in the equilibrium constant entirely depends on whether or not water is a product and if water is the solvent.

If water is the solvent, it would be very normal to ignore the water which is produced or the water which might act as a reactant.

Similarly, any reaction which takes place by solvolysis will exclude the solvent from the equilibrium constant.

Also, you say that if water were alone in a beaker, then obviously all you have is water so the concentration is constant. Consider that water auto-ionizes by:$$\ce{2H2O <=> H3O+ + OH-}$$In this case, even though we know that we don't have just water, we know that the concentration of the other things is tiny compared to the concentration of water, so we just ignore the other stuff and drink the water.

• Pro tip: Type entire equations with a single \ce{...}. $\ce{H2 + Cl2 -> 2 HCl}$ and $\ce{AB + C <=> A + B + C <=>> A + BC}$ give $\ce{H2 + Cl2 -> 2 HCl}$ and $\ce{AB + C <=> A + B + C <=>> A + BC}$, respectively ;)
– Jan
Dec 11, 2015 at 2:37
• So not only water must be a solvent, but also its quantity must be relatively large comparing to other soluble substances of the reaction in order to consider its concentration as constant and attaining the above value of 55.5 mol per liter. Agree? Dec 14, 2015 at 5:45
• Or I guess the solubility (which is the maximum amount of dissolved matter in a given temperature) which would be small will keep the concentration of other substances small relative to water, that's why we won't bother by comparing quantities put into reaction. Dec 14, 2015 at 5:56