Even in an acidic solution with a pH equal to 0, there are only 1 mol/L H3O+ ions, while there are approx. 54 mol/L water molecules

I find this sentence in my high school chemistry book to be confusing:


"Even in an acidic solution with a pH equal to 0, there are only 1 mol/L H3O+ ions, while there are approx. 54 mol/L water molecules."


How did they come to that conclusion and what were their calculations? I understand that the concentration of H3O+ ions is going to be 1 mol/L in a solution with a pH of 0, but how did they find that there are approximately 54 mol/L water molecules in this solution?




From the solution composition and density, it is easy to determine the water mass in $\pu{1 L}$ of the solution. As the molar mass of water is known, it is easy to calculate the molar amount of water in $\ce{1 L}$ too.


For $\mathrm{pH}=0$, instead of $[\ce{H3O+}]=\pu{1 mol/L}$, rather $a(\ce{H3O+})=1$. E.g. $\pu{2 M}\ \ce{H2SO4}$ has $\mathrm{pH}=0$ with just a small error, as activity coefficient of $\ce{H3O+}$ is near 0.5 ( and the 2nd hydrogen of $\ce{H2SO4}$ is mostly not dissociated, with $\ce{HSO4-}$ being a weak acid.)


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