Acid/Base dissociation constants relationship

From what I understand, the concentration of hydronium and hydroxide molecules is constant in pure water (and equals $10^{-7}\ \text{M}$, which is measured experimentally). 


What I don't understand is why this remains true for an aqueous solution with an acid or base in it. Shouldn't it affect these concentrations?


When you write the equilibrium reactions for a weak/weak acid/base pairs (one with its conjugate), and get that $K\_\mathrm a\cdot K\_\mathrm b = [\ce{H+}][\ce{OH- }]$, what's the explanation for it? Why does their multiplication (either as pure water or in a acidic/basic solution) remain constant?




The concentrations do change in the presence of acid or base and are no longer $10^{-7}$M.


Instead, [H3O+][OH-] $= 10^{-14}M^2$


This comes from: 


$\ce{2H2O <=> H3O+ + OH-}$


$K = \frac{a(\ce{H3O+})a(\ce{OH-})}{(a(\ce{H2O}))^2}$


Where "a(x)" is activity of species "x". 


Then, approximating activity of $\ce{H2O}$ as 1, and the activity of the ions as the concentration of the ions there is a constant $K\_w$, such that:


[H3O+][OH-] $= 10^{-14}M^2 = K\_w$


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