Some solutions, like concentrated HCl, act as buffers even though they lack the buffering component. What is the cause for such a behaviour?


I do wonder what your question is, but I have a hunch. Assuming that what you’re asking is essentially:

In a concentrated $\ce{HCl}$ solution, the $\mathrm{pH}$ will more or less be the same value even if we add external $\ce{H+}$ or $\ce{OH-}$. Why?

Then there are multiple effects at work:

  • Aquaeous $\ce{HCl}$ contains a buffering component — the $\ce{H3O+}$ (or closest relative) ion. It is the most acidic ion that can be present is solution, so it ‘buffers down’ the acidity of said solution to what the corresponding $\ce{H3O+}$ acidity would be.

  • $\ce{HCl}$ in itself is an acid, but at the same time $\ce{Cl-}$ is also a (very weak) base. For high total $\ce{HCl}$ concentrations, the ratio of $\ce{HCl}$ to $\ce{Cl-}$ is at equilibrium. If we add additional $\ce{H+}$ to the mixture, we are effectively shifting that equilibrium back to the $\ce{HCl}$ side, reducing the proton count and thus buffering the $\mathrm{pH}$. Therefore also, if we add a proton scavenger (colloquially also known as ‘base’), then that will shift the equilibrium to the $\ce{Cl-}$ side, and as long as there is still enough $\ce{HCl}$ around, the total proton concentration won’t change notably.

In essence, everything that can act as a Brønsted acid or base will buffer at a specific $\mathrm{pH}$ value, even if that $\mathrm{pH}$ value is far from what one would typically call ‘buffered’.


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