[Dimpl]

In an experiment, I measured the pH of 1M NaOH with a glass electrode pH meter, and it read 11.7. I then added distilled water to the solution, and the pH rose to 12.3. Is this expected from the equipment? I am confused :s

[magician4]

standard (!) glass electrodes at high pH will give you false readings with certain substances such as KOH , NaOH

in those cases, you have to use special equipment to measure pH=14 correctly


regards


Ingo

[vansh123]

The ph will decrease because the concentration of OH- will decrease after the dilution, hence the ph value will move towards 0.

[mjpam]

More to the point, though, is the fact that the pH system, rests on the fundamental assumption that the solution is dilute enough that interaction between solute ions is negligible, which allows the only relevant equilbrium to be that of the (de)protonation of the solute ions. Mathematically, this means that systems can be model in terms of a exponential equation, the logarithm of which is the linear pH scale.

As solutions become more concentrated, solute-solute interactions can no longer be ignored, and the deviations from linearity in the pH scale become more easily measurable. Hence, diluting a concentrated base soluton can result in a decreased pH.

[magician4]

Hence, diluting a concentrated base solution can result in a decreased pH.

though in principle all you've said is correct:
with respect to the original setup / question, I don't agree: though the effect you've described exists - no question there - this simply is not relevant for NaOH at or below 1 mol/L (i.e. what Dimple originally was asking about)

the what you've described, i.e. the inversion of the degree of effective dissociation for NaOH , will take place at or above ~ 4 mol/L  earliest.

(electric conductivity [mS/cm] vs. conc. [w/w] ; "Natronlauge" = aq. NaOH )
(from: link )
(data shown in this graphic are taken form: Weast, R.C, Astle, M.J. & Beyer, W.H. (edt.) 1985: CRC Handbook of Chemistry and Physics. 65th edition. Boca Raton (Florida): CRC Press. )

regards

Ingo

[mjpam]

My comment was deliberately vague because the caution about using the pH scale from concentrated acid solution such as >10 M H₂SO₄ much more readily apparent and applicable to the context I originally read it in (Olah and Klumpp's Superelectrophiles and Their Chemistry and Olah et al's Superacid Chemistry). Highly concentrated protic acids can undergo autoprotolysis and/or autodehydration and the concentration-activity of the H₃O⁺ is simply no longer an accurate measure of the ability of the soultion to protonate Lewis bases, due to the fact that H₃O⁺ is, in that context, one of the weakest protonators in the solution.

That said, I recall being cautioned about using the pH scale to describe of >1 M H⁺ or OH^-, which may have more to do with the divergence of the activity co-efficient from unity at high concentrations, but which I have no better way of explaining than waving my hands a mumbling "activity" under my breat.