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pH of Non-Aqueous Solutions and Brønsted–Lowry Acid-Base Theory
While S.P.L. Sørensen was the first to propose the term pH in 1909, the Brønsted–Lowry acid base theory was developed (independently) by Danish Chemist Johannes Brønsted and Englishman Thomas Lowry.
A fascinating topic in it's own right, Brønsted–Lowry Acid Base theory broadly states that when an acid and a base react with each other, the acid forms its conjugate base (a base with a hydrogen [H+] ion added to it), and the base forms its conjugate acid (what's left - the base minus a hydrogen [H+] ion) by exchange of a proton (the hydrogen cation, or H+) - that is any compound that can transfer a proton to any other compound is an acid, and the compound that accepts the proton is a base.
This becomes interesting when you consider non-aqueous soltuions - consider our water molecules in standard motion:-
H2O + H2O ⇌ H3O+ + OH−
The hydrogen [H+] ion, or hydronium [H3O+] ion is a Brønsted–Lowry acid in aqueous solutions. The hydroxide [OH−] ion is the base.
We can draw analogies with a liquid ammonia reaction.
NH3 + NH3 ⇌ NH+
4 + NH−
2
The ammonium [NH+
4] ion plays the same role in liquid ammonia as does the hydronium [H3O+] ion in water and the amide [NH−
2] ion, is comparable to the hydroxide ion. Ammonium salts behave as acids, and amides behave as bases.
We can use the Brønsted-Lowry definitions to discuss acid-base reactions in any solvent.
If we expand our previous definition of pH: pH = - log10aH - then conversely we can say pOH[OH−]= −log[OH−].
Where pH + pOH= 14pH in aqueous solutions at 25°C
Our ammonium reaction wouldn't clarify for this definition as in non-aqueous solutions, the presence of H+ doesn't imply presence of OH-.
For example if a solution of dry sulphuric acid dissociation still takes place - H2SO4 ⇌ HSO4- + H+ where the H+ protonates sulphuric acid so 2H2SO4 ⇌ H3SO4+ + HSO4-
In this example the H3SO4+ is a analogous to our hydronium [H3O+] ion and conjugate base HSO4-. We can define pH for such solution, but we can't calculate pOH as pH + pOH= 14pH in aqueous solutions at 25°C.
Practicalities of measuring pH in non-aqueous solutions
There are other problems associated with measuring the pH of non-aqueous solutons.
Firstly a standard pH electrode may or may not work - the glass bulb of a pH electrode depends on a hydrated gel outer layer to sense the hydrogen activity (and pH) of a solution. The non-aqueous solution being measured may dehydrate or disrupt the hydrated layer, which results in the response becoming slow and precision becoming lost.
Another concern is the polymerised KCL salt bridge solution may not be miscible or may not dissolve into the sample being tested. This can cause a junction potential to develop and introduce bias to the results. In addition, the sample being measured may adversely affect the pH electrode both in terms of the construction of the electrode (epoxy, glass, PVC or PVDF) as well as the reference material and salt bridge.
With some non-aqueous solutions, the solution may have a high impedance and be a poor electrical conductor. We can over come this using a low resistance glass pH sensing bulb (such as the type used on the 9072-10B pH electrode) and using a high impedance pH controller.
Also - as the pH buffer solutions are aqueous solutions - these cannot be directly translated to the non-aqueous pH scale (pH + pOH= 14pH in aqueous solutions at 25°C). Non-aqueous buffer solutions are available, but these must relate to the solution being measured and often we'll find that measuring the pH of non-aqueous solutions becomes an indicative or relative measurement.
If we do begin to make measurements in non-aqueous solutions, then we may experience unstable readings, drift, long response times and measurement errors.
If you do have a requirement to measure non-aqueous solutions - then do let us know. A double junction refillable pH electrode is generally best suited and we will certainly try to assist you with the process.