# Why are strong acids and bases not suitable as primary standards?

This is one of the questions with which I have puzzled over, and can arrive at no definite conclusion. Why are strong acids or bases, such as $\ce{H2SO4}$,$\ce{ HNO3}$, $\ce{HCl}$, and $\ce{NaOH}$ not suitable primary standards?

How will their higher or lower pH affect the accuracy of the results of a titration?

-

• strong Bases in general are not suitable as they will react with $\ce{CO2}$ from the air. This is actually more problematic than hygroscopicity, as it means that you also cannot store diluted NaOH solution (openly) as standard.

• Instead, $\ce{Na2CO3}$ or $\ce{KHCO3}$ are much better alternatives (for strong acids).

• The acids are either a gas (HCl) or are produced from gas. I don't see how you can produce a conentration reliably enough for use as primary standard.

• even if you would manage to produce saturated HCl, it's much more difficult and less reliable than known alternatives (see below).

• but that wouldn't work e.g. for $\ce{H2SO4}$ as $\ce{SO3}$ is soluble in $\ce{H2SO4}$ (Oleum).

• $\ce{HNO3}$: redox reactions are not only not particularly useful, but actually harmful as they use up acid equivalents: $\ce{NO3- + 2e- + \mathbf{2 H+} -> NO2- + H2O}$

• also here, a solid alternative (KH-Phthalate) is available. If you don't have that, you may get away with oxalic acid which is still better than the acids suggested here.

-

There's nothing particular about strong acids or strong bases that disqualifies them from being a primary standard, but most of them are too unstable or difficult to work with to be useful as a primary standard. A primary standard should have these qualities:

1. High purity

2. Stability in presence of air

3. Absence of any water of hydration which might vary with changing humidity and temperature.

4. Cheap

5. Dissolves readily to produce stable solutions in solvent of choice

6. A larger rather than smaller molar mass

Now for the examples you gave:

$\ce{H2SO4}$ is hygroscopic, so it's concentration will change quite readily in air.

$\ce{HNO3}$ is very hygroscopic as well. The $\ce{NO3-}$ anions can also start redoxing other things which isn't particularly useful for titrations.

$\ce{HCl}$ is a gas, so it will be difficult to determine the concentration of the solutions (you'd have to compare it against another primary standard), and high-concentration solutions are unstable.

$\ce{NaOH}$ is hygroscopic as well.

As Georg pointed out below, one generally needs high-purity solid compounds so they can be weighed accurately (and in addition to the cons above, these substances are difficult to find/weigh in solid form). I hope this helps.

-
The only correct point is for NaOH. For HCl the reasoning is silly! diluted HCl solutions are absolutely stable, if only You avoid vaporizing the water in them. The HCl is not volatile at all in diluted solution. For all other the main point misses: one needs to weigh the primary titration standards reliably! This calls for stable solid compounds. –  Georg May 15 '12 at 21:00
@Georg: if you could be sure of the concentration in the first place you could also weigh a liquid... –  cbeleites May 15 '12 at 23:43
I edited to clarify the instability of HCl. Also, good point on the weighing--I added it to the answer. –  LeakyBattery May 16 '12 at 0:25