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I want to answer the question if a drug which is analysed for pH and pKa is ionised or not ionised in a cow stomach with a pH of 6.

The drug has a pH of 8,52 and a pKa of 8,78.

Since the pH is 8,52 I considered it as a base and in that case the drug would be almost completely ionized at a pH of 6, which means it is not absorbed through the stomach wall.

However, the drug has also some acidic side groups (carboxylic and phenols), which could mean that it works as an acid. In that case the equation turns around and it would mean that it is unionized at a pH of 6, which means it is absorbed through the stomach wall.

Is there a way that I can tell if the drug, based on just the pH and pKa, is considered an acid or a base?

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closed as unclear what you're asking by Mithoron, A.K., Todd Minehardt, Jon Custer, Tyberius Mar 22 at 15:31

Please clarify your specific problem or add additional details to highlight exactly what you need. As it's currently written, it’s hard to tell exactly what you're asking. See the How to Ask page for help clarifying this question. If this question can be reworded to fit the rules in the help center, please edit the question.

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    $\begingroup$ If it has multiple functional groups that can act as acid or base, it should have multiple pKa values. I don't understand what you mean when you say "the drug has a pH of...". A solution has a pH, a molecule has a pKa. The pH of a solution containing the drug (only, nothing else) would depend on the concentration of the drug. Does the drug in the solid form contain sodium or chloride counterions? Could you give an example structure of the kind of drug you are asking about? $\endgroup$ – Karsten Theis Mar 21 at 14:48
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Fist things first: a substance with multiple basic and acidic dissociation sites has multiple $pK_a$'s, not only one.

Normally, for basic sites the $pK_a$'s of the corresponding conjugated acids are given.
The highest $pK_a$ is the one of the 'most basic' site, because its conjugated acid is the weakest.

For acidic sites, the $pK_a$'s of the acids themselves are given. The lowest $pK_a$ is the one of the 'most acidic' site, i.e. the strongest acid.

Under some assumptions, you can study the behaviour of a molecule looking only at these two $pK_a$'s.

3-hydroxypyridine has 1 basic $pK_a$ (for the pyridinium $NH^+$, 4.79) and 1 acidic $pK_a$ (for the $OH$, not far from 8.75).
So it will have mostly charge +1 at $pH < 4.79$, mostly charge 0 (and no charges on any atoms) at $4.79 < pH < 8.75$, mostly charge -1 at $pH > 8.75$.

Note however that the most basic $pK_a$ is not always lower than the most acidic $pK_a$.

E.g. glycine (2-amino-acetic acid) has 1 acidic $pK_a$ (for the $COOH$, 2.35) and 1 basic $pK_a$ (for the $NH_3^+$, 9.78).
So it will have mostly charge +1 at $pH < 2.35$, mostly charge 0 (but zwitterionic, i.e. with 1 positive charge and 1 negative charge) at $2.35 < pH < 9.78$, mostly charge -1 at $pH > 9.78$.

So to answer your initial question, i.e. how to calculate the fraction with charge 0 of your drug at a specific pH, you need to know at least the most basic $pK_{a,MB}$ and most acidic $pK_{a,MA}$.
Obviously if there is no acidic or basic site, the molecule is always neutral, and if there is only one acidic or only one basic site you can ignore the other one.
In practice, for the calculation, if there are no acidic $pK_a$'s, you can set the most acidic $pK_a$ is 100; and if there are no basic $pK_a$'s, you can set the most basic $pK_a$ is -100.

Then at a given pH the fraction with charge 0 is:

$$f_{z=0} = \frac 1 {1+10^{-pK_{a,MA}+pH}+10^{+pK_{a,MB}-pH}}, if \ pK_{a,MA} > pK_{a,MB}$$

or:

$$f_{z=0} = \frac 1 {1+10^{+pK_{a,MA}-pH}+10^{-pK_{a,MB}+pH}}, if \ pK_{a,MA} > pK_{a,MB}$$

For instance for 3-hydroxypyridine, at pH 6:

$$f_{z=0} = \frac 1 {1+10^{-8.75+6}+10^{+4.79-6}} \approx 94 \%$$

At pH 3 (like in a human stomach) things would be quite different:

$$f_{z=0} = \frac 1 {1+10^{-8.75+3}+10^{+4.79-3}} \approx 1.6 \%$$

For glycine, at pH 6:

$$f_{z=0} = \frac 1 {1+10^{+2.35-6}+10^{-9.78+6}} \approx 100 \%$$

At pH 3:

$$f_{z=0} = \frac 1 {1+10^{+2.35-3}+10^{-9.78+3}} \approx 82 \%$$

As for your other question, whether a substance is an acid or a base, well as you can see from the explanation here, it depends on the presence of acidic and basic dissociation sites, and the pH we're considering as 'reference'.

Normally the reference physiological pH is $7.4$ (blood).

A substance that has mostly positive charge at the reference pH is usually considered a 'base', and if negative, an 'acid'. But you see how fluid the situation can be...

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