Drug Solubility

[FastLane051]

Advertisement - Members don't see this ad

So, this is my first year in pharmacy school and something that was mentioned the other day confused me.
The professor mentioned that the drug solubility had to be in equilibrium between its ionic and non-polar states.

How does this make sense?
I thought a drug had to be non-polar for absorption through the epithelial walls. So would it non make sense for aspirin, a drug with a PKa of 3.6, to be in an acidic environment where it remains protonated to be absorbed as a non-polar molecule?

 

[owlegrad]

So, this is my first year in pharmacy school and something that was mentioned the other day confused me.
The professor mentioned that the drug solubility had to be in equilibrium between its ionic and non-polar states.

How does this make sense?
I thought a drug had to be non-polar for absorption through the epithelial walls. So would it non make sense for aspirin, a drug with a PKa of 3.6, to be in an acidic environment where it remains protonated to be absorbed as a non-polar molecule?

Sounds like a great question for your professor.

 

[Farmercyst]

So, this is my first year in pharmacy school and something that was mentioned the other day confused me.
The professor mentioned that the drug solubility had to be in equilibrium between its ionic and non-polar states.

How does this make sense?
I thought a drug had to be non-polar for absorption through the epithelial walls. So would it non make sense for aspirin, a drug with a PKa of 3.6, to be in an acidic environment where it remains protonated to be absorbed as a non-polar molecule?

Bare in mind that pH changes throughout the GI tract. The pKa determines both whether and where a drug will be absorbed. Also keep in mind that pH=pKa x log of A-/HA therefore there is a ratio, however small, of protonated vs unprotonated at any pH. Whether it's 1:1 at 3.6 or 1:100 at 5.6 (jejunum) or 100:1 at 1.6 (stomach). Hence, aspirin is more readily absorbed in the small intestine (HA 100:1) than in the stomach (A- 100:1). Finally make sure you're not confusing the definition of equilibrium. Equilibrium doesn't mean equal parts (1:1) but rather constancy of change (ratio is maintained at the same level) therefore if some aspirin is absorbed, more is converted from A- to HA to keep the ratio constant)

Hopefully this hasn't been more or less than you were looking for.

 

[HaloKing]

sounds like a pre-pharmacy question.. nicely answered cyst

 

[Farcus]

mastering and understanding the concept of Henderson-Hasselbalch should be a pre-requisite for pharmacy

but the poster above posted is pretty on point, and you should drill that concept of equilibrium into your head of it not being 1:1 but a dynamic change to maintain some ratio although i'm have trouble understanding your question.

 

[avuong]

Dear OP,

drug must be non-ionic, non-charged to be absorbed; nonpolar lipophilic even better . HA (zero charge) will be absorbed, A-(conjugate-base, negative charged) will not be absorbed.

HA <---> [A-] + [H+]

ka = [A-][H+] / [HA]
-log Ka = -log ([H+][A-]/[HA])
pKa = pH - log [A-]/[HA]

http://chemistry.about.com/od/acidsbase1/a/hendersonhasselbalch.htm

notice:

eg. pKa of 3.6, pH of 5.6
pKa = pH - log [A-]/[HA]
3.6 = 5.6 - log [100]/[1]
3.6 = 5.6 - 2

mathematically, there must be 100 fold more A- than HA.

but why? -- because actually, physically, in the real world, when conditions are basic (high OH-, low H+), you can see with the naked eye the reaction actually drives toward the right to replenish the low H+

eg.
[HA] <--> [A-] + [H+]
under basic conditions, will want to drive toward regenerating more H+
youll end up with:
[HA] ---> [A-] + [H+]
(100 fold more A- than HA)



Conversely,

eg. pH environment of 1.6, pKa 3.6
pKa = pH - log [A-]/[HA]
3.6 = 1.6 - log [1]/[100]
3.6 = 1.6 - (- 2)

mathematically you will find 100 fold more HA than A-.

in the real world, physically, chemically, visually, you will see the reaction actually driving toward the left. Excess H+
HA <--> A- + H+
drives to
HA <--- A- + H+
In a more acidic environment than pKa (excess of H+), reaction will drive toward reactant; left side.

-----
For review:
[HA] --> [A-] + [H+]
you want to be in a low pH (excess H+) to drive the reaction toward the protonated uncharged HA.

for drugs like:
[H20] + --> [BH+] + [OH-]
you actually want the pH to be high (more OH-) to drive the reaction toward the unprotonated, uncharged B.
eg. NH3 <---- NH4+. You want there to be a surplus of OH, deficit of H+ so the H+ will break off the NH4.
---
Therefore
You can think in terms of raw numbers and use the mathematics of Henderson-Hasselbach, or you can think physically in terms of La Chatelier.
Personally, I prefer the physical mode because I can actually see the reactions with my naked eye.

Best of Luck!

 

[avuong]

So would it make sense for aspirin, a drug with a PKa of 3.6, to be in an acidic environment where it remains protonated to be absorbed as a non-polar molecule?

Yes, it would make sense because low pH drives the reaction toward the?:

[HA] <--> [A-] + [H+]

<-------- (left side)

[HA] form, uncharged, non-ionized. non-ionized forms can be absorbed through the fatty bilayer.

Best of luck.