aa and pKa-values

Non-professional visitors (i.e., lay people, high school and undergraduate students) or professionals from other fields should use this forum for general questions regarding molecular biology. No guarantee that they'll be answered but you can always try!

Moderators: mchlbrmn, r.rosati

aa and pKa-values

Postby Mari » Jun 01 2003 11:59 am

Hi,
maybe this is a really stupid question, but I don't quite understand it!
How will changing the pH (both ways, over and below 7) affect the pKa in an aquos environment, and in a hydrofobic environment? (For example for Lys and Asp) How much change in the pH is nesessary for change in pKa? What other factors may affect the pKa, other aa? Are there some general rules about this, or does it differ from protein to protein?

Thanks in advance :-)
Mari
Mari
technician-in-training
technician-in-training
 
Posts: 10
Joined: Jan 16 2003 4:23 am

Postby Ziggy » Jun 02 2003 5:34 am

The pKa is a constant, and will not change in aqueous solutions (irrespective of pH). pKa values can change in e.g. a hydrophobic environment, or the pKa can be affected by other amino acids nearby in a protein. I'm not sure whether there are general rules about this; it probably depends on the type of environment (e.g. are the nearby AA charged, hydrophobic, etc.)
Ziggy
Prolific Post-Master
Prolific Post-Master
 
Posts: 149
Joined: Mar 28 2003 7:08 am

Postby Edi » Jun 04 2003 10:20 am

Actually I feel confused about pH and pKa and raised the following questions:
1. What is the relationship between pH and pKa? Although I know that at 50% acid dissociation, pH will equal to pKa. Are they really the same thing? And what do pH and pKa mean if it is not in the situation of 50% dissociation?
2. A solution has its pH, then will it has its pKa value? In another hand, an amino acid residue has its pKa value, why can't we say it has a pH?
3. How can I know whether I can prepare a buffer at certain pH? because it seems that at certain pH, the solution may not act as a buffer as it has lost its buffering capacity.
4. Refering to AB's message, why pKa is a constant in hydrophilic environment whereas it varys in hydrophobic environment?
Hope someone can answer my silly questions. Thanks!
Edi
supertech
supertech
 
Posts: 74
Joined: May 03 2003 11:19 pm

Postby dpage » Jun 04 2003 10:58 am

pH is a function of the number of free protons (or hydrogen ions) in a solution - hence a solution has a pH, but it is nonsensical to talk about an aa having a pH.

pKa is defined for a molecule as being the pH at which half the molecules in that solution are in the N-H form and half in hte N- H+ form (if I remember correctly, though I stand to be corrected). Therefore molecules have a pKa value which is constant and does not change, as it is a property of the molecule, not of the solution it is in.
Hope this helps!!
David
Genomic Technology and Informatics
http://www.gti.ed.ac.uk
dpage
ModSquad
ModSquad
 
Posts: 111
Joined: Nov 07 2002 4:15 am
Location: Edinburgh, UK

Postby Ziggy » Jun 04 2003 11:33 am

[quote="Edi"]why pKa is a constant in hydrophilic environment whereas it varys in hydrophobic environment?
quote]

The whole theory about pH and pKa is based upon aqueous solutions. Acids and bases may behave very different if the amount of water molecules is very low, since the exchange of H+ or OH- may be different. As far as I know there is very little known about acid-base chemistry in for instance organic solvents, so the pKa of an amino acid is probably difficult to predict in hydrophobic environments.
Ziggy
Prolific Post-Master
Prolific Post-Master
 
Posts: 149
Joined: Mar 28 2003 7:08 am

Postby r.rosati » Jun 04 2003 11:57 am

Erm, this is going to be lengthy and tedious... be warned!

1.
in pH, "p" stands for "potence" (and H is obviously hydrogen ions, or protons). pH is, numerically, the inverse (positive <-> negative, is it "inverse" in english?) of the logarithm of the H+ ions molar concentration. (Or, approximately, the inverse of the exponent of 10 in the concentration).
Example: if H+ ions are 1*10^(-7) Molar, then pH is 7 (which is neutral pH). If we have less ions, then the solution is basic: let's say we have [H+] = 1*10^(-11) M, then pH is 11.
Similarly, pKa is the inverse potence of Ka, which is the dissociation constant for the equation [HA] -><- [A-] + [H+]. So if Ka is 1*10^(-5), then pKa is 5.

2.
An aminoacid has a dissociation constant, so it has a pKa. A solution has a proton concentration, so it has a pH.

3.
Buffer solutions basically work because the buffering agent has (at least) two forms, one of which can take H+ or OH- ions from the solution, thus becoming the other form.
Let's say, acetic acid (AcOH) has two forms, the acid AcOH and the acetate ion AcO- . One relationship between them is that [AcOH] = [AcO-] + [H+], this meaning that if you add protons in the solution (e.g. from a strong acid), then AcO- should take part of them to become AcOH (thus buffering the increase of H+). On the other hand, if you put OH- ions in the solution (e.g. from a strong base), then the OH- will tend to withdraw H+ (becoming H2O together), and AcOH will shift to AcO- making more protons available (thus buffering the decrease of H+). So, the acetic acid form buffers an addition of bases to the solution, while the acetate form buffers additions of acids.
As for this acid, Ka = [AcO-] [H+] / [AcOH],
when [H+] = Ka (that is, when pH = pKa), then [AcO-] = [AcOH].
This means that if pH is equal to pKa, there is an equal quantity of acetic acid and acetate ions.
This is the best condition for buffering activity, because if you add a strong acid, there is a good amount of AcO- to buffer it, and if you add a strong base, there is a good amount of AcOH to buffer it.
But as the two forms are strictly related to [H+], and thus to pH, any increase or decrease of one pH unity will change the concentration of the forms roughly by one potence of 10, or ten-fold!
So, at pH = pKa, a solution of acetic acid / acetate will be 50% acetic acid and 50% acetate... while at pH two units higher than pKa (thus more basic), you will have approximately 1% acetic acid and 99% acetate!
This last solution would be, of course, a risky buffer because it would only buffer acids: to buffer bases, you would need AcOH to be transformed, but AcOH is only 1% of total... so it won't last for long, and it won't work fully either! (As buffers are "weak" acids and bases, there [must] always remain some trace of either form).

4.
pKa, being related to Ka, measures the tendency of an acid to dissociate into ions. So in water, where ions are stable enough, acids have few problems in dissociating. But if you put an acid in a hydrophobic environment, this solution won't host the charged ions as well as water, and acids will tend to retain protons instead of dissociating - thus having a smaller dissociation constant (Ka) and a higher pKa.

*phew* i hope i didn't put mistakes in this!

Be well!

-Rob
r.rosati
ModSquad
ModSquad
 
Posts: 1917
Joined: Nov 04 2002 10:23 am
Location: Brazil

Postby Another God » Jun 05 2003 2:17 am

WeirdOmen wrote: pH is, numerically, the inverse (positive <-> negative, is it "inverse" in english?) of the logarithm of the H+ ions molar concentration. (Or, approximately, the inverse of the exponent of 10 in the concentration).

We normally just say 'pH is the negative log of [H]'

Thanks for a great summary though!
Prove me wrong. This is how I learn.
Another God
technician-in-training
technician-in-training
 
Posts: 5
Joined: Mar 29 2003 8:19 pm
Location: Sydney, Australia

Postby r.rosati » Jun 05 2003 7:51 am

Thanks A.G.!

One note, this sentence of mine
WeirdOmen wrote:This last solution would be, of course, a risky buffer because it would only buffer acids: to buffer bases, you would need AcOH to be transformed, but AcOH is only 1% of total... so it won't last for long, and it won't work fully either

is not fully correct: for additions of acids, there is indeed a lot of AcO- to counteract the increase of [H+], but due to the % difference in the two forms, the pH won't stay stable anyway.
r.rosati
ModSquad
ModSquad
 
Posts: 1917
Joined: Nov 04 2002 10:23 am
Location: Brazil


Return to Student Questions

Who is online

Users browsing this forum: No registered users and 1 guest