[orgo814]

I was supposed to sketch the titration curve for H3N+CH2COOH. This is a diprotic acid so I looked up appropriate pKa values... 4.76 for acetic group and 9.25 for ammonium. However, when I checked my answer, the pKa in the books titration curve is around 2.3-2.4 while the ammonia is around 10. What's the cause of this? I'm so confused

[rwiew]

Ok, so comparing your aminoacid (glycine) to acetic acid is correct, but I would like to change from comparing to ammonium to N-methylammonium (protonated methylamine) - this is because glycine has an alkyl group on the nitrogen, pKa for N-methylammonium is 10.64. So in glycine the carboxylic acid is less acidic than in acetic acid and the amino group is less basic than in methylamine. Clearly the two groups are influencing each other, what can you tell me about the electron withdrawing/donating abilities of the two. Consider what forms (deprotonated/protonated) they are at the points in the titration curve we're talking about as well.

[orgo814]

Well the glycine is more acidic since in the curve they gave me the pK was 2.35ish. And the pK for the NH3+ group was higher. I know that once the first group dissociates it probably makes the overall molecule more basic but I just wish I had something more logical to determine the pK besides just saying "ok, this NH3+ is also on the molecule so the carboxylic acid is obv more acidic etc)

[Borek]

9.78, 2.35

[rwiew]

I just wish I had something more logical to determine the pK besides just saying "ok, this NH3+ is also on the molecule so the carboxylic acid is obv more acidic etc)

Well you do, an experiment. Or you could do some computation if you want to determine this theoretically. I vaguely recall someone came up with an approximate system of calculating pKa's based on the inductive/resonance contributions of neighboring groups, can't remember the name though right now. Or maybe that was just NMR shift calculation that I'm thinking about...

[Yggdrasil]

Well the glycine is more acidic since in the curve they gave me the pK was 2.35ish. And the pK for the NH3+ group was higher. I know that once the first group dissociates it probably makes the overall molecule more basic but I just wish I had something more logical to determine the pK besides just saying "ok, this NH3+ is also on the molecule so the carboxylic acid is obv more acidic etc)

A good way to assess the relative acidity/basicity of two compounds is to look at what factors may be stabilizing the conjugate base form of the group in question.  Here, let's compare the basicity of the carboxylic acid in glycine (NH3-CH2-COO) versus acetic acid (CH3-CH2-COOH).  Acetic acid has a pKa of 4.76.  How would we expect the addition of the amino group in glycine to affect the acidity of the carboxyl group?  Well, let's look at the factors involved in stabilizing the conjugate base of acetic acid (acetate, CH3-CH2-COO^-):

1) Resonance.  The oxygen anion in acetate is stabilized by resonance, which delocalizes its negative charge across the two oxygens in the carboxyl group.  This helps explain why the OH in a carboxyl group is so much more acidic than a normal OH group (pKa ~ 14).  Addition of the amino group in glycine, however, does not affect the resonance stabilization.

2) Charge. Deprotonation of acetic acid creates an overall negative charge.  Deprotonation of glycine, however, takes a molecule with a net positive charge (NH3-CH2-COOH) to a molecule with an overall neutral charge.  Thus we expect the glycine with a deprotonated carboxyl to be more stable than the deprotonated acetic acid. 

3) Atomic radius, electronegativity.  The atomic radius and electronegativity of the ionizable group influences its stability (in general, negative charges and lone pairs are more stable on atoms that have a larger radius.  If the radii are similar, then atoms with a higher electronegativity are more stable).  In both cases, the ionizable atom is an oxygen, so there is no difference in this category between acetic acid and glycine.

4) Inductive effect.  The NH3⁺ in glycine is slightly electron withdrawing than the CH3 group in acetic acid (because nitrogen is more electronegative than carbon).  Thus, the amino group draws electron density away from the negatively-charged carboxyl in its deprotonated form, stabilizing the conjugate base.

Thus, because of the effects of the amino group on the overall charge of the molecule and its contribution through the inductive effect, we'd expect the conjugate base of glycine to be more stable than the conjugate base of acetic acid.  Because the conjugate base is more stable (i.e. easier to form), we'd expect the carboxyl group of glycine to exhibit a lower pKa than acetic acid.

Here's a nice tutorial that explains this line of reasoning for looking at the relative basicity of various groups: http://www.chem.ucla.edu/harding/tutorials/acids_and_bases/mol_str.pdf

[Babcock_Hall]

It might be useful to compare the pK_a values of glycine with those of the dipeptide, glycylglycine (Gly-Gly), which are 3.06 and 8.13.  The carboxylic acid is a weaker acid in the dipeptide.  Thinking about electrostatics might help.