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Typically students are not comfortable when asked to identify the most acidic protons in a molecule.
So here are some general guidelines of principles to look for the help you address the issue....

First, consider the general equation of a simple acid reaction:

equation defining Bronsted acidity

The more stable the conjugate base, A-, is then the more the equilibrium favours the product side.....
The more the equilibrium favours products, the more H+ there is....
The more H+ there is then the stonger H-A is as an acid....
So looking for factors that stabilise the conjugate base, A-, gives us a way to deduce acidity.

For a discussion of the factors that influence acidity, see the page on acidity

Here is a link to an acidity ladder.... a diagram that schematically shows the approximate pKa values of important systems for organic chemistry.

Consider the same basic equation as used for acidity, but think about the factors that affect the availability of the electrons.
Afterall, bases are electron pair donors.
The factors are the same ones that affect acidity. For example a more electronegative atom is a poorer electron donor and therefore a weaker base.

Structure and pKa

The information here is to help you decide which structure of an acid or base will dominate at a particular pH. Let's do a general case.
The equation for an acid is just HA = H+ + A- where = means equilibrium
pKa is defined as -log10 Ka where Ka = [H+][A-] / [HA].
From these expressions it is possible to derive the Henderson-Hasselbalch equation which is
pKa = pH + log [HA] / [A-]
This tells us that when the pH = pKa then log [HA] / [A-] = 0 therefore [HA] = [A-] ie equal amounts of the two forms.
If we make the solution more acidic, ie lower the pH, then pH < pKa and log [HA] / [A-] has to be > 0 so [HA] > [A-]. This makes sense as it tells us that the protonated form dominates in an acidic medium.
If instead we make the solution more basic, ie raise the pH, then pH > pKa and log [HA] / [A-] has to be < 0 so [HA] <[A-]. This makes sense as it tells us that the deprotonated form dominates in the basic medium.

These principles can be extended to poly acidic / basic systems (such as amino acids) by thinking of each pKa value in turn.

Lets look at an example.

To the right are the processes for the amino acid HISTIDINE, which has three acidic groups of pKa's 1.82 (carboxylic acid) 6.04 (pyrrole NH) and 9.17 (ammonium NH). Histidine can exist in the four forms shown, depending on the solution pH, from acidic pH (top) to basic pH. (bottom).
Starting from the top, we can imagine that as we add base, the most acidic proton is removed first (COOH), then the pyrrole NH then finally the amino NH. These takes us through each of the forms in turn.

At pH < 1.82, A is the dominant form.

In the range 1.82 < pH < 6.02 B is the dominant form.

In the range 6.02 < pH < 9.17 C is the dominant form, and when pH > 9.17, D is the major form in solution. OK?

© Dr. Ian Hunt, Department of Chemistry, University of Calgary