Acids are defined as compounds which can reversibly lose protons to the solution. The pH at which this occurs (the halfway point) is the pKa. Strong acids such as HCl will give up protons even at very low pH (hence low pKa) and weak acids will only give up protons if the pH is very high (i.e. the free proton concentration is very low). The pH at which this occurs can easily be above 7.
An important example is tyrosine, whose r-group is phenol; an acid with a pKa ~ 10. Conversely, bases can have pKa < 7.0. Can you think of an important biological BASE with a pKa less than 7.0?
Most biological acids, however, are weaker acids than HCl. The major class of biological acids is carboxylic acids. Because the difference in electronegativity between oxygen and hydrogen in a carboxylic acid is not as dramatic as it is with HCl, the tendency of a carboxylic acid to give up its proton is much less than that of HCl. However, carboxylic acids dissociate more readily than water due to the presence of two electronegative oxygens.
An acid's tendency to dissociate is a function of the strength of the acid and the pH of the solution. Strong acids can still dissociate when the pH is low, whereas weak acids cannot. The convention is to identify the pH at which the acid is half dissociated (e.g. half is protonated and half is deprotonated). This pH value is defined as the pKa of the acid in question. For example:
CH3COO (acetic acid) CH3COO- (acetate ion) + H+ ; pKa = 4.8;
meaning that, at pH = 4.8, half of the molecules are ionized (acetate) and half are not (acetic acid). The stronger the acid, the lower the pKa.