All right, it's time for me to remind myself of the chem 1 I used to tutor back about 100 years ago.
After looking at the Merck Index, I am afraid I have to contradict Slipknottin. It says the pH of 0.1 M solution of NaHCO3 is 8.3.
But let's be a little more general. There are bases and there are bases.
Strong bases, such as NaOH, dissociate more or less completely, generating a lot of OH- in solution, and driving pH way up. However, once dissociated, there's no NaOH left. So what? Well, say you add enough NaOH to have your solution at pH 8.2. What happens when you add acid? The pH drops immediately, i.e., the system isn't buffered. No reserve.
What about weaker bases, like NaHCO3 (sodium bicarbonate)? When you add it to water, it dissociates into Na+ and HCO3-. HCO3 exists in an equilibrium with H2CO3, such that HCO3- + H20 <-> H2CO3 + OH-. What the ultimate pH of the solution will be depends on the inclination of the reaction to go in one direction or the other. In the case of NaHCO3, equilibrium is reached at pH 8.3.
A very important point here is that the reaction does not completely go to H2CO3 + OH-. The means there will be HCO3- available to deal with added acids, i.e., the solution is buffered at about pH 8.3. That's what you're measuring when you measure alkalinity, how well the system can deal with added acid.
Other, weaker, bases will buffer the system at different pHs. For example, sodium benzoate will buffer a solution to about pH 8.0. So, even though it's a base, it will never be able to bring your pH to 8.2.
This probably makes no sense. I have condensed a few hours of lecture into a few paragraphs. The point is that you can add a base to your solution and it may not bring it to the right pH if you are using the wrong one. On the other hand, if you pick the right one, you can add all you want and it will just stabilize the pH exactly where you want it.
[but don't overdose bicarbonate or you will precipitate your Ca, but that's another story]
