A buffer is a solution that maintains a stable pH. Without buffering, your blood chemistry would be extremely susceptible to fluctuations in pH.

 

The pH of your blood is around 7.5. A rise in pH (alkalosis) to just 8.0 can be fatal, as is a drop in pH to 7.0 (acidosis). How does your blood maintain pH stability? The secret lies in the combination of weak acid, H₂CO₃, and a weak base, HCO₃^-1, that are naturally dissolved in your bloodstream. H₂CO₃, the same acid that is present in carbonated beverages, is formed as your blood picks up waste CO₂. (You might remember that carbon dioxide is a byproduct of cell respiration and combustion.) In solution, it disintegrates into a weak base, HCO₃^-1, sometimes called bicarbonate ion. Both H₂CO₃ and HCO₃^-1 are normally present in blood. In chemistry, such a combination is called a weak conjugate acid-base pair.

 

The presence of the weak acid (H₂CO₃) in blood is important. It absorbs base, and prevents pH of your blood from becoming too high. The weak base (HCO₃^-1) is equally important since it absorbs excess acid that might otherwise build up it your blood.

 

The purpose of this activity is to let you compare a nonbuffered system (plain system) to a buffer so that you will learn that buffers are highly efficient in keeping the pH of a solution from fluctuating.

magnetic stirrer with stir bar

100 mL beaker

labeled piper containing 1 M HCl

labeled pipette containing 1 M NaOH

100 mL graduated cylinder or 25 mL syringes

tonic water

1.0 M NaHCO₃: prepared by adding 84g NaHCO₃ (baking soda) to enough water to create 1000 mL of solution

1. Obtain a beaker (or disposable cup) and measure out 50.0mL of water. Set up a magnetic stirrer and place the beaker on it, setting the stirrer on low speed. We will keep the solution circulating, since your blood is also constantly in motion. Place a pH probe into the water and record the pH.

Now, add 5 drops of 1 M HCl to the beaker of water. HCl is a strong acid and even a tiny amount of it will cause the of pH of fall. Record the pH of the solution. If your blood pH dropped to the level, you would not survive long!

Now, add 5 drops of 1 M NaOH (a strong base) to the beaker of the water to neutralize the acid. The pH should rise to approximately the same pH as the water you started with. Add 5 more drops of NaOH. This simulates the condition of having too much base in you bloodstream. Record the pH.

Water does not contain the conjugate acid-base pair needed to maintain a constant pH. The pH of an unbuffered solutions changes quickly when acid or base are added.

2. Now, let's create a buffer that has a similar chemical combination to the agents that buffer your blood by combining 25mL of tonic water (carbonated beverages always contain this weak acid, H₂CO₃) and 25mL of a 1.0M solution of baking soda, HCO₃^-1.  Measure and record the initial pH of your buffer.

3. Will this buffer resist a change in pH better than the water did?  Add 5 drops of 0.1M HCl.  Record the pH.

Now, add 10 drops of the 0.1M NaOH.  Record the pH.

You may wish to investigate further by finding out how much of the strong acid or the strong base is needed to "break the buffer", or overload it to the point where it no longer maintains a stable pH.  You should now realize that a buffered solution can absorb the addition of a limited amount of acid or base without changing the pH radically.  Buffers do have boundaries, however.  They cannot absorb unlimited amounts of acid or base.