AlgebraLAB
 
 
Site Navigation
Site Directions
Search AlgebraLAB
Activities
Career Profiles
Glossary
Lessons
Reading Comprehension Passages
Practice Exercises
StudyAids: Recipes
Word Problems
Project History
Developers
Project Team






Buffered Solutions
Background Information:
 
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.
 
Equipment:
 
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
 
Procedure: Let's simulate what would happen to your blood pH if the H2CO3 and HCO3-1 were not present.
 
  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.
 
  1. 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.
 
  1. 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.
 
 
Extension: Investigate the terms acidosis and alkalosis, by conducting an internet search. What conditions would cause your blood pH to fluctuate to a fatal degree?

Examples
Data Analysis
Example How much [H3O+] is present in a buffered solution that has a pH of 5.5?
What is your answer?
 
Example If the pH decreases by one unit, to 4.5, by what factor does the [H3O+] change?
What is your answer?
 
Example If the pH decreases by two units, to 3.5, by what factor does the [H3O+] change?
What is your answer?
 



E Saylor

Show Related AlgebraLab Documents


  Return to STEM Sites AlgebraLAB
Project Manager
   Catharine H. Colwell
Application Programmers
   Jeremy R. Blawn
   Mark Acton
Copyright © 2003-2017
All rights reserved.