What do buffering systems do

In this example we will continue to use the hydrofluoric acid buffer. We will discuss the process for preparing a buffer of HF at a pH of 3. This is simply the ratio of the concentrations of conjugate base and conjugate acid we will need in our solution. How much Sodium Fluoride would we need to add in order to create a buffer at said pH 3. From a table of molar masses, such as a periodic table, we can calculate the molar mass of NaF to be equal to Using this information, we can calculate the amount of F - we need to add.

We could use ICE tables to calculate the concentration of F - from HF dissociation, but, since K a is so small, we can approximate that virtually all of the HF will remain undissociated, so the amount of F - in the solution from HF dissociation will be negligible.

Thus, the [HF] is about 1 M and the [F - ] is close to 0. This will be especially true once we have added more F - , the addition of which will even further suppress the dissociation of HF.

Thus, [F - ] should be about 0. For mL of solution, then, we will want to add 0. Since we are adding NaF as our source of F - , and since NaF completely dissociates in water, we need 0. Thus, 0. Such dilute solutions are rarely used as buffers, however. Recall that the amount of F - in the solution is 0.

Let's double check the pH using the Henderson-Hasselbalch Approximation , but using moles instead of concentrations:. Now let's see what happens when we add a small amount of strong acid, such as HCl. The Cl - is the conjugate base of a strong acid so is inert and doesn't affect pH, and we can just ignore it. Proteins are made up of amino acids, which contain positively charged amino groups and negatively charged carboxyl groups.

The charged regions of these molecules can bind hydrogen and hydroxyl ions, and thus function as buffers. Buffering by proteins accounts for two-thirds of the buffering power of the blood and most of the buffering within cells.

Hemoglobin is the principal protein inside of red blood cells and accounts for one-third of the mass of the cell. During the conversion of CO 2 into bicarbonate, hydrogen ions liberated in the reaction are buffered by hemoglobin, which is reduced by the dissociation of oxygen. This buffering helps maintain normal pH. The process is reversed in the pulmonary capillaries to re-form CO 2 , which then can diffuse into the air sacs to be exhaled into the atmosphere.

This process is discussed in detail in the chapter on the respiratory system. Acids and bases are still present, but they hold onto the ions.

The bicarbonate-carbonic acid buffer works in a fashion similar to phosphate buffers. The bicarbonate is regulated in the blood by sodium, as are the phosphate ions. When carbonic acid comes into contact with a strong base, such as NaOH, bicarbonate and water are formed. As with the phosphate buffer, a weak acid or weak base captures the free ions, and a significant change in pH is prevented.

Bicarbonate ions and carbonic acid are present in the blood in a ratio if the blood pH is within the normal range. With 20 times more bicarbonate than carbonic acid, this capture system is most efficient at buffering changes that would make the blood more acidic. Carbonic acid levels in the blood are controlled by the expiration of CO 2 through the lungs. In red blood cells, carbonic anhydrase forces the dissociation of the acid, rendering the blood less acidic.

Because of this acid dissociation, CO 2 is exhaled see equations above. The level of bicarbonate in the blood is controlled through the renal system, where bicarbonate ions in the renal filtrate are conserved and passed back into the blood. However, the bicarbonate buffer is the primary buffering system of the IF surrounding the cells in tissues throughout the body. The respiratory system contributes to the balance of acids and bases in the body by regulating the blood levels of carbonic acid Figure CO 2 in the blood readily reacts with water to form carbonic acid, and the levels of CO 2 and carbonic acid in the blood are in equilibrium.

When the CO 2 level in the blood rises as it does when you hold your breath , the excess CO 2 reacts with water to form additional carbonic acid, lowering blood pH. The loss of CO 2 from the body reduces blood levels of carbonic acid and thereby adjusts the pH upward, toward normal levels.

As you might have surmised, this process also works in the opposite direction. Excessive deep and rapid breathing as in hyperventilation rids the blood of CO 2 and reduces the level of carbonic acid, making the blood too alkaline.

This brief alkalosis can be remedied by rebreathing air that has been exhaled into a paper bag. Rebreathing exhaled air will rapidly bring blood pH down toward normal. Minor adjustments in breathing are usually sufficient to adjust the pH of the blood by changing how much CO 2 is exhaled.

This situation is common if you are exercising strenuously over a period of time. To keep up the necessary energy production, you would produce excess CO 2 and lactic acid if exercising beyond your aerobic threshold. In order to balance the increased acid production, the respiration rate goes up to remove the CO 2.

This helps to keep you from developing acidosis. The body regulates the respiratory rate by the use of chemoreceptors, which primarily use CO 2 as a signal. Acid buffer solutions have a pH less than 7. Commonly used acidic buffer solutions are a mixture of ethanoic acid and sodium ethanoate in solution, which have a pH of 4. You can change the pH of the buffer solution by changing the ratio of acid to salt, or by choosing a different acid and one of its salts.

Alkaline buffer solutions have a pH greater than 7 and are made from a weak base and one of its salts. A very commonly used example of an alkaline buffer solution is a mixture of ammonia and ammonium chloride solution. If these were mixed in equal molar proportions, the solution would have a pH of 9. Similarly when NaOH strong base is added to this buffer system, the ammonium ion donates a proton to the base to become ammonia and water thus neutralizing the base without any significant pH change.