Recall the equation for pH. Here is the calculation that should be performed:

Our main blood buffer system works to protect against large pH changes in the blood. This system relates bicarbonate, carbonic acid, and carbon dioxide via the following equilibria:

While the phosphate buffer system is an important biological buffer, it is not the main buffer system in human blood.

Here are the equations we need to use to find the pOH of our solution of sulfuric acid:

In this question, we're told that a buffer solution consisting of carbonic acid and bicarbonate needs to be prepared. The solution needs to have a concentration of , a volume of , and a pH of .

First, we need to recognize that in order to solve this problem, we'll need to utilize the Henderson-Hasselbalch equation.

The acid in the above expression will be carbonic acid, and its conjugate base is bicarbonate. We can plug in the values we have for pH and pKa to obtain the ratio of base to acid.

Now that we have the ratio of base to acid, we can figure out what fraction of our solution will be base, and what fraction of it will be acid. To find this fraction, we have to realize that for every  moles of base, there is   mole of acid. Thus, there are a total of  moles of acid and base.

Next, we need to take into account the volume and molarity that we want in our desired solution in order to find the total number of moles for our final solution.

Equipped with knowing how many total moles we want in our desired solution, in addition to the portion of our solution that is acid and the portion that is base, we can at last calculate how many moles each of carbonic acid and bicarbonate we need in our solution.

Carbonic acid:

Bicarbonate:

Carbon dioxide concentration is low in the blood surrounding the lungs and high in those around muscle, relatively speaking, because cell respiration by the muscles produces carbon dioxide. Therefore, hemoglobin binds oxygen quite well at the lungs, as it should in order for us to be able to breathe, and then binds it much less effectively around the muscles. This allows the oxygen to unbind and enter the muscle tissue (this process is facilitated by myoglobin) where it is used by the actively respiring muscle cells. 

Carbonic anhydrase in red blood cells catalyzes the following equilibrium reaction:



The carbonic acid formed then naturally gives up a proton to form . This is why elevated  levels lower the blood pH. When more  is removed through more rapid breathing, this shifts the equilibrium reaction back to the left (Le Chatelier's principle), prompting more  and  to turn into  which then turns into  and . Thus, carbonic anhydrase now catalyzes the reaction toward the left in this scenario, raising the pH.  is neither an acidic nor basic molecule by itself (and if it were basic then ridding of it would actually lower the pH further), and while the answer involving oxygen may be tempting, oxygen is not found dissolved in blood plasma; it is bound to hemoglobin, and thus it does not bind protons in the blood to form water.  

Using the Henderson Hasselbach equation:

Thus, the ratio of acid to base = 

This question requires you to use the Henderson-Hasselbach equation, one of the most important equations in biochemistry. The equation is:

where is the concentration of the conjugate base, and  is the concentration of the acid. In this scenario,  is the conjugate base, while  is the acid. With the numbers given in this question, the equation should look like this:

Here is the equation used to find the correct answer:

Water, , is an amphoteric substance--it can act as either an acid or a base. In certain circumstances, water can act as a Bronsted-Lowry acid by donating a proton.

As seen above,  donated one of its hydrogen atoms, becoming .

In other cases, water can act as a Bronsted-Lowry base by accepting a proton. 

As seen above,  accepted a hydrogen atom to become .

Therefore, water can act as either an acid or a base depending on the situation. There are other amphoteric substances, but water is definitely the most common.