Acid - Base Balance

As stated in Egan's, the body's metabolic processes continually generate hydrogen ions . . . even small hydrogen ion concentration changes can cause metabolic processes to fail. Acid-base balance relies on the mechanisms that keep hydrogen ion concentration of body fluids compatible with life by maintaining the pH between 7.35 and 7.45.

Three major organ systems devote much of their activity to regulating acid and alkali levels: the lungs, the blood, and the kidneys. Buffering of acids is largely carried out by the blood and kidneys; excretion of acids is handled by the lungs and kidneys. Of the 12,000mEq pf H+ produced per day, 99% of this is excreted by the lungs. (Respiratory Anatomy and Physiology, Martin & Youtsey)

Everything we do as Respiratory Care Practitioners has this goal - even oxygen administration, by preventing anaerobic metabolism and the generation of lactic acid, helps to maintain proper pH.

 

Hydrogen Ion Regulation in Body Fluids

Hydrogen ions come from volatile or fixed acids. Most H+ is the result of the breakdown of proteins, anaerobic metabolism, or the metabolism of body fats. These are called fixed acids. H+ that are the result of CO2 hydrolysis (H2CO3) is called a volatile acid as it is an acid in equilibrium with a dissolved gas.

Fixed acids are not in equilibrium with a dissolved gas and come from the catabolism of proteins:

Sulfuric and phosphoric acids

Lactic acid (from anaerobic glycolysis)

Ketoacidosis (from fatty acid metabolism)

H+ of fixed acids combines with HCO3- to form H2CO3 and causes an equilibrium imbalance that drives the reaction to the left to form CO2 and H2O

Two major mechanisms that stabilize pH in response to CO2 generated by metabolism are:

Isohydric buffering: H+ is carried by deoxygenated hemoglobin and is released as CO2 in the lungs

Ventilation: elimination of carbonic acid = production

Strong Acids

Acidity of a solution reflects only the free hydrogen ions, not the hydrogen ions still combined with anions. 

Thus, strong acids, which dissociate completely (i.e., they liberate all the H+) and irreversibly in water, dramatically change the pH of the solution.

For example, if 100 hydrochloric (HCl) acid molecules were placed in 1 mL of water, the hydrochloric acid would dissociate into 100 H+ and 100 Cl ions. There would be no undissociated hydrochloric acid molecules in the solution.

Weak Acids

Do not dissociate completely in a solution and have a much smaller effect on pH.

Although weak acids have a relatively small effect on changing pH levels, they have a very important role in resisting sudden pH changes.

Examples of weak acids are carbonic acid (H2CO3) and acetic acid (HC2H3O2

If 100 acetic acid molecules were placed in 1 mL of water, the following reaction would occur:

weak acid.jpg

Because undissociated acids do not alter the pH, the acidic solution will not be as acidic as the HCl solution discussed above

The dissociation of acetic acid can be written as follows:

acetic acid.jpg

Using this equation, it can be seen that:

When H+ (released by a strong acid) is added to the acetic acid solution, equilibrium moves to the left as some of the additional H+ bonds with C2H3O2 to form HC2H3O2

On the other hand, when a strong base is added to the solution (adding additional OH and causing the pH to increase), the equilibrium shifts to the right. This occurs because the additional OH consumes the H+, causing more HC2H3O2 molecules to dissociate and replenish the H+ .

Weak acids play a very important role in the chemical buffer systems of the human body

Strong Bases

 Remember, bases are proton acceptors 

Strong bases dissociate easily in water and quickly tie up H+

Hydroxides are strong bases

Weak Bases

In contrast, weak bases:

Sodium bicarbonate or baking soda is a weak base

Dissociate incompletely and reversibly and are slower to accept protons 

Because sodium bicarbonate accepts a relatively small amount of protons

Its released bicarbonate ion is described as a weak base