Biochemical Buffers and pH Regulation: Pharmacological Implications

Introduction to pH Regulation

Changes to the pH of a solution can drastically change its properties. The basic concepts of pH are crucial to apply in regard to the highly regulated pH balances of the blood. The body uses organ systems to maintain pH in the body at a slightly basic level, 7.35-7.45. A pH imbalance in the body can lead to adverse health effects and can be characterized as an acidosis (pH < 7.35) or alkalosis (pH > 7.45). Additionally, pH imbalances can be further classified based on their altered system; a metabolic pH imbalance primarily relates to the kidneys but also the digestive system whereas a respiratory imbalance relates to the lungs, respiratory rate, and CO2 levels. Compensation can occur in the body to aid in the imbalance which will be carried out by the opposite system. It is important for pharmacists to understand pH imbalances since drugs can contribute to an imbalance and imbalances can impact drug metabolism which is crucial to interpret in terms of patient safety with medications.

Terminal objective: To explain the function of buffers in physiological conditions and to interpret patient blood gas data related to blood pH.

Enabling Objectives for pH Understanding

1. To describe the basic concepts of pH, acids, and bases.
2. To explain the purpose of biochemical buffers and function of pKa in reference to the system pH.
3. To describe the bicarbonate buffer system at the normal blood pH.
4. To explain how the body maintains normal pH related to the respiratory and renal systems.
5. To interpret the four types of pH imbalances and compensation status based on arterial blood gas data and their possible causes.

Recommended reading: Lippincott's Illustrated Reviews: Biochemistry, 8th Edition, Chapter 1 "Amino Acids and the Role of pH" pages 6-101. Introduction to pH

pH Definition and Scale

a. pH = - log ([H+]) where H+ represents hydrogen ions and pH ranges from 1(strong acid) - 14 (strong base); concentration is represented by brackets [ ].

i. pH = 7 is neutral, which signifies an equal concentration of protons (H+) and hydroxide (OH-) ions 1. [H+] = [OH ] = 1 x 10-7 mol/L -> pH = - log(1 x 10-7) = 7

ii. pH < 7 is acidic

iii. pH > 7 is basic

Strong Acids and Bases

b. Strong acids and bases will fully ionize in aqueous solutions which include blood. For acids, that means ionization to form H+ ions and for a base, OH- ions.

i. Acid - H+ will be present when the acid is dissolved in water 1. Ex. Hydrochloric acid (HCI) dissociates to H+ and CL

ii. Base - OH- ions present when dissolved in water 1. Ex. Sodium hydroxide (NaOH) in water dissociates to Na+ and OH-

Weak Acids and Bases

c. Weak acids and bases will only partially ionize in aqueous solution and the degree of this ionization is based on the dissociation constant, Ka. The dissociation constant represents the strength of the acid. Dissociation will form conjugate acid-base pairs.

i. Ka = [HA]

Henderson-Hasselbalch Equation

d. Henderson-Hasselbalch equation: pH = pKa + log [HA]

i. pKa = - logKa

ii. A- is the conjugate base (or salt); it is the ionized form of the acid, HA.

iii. The Henderson-Hasselbach equation is a conversion of the Ka equation. By solving for the concentration of H+ and then taking the logarithm of both sides, multiplying by -1, and substituting log equivalents of pH and pKa.

iv. Henderson-Hasselbalch equation gives an idea of the relationship between pH and concentrations of the weak acid and conjugate base.

v. The higher the Ka, the lower the pKa, the stronger the acid - more dissociation into H+ and A-

vi. The lower the Ka, the higher the pKa, the weaker the acid - less dissociation into H+ and A-

vii. Physiologically relevant weak acids present in human blood that will ionize to the conjugate base plus hydrogen ions: 1. Weak acid = Lactic acid; conjugate base = lactate 2. Weak acid = Acetic acid; conjugate base = acetate 3. Weak acid = Aspartic acid; conjugate base = aspartate 4. Weak acid = Ammonium (NH4+); conjugate base = ammonia (NH3)

viii. Law of mass action - at a specific temperature, the reaction rate is proportional to the product of the concentrations of the reactants

ix. Chemical equilibrium - concentrations of the reactants and products stay constant over time

Biochemical Buffers

Buffer definition (Lippincott's) - "solution that resists a change in pH following the addition of an acid or base and can be created by mixing a weak acid (HA) with its conjugate base (A-)"

Buffering Capacity

a. Maximum buffering capacity occurs when the weak acid concentration [HA] is equal to the conjugate base concentration [A ], meaning pH = pKa.

i. When those concentrations are equal, the equation would be pH = pKa + log(1) and log of 1 is 0 hence why pH would equal pKa.

ii. pKa is determined by the titration curve of the buffer solution.

b. Acid-base conjugate base pairs can buffer efficiently (not maximally) at +1 pH unit.

c. If acid is added to the buffer solution, the conjugate base can neutralize it and produce HA.

d. Similarly, if a base is added to a buffer solution, the weak acid can neutralize it and convert to the conjugate base.

e. The diagram to the right provides an example for acetic acid. The pKa = 4.8, so pH = 4.8 would be the mass buffering capacity but there will still be a resistance to change in pH in the buffer region of pH = 3.8-5.8.

OH" H2O CH3COOH FORM I (acetic acid, HA) CH3COO FORM II H+ (acetate, A“)

Buffer region 1.0 Equivalents OH" added

[]] = []]] pKa = 4.8 0.5- [U] > [] 0 - T T 0 3 4 5 6 7 pH

Buffering System in the Blood

a. Blood pH is slightly basic at 7.4 with a range of 7.35-7.45 for arterial pH, and it is crucial that the body maintains this pH to avoid adverse health effects with the help of the bicarbonate buffer system.

i. It is necessary to be at this pH range for enzymes and proteins to function optimally in the body.

ii. Some enzyme exceptions include various digestive enzymes that function optimally at the acidic pH of the stomach (pH 1.5-3.5) and lysosomal enzymes at the acidic pH of 4.5-5.0.

Acid Production and Susceptibility

b. Metabolism of carbohydrates, fats, and proteins can lead to production of acids so humans are more susceptible to acidosis than alkalosis. Examples:

i. Lactic acid produced during metabolism.

ii. CO2 generated from glucose and fatty acid oxidation.

iii. Amino and fatty acids themselves are acids.

iv. Citric acid cycle intermediates

v. Ketoacids produced in high amounts in cases of Type I Diabetes Mellitus

vi. Some dietary sources like citrus are acidic but unlikely to eat in high enough amounts to have a substantial impact on blood pH.

Factors Influencing Blood pH

c. pH of the blood is influenced by the concentration of bicarbonate (HCO3 ) and carbon dioxide (CO2) with major impact by the kidneys and lungs.

d. There is an alkaline, bicarbonate (HCO3 ), reserve to aid with the higher production of acids vs bases in the body.

e. CO2 is a volatile, weak acid that is produced from oxidative metabolism.

i. Bicarbonate buffer system: CO2(dissolved) + H20 >> H2CO3 <> HCO3 + H+

ii. CO2 that is dissolved in water in the plasma (CO2 is relatively water-insoluble) will form its aqueous form, carbonic acid (H2CO3), which can further dissociate to water-soluble HCO3 and H+.

iii. CO2 is the anhydrous form (without water) of H2CO3; dissolved CO2 will diffuse into red blood cells and be converted to H2CO3 by the enzyme carbonic anhydrase.

iv. pKa of this bicarbonate buffer system is 6.1

v. Henderson-Hasselbalch equation with CO2 as the weak, conjugate acid [HA] and HCO3 as the conjugate base: pH = pKa + log [CO2] [HCO3] so for typical blood values we arrive at pH of 7.4 with pka = 6.1, [HCO3 ] = 24.0 mM, [CO2] = 1.2 mM -> pH = 6.1 + log(24/1.2) = 7.4.

1. There is more bicarbonate present at pH of 7.4 to help buffer the acids produced from metabolism and other sources.

2. Looking at the equation it is clear that an increase in bicarbonate will increase the pH, making it more basic, and an increase in carbon dioxide will decrease the pH, making it more acidic.

3. An increase in H+ concentration due to increase of organic and metabolic acids in the system will shift the equation to the left to restore equilibrium (law of mass action). Some of the H+s will associate with HCO3 to form carbonic acid and ultimately lead to CO2 produced and expelled in breath leading to a lesser reduction of pH than if there was no bicarbonate buffer system.

pH Imbalances

The main systems involved in maintaining normal blood pH is the respiratory and renal systems involving the lungs and kidneys, respectively. They help to provide acid-base regulation.

Types of Imbalances

a. Four possible imbalances

i. Respiratory acidosis - caused by hypoventilation, decreased respiration rate

ii. Respiratory alkalosis - caused by hyperventilation, increased respiration rate

iii. Metabolic acidosis - loss of bicarbonate or retention of acids

iv. Metabolic acidosis - retention of bicarbonate or loss of acids

Respiratory System Role in pH

b. Respiratory system - takes in O2 in and expels CO2, a weak acid.

i. Lungs can regulate the rate of respiration which determines how much CO2 is excreted from the body and thus controls the blood pH.

ii. Hypoventilation will lead to respiratory acidosis (pH < 7.35) because of the higher retention of CO2 and therefore of an acid, leading to a decrease of pH.

1. CO2 in the blood is given in terms of pCO2, with the p standing for the partial pressure, the pressure exerted by CO2 dissolved in the blood.

2. A pCO2 >45 mmHg is typical of hypoventilation.

iii. Hyperventilation will lead to respiratory alkalosis (pH>7.45) because of the higher expulsion of CO2, leading to an increase of pH.

1. A pCO2 <35 mmHg is typical of hyperventilation.

iv. When delivering CO2 in red blood cells produced from metabolism, it is in the form of carbonic acid and is transported from the tissues to the lungs.

1. It is then converted to bicarbonate ion which is exported from the red blood cell through the chloride-bicarbonate exchange protein leading to the import of chloride ions and a chloride shift.

2. In the lungs, the bicarbonate ion re-enters the red blood cells and converts to CO2 which will dissociate and get exhaled from the lungs.

Renal System Role in pH

c. Renal system - retention and elimination control of acids and bases

i. The kidneys can regulate the elimination and retention of H+ and bicarbonate (a weak base) ions thus controlling the blood pH.

Clinical Acid-Base Imbalance: Respiratory Acidosis

d. Clinical acid-base imbalance: Respiratory acidosis

i. High pCO2 due to hypoventilation which prevents the elimination of CO2 by the lungs

1. Law of mass action of the bicarbonate buffer system: shift of the equilibrium will go to the right, carbonic acid production will increase and shift to produce more HCO3 and H+ ions. This resulting increase in H+ ions lowers the pH. While HCO3"