MLAB 2401: Clinical Chemistry Keri Brophy-Martinez

1. MLAB 2401: Clinical ChemistryKeri Brophy-Martinez Acid-Base Balance: Overview

2. Terms • Acid • Any substance that can yield a hydrogen ion (H+) or hydronium ion when dissolved in water • Release of proton or H+ • Base • Substance that can yield hydroxyl ions (OH-) • Accept protons or H+

3. Terms • pK/ pKa • Negative log of the ionization constant of an acid • Strong acids would have a pK <3 • Strong base would have a pK >9 • pH • Negative log of the hydrogen ion concentration • pH= pK + log([base]/[acid]) • Represents the hydrogen concentration

4. Terms • Buffer • Combination of a weak acid and /or a weak base and its salt • What does it do? • Resists changes in pH • Effectiveness depends on • pK of buffering system • pH of environment in which it is placed

5. Terms • Acidosis • pH less than 7.35 • Alkalosis • pH greater than 7.45 • Note: Normal pH is 7.35-7.45

6. Acid-Base Balance • Function • Maintains pH homeostasis • Maintenance of H+ concentration • Potential Problems of Acid-Base balance • Increased H+ concentration yields decreased pH • Decreased H+ concentration yields increased pH

7. Regulation of pH • Direct relation of the production and retention of acids and bases • Systems • Respiratory Center and Lungs • Kidneys • Buffers • Found in all body fluids • Weak acids good buffers since they can tilt a reaction in the other direction • Strong acids are poor buffers because they make the system more acid

9. Blood Buffer Systems • Why do we need them? • If the acids produced in the body from the catabolism of food and other cellular processes are not removed or buffered, the body’s pH would drop • Significant drops in pH interferes with cell enzyme systems.

10. Blood Buffer Systems • Four Major Buffer Systems • Protein Buffer systems • Amino acids • Hemoglobin Buffer system • Phosphate Buffer system • Bicarbonate-carbonic acid Buffer system

11. Blood Buffer Systems • Protein Buffer System • Originates from amino acids • ALBUMIN- primary protein due to high concentration in plasma • Buffer both hydrogen ions and carbon dioxide

12. Blood Buffering Systems • Hemoglobin Buffer System • Roles • Binds CO2 • Binds and transports hydrogen and oxygen • Participates in the chloride shift • Maintains blood pH as hemoglobin changes from oxyhemoglobin to deoxyhemoglobin

13. Oxygen Dissociation Curve Curve B: Normal curve Curve A: Increased affinity for hgb, so oxygen keep close Curve C: Decreased affinity for hgb, so oxygen released to tissues

14. Bohr Effect • It all about oxygen affinity!

15. Blood Buffer Systems • Phosphate Buffer System • Has a major role in the elimination of H+ via the kidney • Assists in the exchange of sodium for hydrogen • It participates in the following reaction • HPO-24 + H+ H2PO – 4 • Essential within the erythrocytes

16. Blood Buffer Systems • Bicarbonate/carbonic acid buffer system • Function almost instantaneously • Cells that are utilizing O2, produce CO2, which builds up. Thus, more CO2 is found in the tissue cells than in nearby blood cells. This results in a pressure (pCO2). • Diffusion occurs, the CO2 leaves the tissue through the interstitial fluid into the capillary blood

17. Bicarbonate/Carbonic Acid Buffer Excreted by lungs Carbonic acid Conjugate base Bicarbonate Excreted in urine

18. Bicarbonate/carbonic acid buffer system • How is CO2 transported? • 5-8% transported in dissolved form • A small amount of the CO2 combines directly with the hemoglobin to form carbaminohemoglobin • 92-95% of CO2 will enter the RBC, and under the following reaction • CO2 + H20 H+ + HCO3- • Once bicarbonate formed, exchanged for chloride

19. Henderson-Hasselbalch Equation • Relationship between pH and the bicarbonate-carbonic acid buffer system in plasma • Allows us to calculate pH

20. Henderson-Hasselbalch Equation • General Equation • pH = pK + log A- HA • Bicarbonate/Carbonic Acid system • pH= pK + log HCO3 H2CO3 ( PCO2 x 0.0301)

21. Henderson-Hasselbalch Equation • pH= pK+ log H HA • The pCO2 and the HCO3 are read or derived from the blood gas analyzer pCO2= 40 mmHg HCO3-= 24 mEq/L • Convert the pCO2 to make the units the same pCO2= 40 mmHg * 0.03= 1.2 mEq/L • Lets determine the pH: • Plug in pK of 6.1 • Put the data in the formula pH = pK + log 24 mEq/L 1.2 mEq/L pH = pK + log 20 pH= pK+ 1.30 pH= 6.1+1.30 pH= 7.40

22. The Ratio…. Normal is : 20= Bicarbonate = Kidney = metabolic 1 carbonic acid Lungs respiratory • The ratio of HCO3- (salt/bicarbonate) to H2CO3 (acid/carbonic acid) is normally 20:1 • Allows blood pH of 7.40 • The pH falls (acidosis) as bicarbonate decreases in relation to carbonic acid • The pH rises (alkalosis) as bicarbonate increases in relation to carbonic acid

23. Physiologic Buffer Systems • Lungs/respiratory • Quickest way to respond, takes minutes to hours to correct pH by adjusting carbonic acid • Eliminate volatile respiratory acids such as CO2 • Doesn’t affect fixed acids like lactic acid • Body pH can be adjusted by changing rate and depth of breathing “blowing off” • Provide O2 to cells and remove CO2

24. Physiologic Buffer Systems • Kidney/Metabolic • Can eliminate large amounts of acid • Can excrete base as well • Can take several hours to days to correct pH • Most effective regulator of pH • If kidney fails, pH balance fails

26. References Bishop, M., Fody, E., & Schoeff, l. (2010). Clinical Chemistry: Techniques, principles, Correlations. Baltimore: WoltersKluwer Lippincott Williams & Wilkins. Carreiro-Lewandowski, E. (2008). Blood Gas Analysis and Interpretation. Denver, Colorado: Colorado Association for Continuing Medical Laboratory Education, Inc. Sunheimer, R., & Graves, L. (2010). Clinical Laboratory Chemistry. Upper Saddle River: Pearson .