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Acid-Base Balance: Overview MLAB 2401: Clinical Chemistry Keri Brophy-Martinez.

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Presentation on theme: "Acid-Base Balance: Overview MLAB 2401: Clinical Chemistry Keri Brophy-Martinez."— Presentation transcript:

1 Acid-Base Balance: Overview MLAB 2401: Clinical Chemistry Keri Brophy-Martinez

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

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

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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 CO 2 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 H + H 2 PO – 4 Essential within the erythrocytes

16 Blood Buffer Systems Bicarbonate/carbonic acid buffer system  Function almost instantaneously  Cells that are utilizing O 2, produce CO 2, which builds up. Thus, more CO 2 is found in the tissue cells than in nearby blood cells. This results in a pressure (pCO 2 ).  Diffusion occurs, the CO 2 leaves the tissue through the interstitial fluid into the capillary blood

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

18 Bicarbonate/carbonic acid buffer system How is CO 2 transported?  5-8% transported in dissolved form  A small amount of the CO 2 combines directly with the hemoglobin to form carbaminohemoglobin  92-95% of CO 2 will enter the RBC, and under the following reaction CO 2 + H 2 0 H + + HCO 3 -  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 o pH= pK + log HCO 3 H 2 CO 3 ( PCO 2 x )

21 Henderson-Hasselbalch Equation 1. pH= pK+ log H HA 2. The pCO 2 and the HCO 3 are read or derived from the blood gas analyzer pCO 2 = 40 mmHg HCO 3 - = 24 mEq/L 3. Convert the pCO 2 to make the units the same pCO 2 = 40 mmHg * 0.03= 1.2 mEq/L 3. Lets determine the pH: 4. Plug in pK of Put the data in the formula pH = pK + log 24 mEq/L 1.2 mEq/L pH = pK + log 20 pH= pK pH= pH= 7.40

22 The Ratio…. Normal is : 20 = Bicarbonate = Kidney = metabolic 1 carbonic acid Lungs respiratory The ratio of HCO 3 - (salt/bicarbonate) to H 2 CO 3 (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 CO 2  Doesn’t affect fixed acids like lactic acid  Body pH can be adjusted by changing rate and depth of breathing “blowing off”  Provide O 2 to cells and remove CO 2

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

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26 References Bishop, M., Fody, E., & Schoeff, l. (2010). Clinical Chemistry: Techniques, principles, Correlations. Baltimore: Wolters Kluwer 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. 26


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