1. MLAB 2401: Clinical Chemistry Keri 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 Alkalosis Note: Normal pH is 7.35-7.45
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

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 + H 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 )

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 mEq/L
1.2 mEq/L
pH = pK + log 20
pH= pK+ 1.30
pH=
pH= 7.40

22. The Ratio….
Normal is : = Bicarbonate = Kidney = metabolic
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: 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 .