Acid-Base Balance: Overview

MLAB 2401: Clinical ChemistryKeri Brophy-Martinez

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+

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

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

Terms

Acidosis pH less than 7.35

Alkalosis pH greater than 7.45

Note: Normal pH is 7.35-7.45

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

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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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.

Blood Buffer Systems

Four Major Buffer Systems Protein Buffer systems

Amino acids Hemoglobin Buffer system

Phosphate Buffer system Bicarbonate-carbonic acid Buffer system

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

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

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

Bohr Effect

It all about oxygen affinity!

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 –

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• Essential within the erythrocytes

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

Bicarbonate/Carbonic Acid Buffer

Carbonic acid

Bicarbonate

Conjugate base

Excreted in urine

Excreted by lungs

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

Henderson-Hasselbalch Equation

Relationship between pH and the bicarbonate-carbonic acid buffer system in plasma

Allows us to calculate pH

Henderson-Hasselbalch Equation

General Equation

pH = pK + log A-

HA

Bicarbonate/Carbonic Acid system

o pH= pK + log HCO3

H2CO3 ( PCO2 x 0.0301)

Henderson-Hasselbalch Equation

1. pH= pK+ log H

HA

2. The pCO2 and the HCO3 are read or derived from the blood gas analyzer

pCO2= 40 mmHg

HCO3-= 24 mEq/L

3. Convert the pCO2 to make the units the same

pCO2= 40 mmHg * 0.03= 1.2 mEq/L

3. Lets determine the pH:

4. Plug in pK of 6.1

5. 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

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

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

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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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 .

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