Transcript Acid-Base balance: overview

MLAB 2401: Clinical
Chemistry
Keri Brophy-Martinez
Acid-Base Balance: Overview
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


Alkalosis


pH less than 7.35
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 – 4
•
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
Conjugate
base
Bicarbonate
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.
2.
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
3.
Convert the pCO2 to make the units the same
pCO2= 40 mmHg * 0.03= 1.2 mEq/L
3.
Lets determine the pH:
Plug in pK of 6.1
4.
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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