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TUMS
Azin Nowrouzi, PhD
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Functions of H+
• [H+] and therefore pH is important for many
biochemical processes:
– Transport of oxygen in blood
– Enzymatic reactions
– Generation of metabolic energy during respiration
and photosynthesis
– Metabolism and its regulation.
– Biomolecules amino acids , proteins, sugars,
lipids, nucleic acids, vitamins, and …. are
sensitive to pH variations.
– Measurement of pH of blood and urine are commonly
used in medical diagnoses.
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Acids and Bases
• Acid: a substance that can donate H+.
• Base: a substance that can accept H+.
• Water (H2O): Both an acid and a base.
One H2O molecule, acting as an acid, may
donate a proton to another H2O molecule,
acting as a base.
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Self-ionization (autoprotolysis) of water
Keq = 1.8 x 10-16
[H2O]=55.5 M
because (1000g/L)/(18.015
g/mol)
(55.6)Keq = [H+][OH-]
(55.6)(1.8 x 10-16) = [H+][OH-] = Kw
Kw = 1.0 x 10-14 = [H+][OH-]= [H+]2
[H+] = [OH-] = = 1.0 x 10-7 M
pH + pOH = 14
If we put an acid into a solution to decrease the pH the pOH must
increase to keep their sum equal to pKw.
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Definition of buffer
• Buffer action = the ability to oppose changes in
pH when small amounts of acid or base are
added to the solution.
• Buffers are mixtures of weak acids and their
conjugate bases.
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Buffers
• The mathematical basis of buffer action is given by
the Henderson-Hasselbalch equation:
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• Liquid inside cells (Intracellular fluid) - 40 %
• Extracellular fluids - 15 %
• Liquid making the plasma of blood - 5 %
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Fluid spaces in the human body
The ICF and ECF are called
fluid compartments because
they behave as distinct
entities.
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Intracellular fluid
35%-40% BW
Total body
water
Blood plasma
(intravascular fluid)
50%-60% BW
4%-5% BW
Extracellular fluid
15%-20% BW
Interstitial fluid
BW = Body Weight
Lymph
(extravascular)
11%-15%BW
Transcellular fluid:
Cerebrospinal fluid
Intraocular fluid
Synovial fluid
Pericardial fluid
Pleural fluid
Peritoneal fluid 9
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Sources of acid and base in the body
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Types and sources of Acids in the Body
•
•
•
•
Acid-forming foods are flesh foods,
grains, dairy products, a majority of
nuts and seeds, beans and peas,
simple sugars, fats, and proteins.
•
Fats as a general class are acidforming because their metabolism
produces acetic acid. Simple
carbohydrates like white sugar are acid- •
forming because they enter the system
too quickly and burn too fast.
Alkaline-forming foods are
vegetables and fruits. Exceptions
include asparagus, cranberries, plums,
and prunes, which are all slightly acidic.
While fruits such as lemons, oranges,
grapefruits, tomatoes and pineapples,
are acid when you eat them, but, by the
process of enzymatic digestion; they
are turned into alkaline substances.
Volatile (CO2, Carbonic acid)
– Can leave the solution and
enter the atmosphere.
– End product of oxidation of
glucose and triglyceride during
aerobic metabolism.
Fixed (non-carbonic)
•
Acids that do not leave
solution.
– Inorganic acids (Sulfuric Acid,
Phosphoric Acid)
•
Catabolism of amino acids,
nucleic acids, and
phospholipids.
– Organic (Lactic acid, Ketone
Bodies)
•
By products of aerobic
metabolism, anaerobic
metabolism, starvation and
diabetes.
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Production of
Dietary Source
Quantity Produced
(mEq/day)
Acid
Alkali
Acid
Alkali
Carbohydrate
0
0
0
0
Fat
0
0
0
0
Amino acids
Sulfur-containing (Cys, Met)
Cationic (Lys, Arg, His)
H2SO4
70
HCl
135
100
HCO3-
Anionic (Asp, Glu)
Organic anions
(Acetate, Citrate, Gluconate, Lactate, Malate)
Phosphate
Net acid production
60
HCO3H2PO4-
30
75
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• The role of
intracellular
and
extracellular
buffers,
respiratory,
and renal
mechanisms
in maintaining
normal blood
pH.
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Different Lines of Defense against
Acid –Base disorders
• First Line of defense:
Buffer System
(buffering)
The body has a HUGE buffering capacity, and
This system is essentially IMMEDIATE in effect.
• Second Line of defense: Respiratory System
Alteration; in arterial pCO2
(compensation)
minutes to hours
• Third Line of defense:
Renal System
Alteration in HCO3- excretion
(correction)
several days
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Acid-Base Balance
I.
Buffering systems of the body
1.
The bicarbonate/carbonic acid system
•
2.
3.
The phosphate/phosphoric acid buffer system
Protein buffering
•
•
•
4.
II.
III.
The role of the erythrocyte in the bicarbonate system
The hemoglobin (Hb) system
Cellular buffering
Bone buffering
Ammonia buffering system
The lungs
The kidneys
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Important buffers at physiologic [H+]
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1. Bicarbonate/carbonic acid buffer system
CO2 as an acid
1.
The greatest source of H+.
2.
End product of oxidation of glucose and triglyceride during aerobic metabolism.
3.
Most important weak acid in body fluids.
[H+] x [HCO3-] = 24 x PCO2
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Importance of bicarbonate buffer system
• Most important
physiological buffer
because:
 It is an open system: the
acid component of the
system is volatile (CO2),
and can be excreted
rapidly by pulmonary
respiration.
 It transports H+ to the
kidney and transfer it to
other buffers like
phosphates to be
excreted in urine.
 It’s in high concentration
in plasma (HCO3- = 24
meq/l).
 It’s the bicarbonate buffer
system that is regulated.
H2CO3 = PCO2 X 0.03
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The role of the erythrocyte in the
bicarbonate system
It is an open system: the acid component of the system is
volatile (CO2), and can be excreted rapidly by pulmonary
respiration.
Carbonic anhydrase (CA): Located in
• The cytosol of cells (renal cells and RBC’s).
• Along the brush border of the proximal tubule. Membrane-bound
enzyme with the active site facing the nephron lumen.
• Not found in plasma.
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Hydrogen ion (H+) secretion and bicarbonate ion
(HCO3-) reabsorption in a renal tubular cell
It transports H+ to the kidney and transfer it to other
buffers like phosphates to be excreted in urine.
• Carbon dioxide (CO2)
diffuses from the blood or
urine filtrate into the tubular
cell, where it combines with
water in a carbonic
anhydrase-catalyzed
reaction that yields carbonic
acid (H2CO3).
• The H2CO3 dissociates to
form H+ and HCO3-.
• The H+ is secreted into the
tubular fluid in exchange for
Na+.
• The Na+ and HCO3- enter
the extracellular fluid. "
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Acid-base homeostasis
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2. Phosphate/phosphoric acid buffer sytem
• Phosphate concentration in the blood is so low that it is
quantitatively unimportant.
• Phosphates are important buffers intracellularly and in urine where
their concentration is higher.
• Phosphoric acid is triprotic acid that exists in four forms.
Phosphoric acid has a pKa value for each of the three dissociations
with three pK values of 2.1, 6.8, and 12.7:
H3PO4 ↔ H2PO4- ↔ HPO4-2 ↔ PO4-3
• When a solution of phosphoric acid (H3PO4) is titrated completely,
there will exist three buffering regions.
• Physiologically, we only need to deal with the pH range between pH
4.5 (acidic urine) and 8.0 (alkaline urine).
• For this reason, only two forms of phosphate (H2PO4- and HPO4-2)
and one pK value (6.8).
• H2PO4- is the acid form, and HPO4-2 is the salt form.
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The renal phosphate buffer system
• The monohydrogen
phosphate ion (HPO42-)
enters the renal tubular
fluid in the glomerulus. A
H+ combines with the
HPO42- to form H2PO4and is then excreted into
the urine in combination
with Na+. The HCO3moves into the
extracellular fluid along
with the Na+ that was
exchanged during
secretion of the H+.
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3. Hemoglobin as a buffer system
•
•
•
•
•
Located in RBC.
Very important buffer of plasma pH.
pK of hemoglobin is dependent on oxygenation.
The drop in pH due to CO2 addition to the blood is less than predicted
(from 7.40 to 7.37 and not 7.32) because of the action of hemoglobin.
This is because in the RBC, hemoglobin acts as a buffer for H+.
•
•
•
Hemoglobin protein can reversibly
bind either H+ (to the protein) or O2
(to the Fe of the heme group).
But that when one of these
substances is bound, the other is
released (as explained by the Bohr
effect).
During exercise, hemoglobin helps
to control the pH of the blood by
binding some of the excess
protons that are generated in the
muscles. At the same time,
molecular oxygen is released for
use by the muscles.
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• Mechanisms of carbon dioxide transport. CA, Carbonic Anhydrase.
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H+ attachment to hemoglobin
T state
R state
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Cellular (protein) buffer system
• Proteins can act as buffers.
• Because each protein molecule is both a weak acid and a weak base.
• The weakly acidic carboxylic acids counteract rising pHs while the
weakly basic amino groups can counteract falling pHs.
• Plasma proteins and hemoglobin (blood's oxygen-carrying pigment, a
protein) enhance the blood's buffering capacity.
• Human albumin contains 585 amino acid residues, which includes 61
glutamic acid, 24 arginine, 26 aspartic acid, and 16 histidine amino
acids. Dependent on the pH, albumin can act as a conjugate base
(accepts protons) or a conjugate acid (donates protons).
(1)
(2)
(3)
(4)
(5)
If pH = 3.0, its net charge is +60 (it has 60 protons to donate).
If pH = 4.9, it is a zwitterion, electrically neutral.
If pH = 7.5, its net charge is -10 (it can accept 10 protons)
If pH = 8.6, its net charge is -20
If pH = 11.0, its net charge is -60
• Proteins also help with buffering within cells.
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The ammonia buffer systems
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4. Bone buffer systems
• When protons are added to the body
– The immediately available buffers are those of the blood and
extracellular fluid
– But, after some time, intracellular buffers and bone play their
role.
• The uptake of protons by bone can occur in exchange
for surface Na+ and K+ and by dissolution of bone
mineral (bone carbonate).
– Bone contains 80% of total CO2 (CO32-, HCO3- and CO2) in the
body.
– Two-thirds in the form of CO32- complexed with Ca2+, Na+, and
other cations in bone crystals.
– One-third in the form of HCO3- in the hydroxyapatite crystals.
• Chronic metabolic acidosis is dangerous because of
progressive bone lysis and in children, growth
retardation.
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Acid-basic disorder evaluation
Constant
[HCO3- ]
[CO2]
pH= pK +log
24
＝ 6.1+log
1.2
[HCO3- ]
[CO2]
Changes lead to acidosis or Alkalosis
Metabolic
Respiratory
＝6.1+1.3= 7.4
> 1.3
Alkalosis
kidneys
Lungs
or
Normal
< 1.3
Acidosis
[HCO3- ]
[CO2] 35
Respiratory Acidosis: [CO2] ↑
• A process that drives blood pH below 7.35.
• Occurs when CO2 is not eliminated from the body at
normal rate.
• Usual cause is decreased respiration or hypoventilation.
• When respiration is restricted, the concentration of
dissolved carbon dioxide in the blood increases, making
the blood too acidic. Such a condition can be produced by
asthma, pneumonia, emphysema, or inhaling smoke.
• An elevated plasma PCO2 is called hypercapnia.
• Respiratory system is unable to eliminate all of the CO2
generated by peripheral tissues.
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Respiratory Alkalosis: [CO2] ↓
• A process that drives blood pH above 7.45
• Results from excessive breathing or hyperventilation.
• Hyperventilation causes too much dissolved CO2 to be
removed from the blood, which decreases the carbonic
acid concentration, which raises the blood pH.
• Often, the body of a hyperventilating person will react by
fainting, which slows the breathing.
• Problems with this are relatively uncommon.
• When plasma PCO2 is below normal levels, the condition
is called hypocapnia.
• Temporary hypocapnia can be produced by
hyperventilation.
• Condition seldom persists long enough to cause a
clinical emergency.
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Metabolic Acidosis: [HCO3- ]↓
• Decrease in blood pH that results when excessive
amounts of acidic substances (metabolic acids such as
lactic acid or ketone bodies) are released into the blood.
• This can happen through
– prolonged physical exertion,
– diabetes
– restricted food intake (starvation).
• Occurs when bicarbonate ion concentrations decrease.
The acids are buffered by HCO3- lowering its levels.
• The normal body response to this condition is to
increase breathing to reduce the amount of dissolved
carbon dioxide in the blood. This is why we breathe more
heavily after climbing several flights of stairs.
• Can be caused by the inability to excrete hydrogen ions
from the kidneys and any condition causing severe
kidney damage.
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Metabolic Alkalosis: [HCO3- ]↑
• Occurs when bicarbonate ion concentrations
become elevated.
• This can result from the ingestion of alkaline
materials, and through overuse of diuretics.
• Bicarbonate ions (from alkaline foods) interact
with hydrogen ions in solution to form carbonic
acid and the reduction of H+ ions causes
symptoms of alkalosis.
• Again, the body usually responds to this
condition by slowing breathing, possibly through
fainting.
• Cases of severe metabolic alkalosis are
relatively rare.
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The four primary acid-base disturbances
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Normal and compensated states of pH
and acid-base balance represented as a
balance scale
A.
B.
C.
D.
E.
•
When the ratio of bicarbonate (HCO3-) to carbonic acid (H2CO3,
arterial PCO2  .03) = 20:1, the pH = 7.4.
Metabolic acidosis with an HCO3-;H2CO3 ratio of 10;1 and a pH of
7.1.
Respiratory compensation lowers the H2CO3 to 0.6 mEq/L and
returns the HCO3-:H2CO3 ratio to 20;1 and the pH to 7.4.
Respiratory alkalosis with an HCO3-;H2CO3 ratio of 40;1 and a pH of
7.7.
Renal compensation eliminates HCO3-, reducing serum levels to 12
mEq/L, returning the HCO3-;H2CO3 ratio to 20;1 and the pH to 7.4.
Normally, these compensatory mechanisms are capable of buffering
large changes in pH, but may not return the pH completely to normal,
as illustrated here.
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Determining acid/base disorder
Practice using the following gases:
pH=7.3
PCO2= 60
HCO3-= 24
"______?________ , ______?________ , _______?______"
Step one: Determine whether pH is acidic or
basic
Step two: Evaluate PCO2
Step three: Evaluate HCO3- or base excess
Step 4: Put all the information together and
look for compensation.
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Step One:
Determine whether pH is acidic or basic
• pH must always be determined first when assessing
acid/base balance.
– If pH is > 7.40, blood is more basic. If pH is < 7.40, blood is acidic.
– If pH is between 7.35 and 7.45, the acid base disturbance is
compensated.
• Determine which the pH is: an acid or base; then mark an
‘A’ for acid, ‘B’ for base or ‘N’ for normal beside the pH.
You have now determined which imbalance, (acidosis or
alkalosis) is present.
• From the above example:
• ph = 7.3 " A“, PCO2=60 ________, HCO3=24 _________
"______?________, ______?_______ , ___Acidosis___"
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Step Two:
Evaluate PCO2
• PCO2 is the respiratory component of acid/base
balance.
**Remember that PCO2 is an acid.
• A PCO2 > 40 favors resp. acidosis
• A PCO2 < 40 favors resp. alkalosis
•
Normal range for PCO2 is 35 - 45.
• Mark and ‘A’ , ‘B’ or ‘N’ beside the PCO2 value.
• In the example, PCO2 = 60 so we put an "A" beside
the PCO2.
"______?______, Respiratory Acidosis"
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Step three:
Evaluate HCO3- or base excess (BE)
• HCO3- < 22 favors metabolic acidosis
HCO3- > 26 favors metabolic alkalosis
• A positive base excess (BE) favors
metabolic alkalosis; a negative base
excess favors acidosis.
• Mark a ‘A’ , ‘B’ or ‘N’ beside the value. In
our example HCO3 = 24 so we put an "N"
beside the value.
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Step 4:
Put all the information together and look
for compensation
• Circle all the letters. Match the ones that are
the same.
• In our example, the pH is acidic, the PCO2 is
acidic and the HCO3 is normal. Therefore the
disorder is a respiratory acidosis because
the PCO2 and the pH match.
• So, we can fill in 2 blanks:
"___________ respiratory acidosis."
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Determine the compensation
1. Uncompensated-The gases are
uncompensated if the pH is outside the
normal range ie <7.35 or > 7.45
– In the above example the gases are
uncompensated because the pH is 7.3
2. Partial compensation: pH is within the
normal range of 7.35 - 7.45 but is not
7.40
3. Complete compensation: pH is 7.4
“Uncompensated Respiratory Acidosis”
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Hypoxemia
• Normal PO2 is 80 to 100 mmHg (for a healthy
person, less than 60 years old, breathing
room air).
• Hypoxemia- Definition: deficient oxygenation of
the blood.
**Measured by PO2
• PO2 < 80 : mild hypoxemia
PO2 < 60 : moderate hypoxemia
PO2 < 40 : severe hypoxemia
• age: For every year above 60, a patients’ normal
PO2 drops by 1 mm Hg
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Examples: Arterial blood gas interpretation
PO2=70
pH= 7.33 A B or N
PCO2= 48.3 A B or N
HCO3=25.8 A B or N
PO2=88
pH= 7.45 A B or N
PCO2= 43 A B or N
HCO3=30.1 A B or N
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Example: Evaluation of acid-base status
•
•
•
•
•
•
•
•
In a patient with chronic respiratory acidosis:
Arterial PCO2 = 60mmHg
Plasma [HCO3-] = 32 mmoles/L
Arterial pH = 7.35
BE = [HCO3-]observed – [HCO3-]predicted
Observed [HCO3-] is calculated from pH and pCO2
Predicted [HCO3-] = 24 + (2.733 + 0.52Hb)(7.4-pH)
Predicted [HCO3-] = 24 + 10(7.4-pH) if Hb is normal
and equal to 15g/100ml of blood.
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Davenport diagram
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Modified Henderson Equation
This is the equation used to derive [HCO3-] from pH and PCO2.
[H+] x [HCO3-] = 24 x PCO2
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The acid-base nomogram
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Protonation and deprotonation of histidine residues in Hb and
other proteins.
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