Acid-Base Physiology
2
pH scale – to express hydrogen ion concentration.
The [H+] of ECF is very low (0.00004 mEq/L = 40 nmoles/L).
Normal variations are are markably small 3-5 mEq/L. It is
customary to express these very small numbers using the
logarithmic pH scale.
pH = - log10 [H+] or
pH = log 1 / [H+] log to the base 10 of the
reciprocal of hydrogen-ion concentration.
1) Because [H+] is in the denominator,
A high [H+] low pH and
A low [H+] high pH.
2) pH unit change of 1 = 10X change in [H+]
The Conceptual Problem with pH
• Because it’s a logarithmic scale, it doesn’t make
“sense” to our brains.
• EASY TO REMEMBER FACTS :– Every factor of 10 difference in [H+] represents 1.0 pH
units,
– Every factor of 2 difference in [H+] represents 0.3 pH
units.
• Therefore, even numerically small differences in
pH, can have profound biological effects…
4
ACIDS
• Acids are H+ donors.
• Acids can be:
– Strong – dissociate completely in solution
–HCl
– Weak – dissociate only partially in solution
–Lactic acid, carbonic acid
6
Volatile and Fixed Acids
• VOLATILE ACIDS :- carbonic acid
– Nearly 20,000 mEq of carbonic acid /day
• FIXED ACID :- lactate , keto acids, sulphuric
acid, phosphoric acid
-- Nearly 60-80mEq of fixed acids/day
7
BASES
• Bases are acceptors of H+(protons) or give up
OH- in solution
• Bases can be:– Strong – dissociate completely in solution
-NaOH
– Weak – dissociate only partially in solution
– NaHCO3
9
Weak acids thus are in equilibrium with their
ionized species:
Governed by the Law of Mass Action,
and characterized by an equilibrium
constant:
H+
HA
+
[H ][A ]
Ka = [HA],
+A
-
pKa = -log Ka
Derivation of the Henderson-Hasselbalch equation
• Ka = [H+] [A-]
[HA]
• so [H+] = Ka [HA]
[A-]
• TAKING THE NEGATIVE LOG OF BOTH SIDES
• As pH = - log [ H+],
• pH = -log Ka [HA]
[A-])
• pH = -log(Ka)-log([HA]
[A-])
• pH = pKa + log([A-]/[HA])
The Henderson Hasselbalch Equation
pH = pKa + log
[A ]
[HA]
L J HENDERSONK A HASSELBALCH
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Simplified form……
• pH = pKa + log ([A-]
[HA])
• pH = pKa + log(Conjugate base
Conjugate acid)
• pH = pKa + log(Proton acceptor
Proton donor )
The Body and pH
•
•
•
•
•
•
Homeostasis of pH is tightly controlled
Extracellular fluid = 7.4
Blood = 7.35 – 7.45
< 6.8 or > 8.0 death occurs
Acidosis (acidemia) below 7.35
Alkalosis (alkalemia) above 7.45
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Importance Of Maintenance Of pH Between 7.35 - 7.45(7.4)
Acidosis pH<7.35 and AlkalosispH>7.45.
Death occurs if pH falls outside the range of 6.8 to 8.0
• Altered [H+] results in changes in protein structure
(Enzymes, Receptors and ligands, Ion
channels,Transporters,Structural proteins)
• Function of excitable tissues
– Acidosis: hypoexcitability, CNS depression
– Alkalosis: hyperexcitability, tetany
• Affects K+ levels in the body.
Continuous addition of H+ ions to the body fluids and
3 Lines Of Defense Against pH Changes due to this:
• Buffering
• Changes in ventilation
• Changes in renal handling of H+ and HCO3-
Mechanisms of Regulation of pH
• FIRST LINE OF DEFENSE : BLOOD BUFFERS
• SECOND LINE OF DEFENSE :- RESPIRATORY
REGULATION
• THIRD LINE OF DEFENSE :RENAL REGULATION
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FIRST LINE OF DEFENSE : BLOOD
BUFFERS
• Buffer systems. Buffers act quickly to
temporarily bind H+ removing the highly
reactive, excess H+ from solution. Buffers thus
raise pH of body fluids but do not remove H+
from the body.
• Buffers function almost instantaneously
Buffers Are The1st Line Of Defense. They
Minimize (But Do Not Prevent) Changes In pH.
Buffer + H+ ↔ Hbuffer
Buffering of hydrogen Ions in the body fluids
•Bicarbonate buffer system
•Intracellular protein
•Hemoglobin Buffer system.
•Phosphate buffer system
Bicarbonate Buffer
• The most important buffer in plasma.
• 65% of buffering capacity.
• BASE CONSTITUENT :- (HCO3-)  Renal
Regulation
• ACID CONSTITUENT :- (H2CO3)  Respiratory
Regulation
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Bicarbonate buffer
CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3• Sodium Bicarbonate (NaHCO3) and carbonic acid
(H2CO3)
• Maintain a 20:1 ratio : HCO3- : H2CO3
HCl + NaHCO3 ↔ H2CO3 + NaCl ; {excess H2CO3 ,
excess CO2}
NaOH + H2CO3 ↔ NaHCO3 + H2O; { decre H2CO3 ,dec
CO2}
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• Normal bicarbonate level of plasma is
24mmol/L
• The normal pCO2 is 40mm Hg
• The normal carbonic acid concentration is 1.2
mmol/L
Remember these
values!!
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• pKa for carbonic acid is 6.1
• So, applying Henderson –Hasselbalch’s equation
pH= pKa + log [HCO3- ]
[H2CO3]
= 6.1 + log 24
1.2
= 6.1 + log 20
= 6.1 +1.3
= 7.4
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What Is The Central Message Of HendersonHasselbalch?
pH = pKa +
log(HCO3
/ H2CO3)
Plasma pH is a simple function of the
HCO3- :H2CO3 ratio
HCO3- : H2CO3 ↑ = pH ↑ (ALKALOSIS) :
Could occur due to either HCO3- ↑(Metabolic alkalo
or PCO2 ↓ (respiratory alkalosis)
HCO3- : H2CO3 ↓ = pH ↓( ACIDOSIS) :
Could occur either HCO3- ↓(metabolic acidosis)
or PCO2 ↑ (respiratory acidosis)
Phosphate buffer:
• Major intracellular buffer
• The main elements of the phosphate buffer
system are H2PO4– and HPO4=.
• H+ + HPO42- ↔ H2PO4-
• OH- + H2PO4- ↔ H2O + H2PO42-
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Protein Buffers
• Buffering capacity of protein dependson the
pKa value of the ionizable side chains.
• Includes hemoglobin
• In general ,
– Carboxyl group gives up H+
– Amino Group accepts H+
– Side chains that can buffer H+ are present on
amino acids.
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Protein Buffer System
• The free carboxyl group at one end of a
protein acts like an acid by releasing H+ when
pH rises; it dissociates as follows:
ACTION OF HEMOGLOBIN
• GENERATES BICARBONATE BY CARBONIC
ANHYDRASE
• In tissues :CO2 + H2O Carbonic Anhydrase
H2CO3
HCO3- + H+
H+ + HbHHb
H2CO3
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SECOND LINE OF DEFENSE :RESPIRATORY REGULATION
• Exhalation of carbon dioxide. By increasing the
rate and depth of breathing, more carbon
dioxide can be exhaled. Within minutes this
reduces the level of carbonic acid in blood,
which raises the blood pH (reduces blood H+
level).
• Respiratory mechanisms take several minutes
to hours
Respiratory mechanisms
•
•
•
•
•
•
2nd Line of Defence
Exhalation of carbon dioxide
Powerful, but only works with volatile acids
Doesn’t affect fixed acids like lactic acid
CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3Body pH can be adjusted by changing rate and
depth of breathing
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Volatile and
Non Volatile acid
secretion
Respiratory System is the Second Line of Defense
The peripheral
chemoreceptors also
respond to pH changes
caused by PCO2
changes, however they
directly monitor
changes in the arterial
blood, not the
cerebrospinal fluid as
the central
chemoreceptors do.
↑
↓
↑
↑
↑
↑
The peripheral
chemoreceptors also
respond to acids
such as lactic acid,
which is produced
during strenuous
exercise
↑
↑
↓
↑
Increased Hydrogen Ion Concentration
Stimulates Alveolar Ventilation
Increasing Alveolar Ventilation
Decreases Extracellular Fluid Hydrogen
Ion Concentration and Raises pH
THIRD LINE OF DEFENSE :RENAL
REGULATION
• Kidney excretion of H ion. The slowest
mechanism, but the only way to eliminate non
volatile acids, is through their excretion in
urine.
• Renal mechanisms may take several hours to
2-3 days.
• Changes are slow but powerful.
The Renal System mechanism
1. Regulation of plasma HCO32. Excretion of fixed (metabolic) acid load
…..Most of the time the urine is acidic to
balance metabolic acid production
MAJOR MECHANISMS OF RENAL REGULATION
1. Reabsorption of filtered HCO32. Excretion of fixed H+
Reabsorption of filtered HCO3-
• The process results in net reabsorption of
filtered HCO3-. However, it does not result in
net secretion of H+.
Regulation of reabsorption of filtered HCO3-
 Filtered load
 PCO2
 ECF volume
 Angiotensin II
2. Excretion of fixed H+
• Fixed H+ produced in the body is excreted by
two mechanisms,
• Simultaneously excreting urinary buffers
(titratable acid )
• Attaching H+ ion to ammonia (NH3) and
excrete it as (NH4+)
a. Excretion of H+ as titratable acid
(H2PO4-)
b. Excretion of H+ as NH4+
• In Alkalosis
• there is an excess of
HCO3– over H+ in the
tubular filterate, the
excess HCO3– cannot
be reabsorbed;
therefore, the excess
HCO3– is left in the
tubules and eventually
excreted into the urine,
which helps correct the
metabolic alkalosis.
• In Acidosis
• there is excess H+
relative to HCO3–,
causing complete
reabsorption of the
bicarbonate; the excess
H+ passes into the
urine.
• The excess H+ is
buffered in the tubules
by phosphate and exc
• And excreted as
ammonium ion
PLEASE REMEMBER !!!
Normal Values
pH
7.35 – 7.45
Bicarbonate
22-26mmol/L
Chloride
96-106mmol/L
Potassium
3.5-5mmol/L
Sodium
136-145mmol/L
pO2
95(85-100) mmHg
pCO2
40(35-45) mmHg
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COMA
CRAMPS
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Acid-Base Imbalances
• pH< 7.35 acidosis
• pH > 7.45 alkalosis
• The body response to acid-base imbalance is
called compensation
• May be complete if brought back within
normal limits
• Partial compensation if range is still outside
norms.
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Case #2
• 36 year old heroin addict found unresponsive
with needle in arm
• Pulse = 102, BP = 110/80, Temp = 35.2 C
• ABG(Arterial Blood Gas): PaO2 = 70, PaCO2 = 80,
• pH = 7.00, HCO3- = 23mEq/L
Respiratory Acidosis
• Carbonic acid excess caused by blood levels of
CO2 above 45 mm Hg.
• Hypercapnia – high levels of CO2 in blood
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Causes
DEPRESSION OF THE
RESPIRATORY CENTRE
• Head Injury
• Anaesthetics, sedatives
(morphine )
DECREASED
FUNCTIONING OF LUNGS
•
•
•
•
Pneumonia
Bronchitis
Asthma
Pneumothorax
• COPD (Emphysema)
• ARDS- Adult Respiratory Distress
Syndrome
• Motor neuron disease
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Treatment of Respiratory Acidosis
• IV lactate solution
• Treat underlying dysfunction or disease
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Question :- Why is lactate used ??
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Case #3
• 16 year old with closed head injury after a fall
from 15 feet
• P = 132, BP = 115/90,
• T = 37.2 C
• ABG: PaO2 = 110, PaCO2 = 26,
• pH = 7.52, HCO3- = 22
Respiratory Alkalosis
• Carbonic acid deficit
• pCO2 less than 35 mm Hg (hypocapnea)
• Most common acid-base imbalance
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Causes
• Hyperventilation(most common )
– Anxiety, Hysteria etc
• Conditions that stimulate respiratory center:
– Oxygen deficiency at high altitudes
– Pulmonary disease and Congestive heart failure – caused
by hypoxia
– Acute anxiety
– Fever, anemia
– Meningitis
– Cirrhosis
– Gram-negative sepsis
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Compensation of Respiratory Alkalosis
• Kidneys conserve hydrogen ion
• Excrete more bicarbonate ion( i.e less is
resorbed)
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Treatment of Respiratory Alkalosis
• Treat underlying cause
• Breathe into a paper bag
• IV Chloride containing solution – Cl- ions
replace lost bicarbonate ions
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Case #4
•
•
•
•
•
22 year old diabetic found unresponsive
P = 102, BP = 110/80,
T = 36.2 C
ABG: PaO2 = 90, PaCO2 = 36,
pH = 7.12, HCO3- = 8
Metabolic Acidosis
• Bicarbonate deficit - blood concentrations of
bicarb drop below 22mEq/L
• Causes:
– Loss of bicarbonate through diarrhea or renal
dysfunction(Type 2 RTA)
– Accumulation of acids (lactic acid or ketones)
– Failure of kidneys to excrete H+ (Type 1 and Type
4 RTA)
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Anion gap; Difference b/w measured cations and
measured anions .
76
• Actually the sum of CATIONS and ANIONS in
ECF is always equal.
• There is no gap whatsoever .
• The unmeasured anions constitute the anion
gap .( 10± 2mmol/L)
77
Anion Gap In Metabolic Acidosis
• Anion gap:
[Na+] - ([Cl-] + [HCO3-]) = 8-12 mmol/L
• If > 18, there are unmeasured anions, such as:
–
–
–
–
–
lactate
ketones
salicylate
ethanol
ethylene glycol (anti-freeze)
M
U
D
P
I
L
E
S
High Anion-Gap Acidosis
MUDPILES (methanol,
uremia, diabetic
•Diabetic ketoacidosis
•Starvation ketoacidosis ketoacidosis, propylene
glycol, isoniazid, lactic
2. Lactic Acidosis
acidosis, ethylene glycol,
salicylates)
1. Ketoacidosis
3. Renal Failure- Excretion of H+ and regeneration of HCO3DEFICIENT
4. Toxins
•Ethylene glycol
•Methanol
•Salicylates
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Normal Anion-Gap Acidosis
(Loss of both CATIONS AND ANIONS)
1. Renal Causes
•Renal tubular acidosis
•Carbonic anhydrase inhibitors
2. GIT Causes
•Severe diarrhoea
•Uretero-enterostomy or Obstructed ileal conduit
•Drainage of pancreatic or biliary secretions
•Small bowel fistula
3. Other Causes
•Addition of HCl, NH4Cl
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Compensation for Metabolic Acidosis
•
•
•
•
•
Increased ventilation- to decrease volatile acid
Increased reapsorption of HCO3- by kidneys
Renal excretion of hydrogen ions if possible
K+ exchanges with excess H+ in ECF
( H+ into cells, K+ out of cells)
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Case #5
• 6 week old infant is lethargic with history of
vomiting increasing for 1 week
• P = 122, BP = 85/60,
• T = 37.2 C
• ABG analysis: PaO2 = 90, PaCO2 = 44,
• pH = 7.62, HCO3- = 36,
Metabolic Alkalosis
• Bicarbonate excess - concentration in blood
is greater than 26 mEq/L
• Causes:
– Excess vomiting = loss of stomach acid
– Excessive use of alkaline drugs,antacids(NaHCO3)
– Excess aldosterone
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Diagnosis of Acid-Base Imbalances
1. Note whether the pH is low (acidosis) or high
(alkalosis)
2. Decide which value, pCO2 or HCO3- , is outside
the normal range and could be the cause of the
problem.
•
•
If the cause is a change in pCO2, the problem is
respiratory.
If the cause is HCO3- the problem is metabolic.
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3. Look at the value that doesn’t correspond to the
observed pH change.
If it is inside the normal range, there is no
compensation occurring.
If it is outside the normal range, the body is
partially compensating for the problem.
90
Example
• A patient is in intensive care because he
suffered a severe myocardial infarction 3 days
ago. The lab reports the following values from
an arterial blood sample:
– pH 7.3
– HCO3- = 20 mEq / L ( 22 - 26)
– pCO2 = 32 mm Hg (35 - 45)
91
Diagnosis
• Metabolic acidosis
• With compensation
92
Q
• pH 7.58;
• Pa.CO2 23 mm Hg;
• [HCO3-] 18 mEq/L
acid base condition
pH 7.58; Pa.CO2 23 mm Hg; [HCO3-] 18 mEq/L
1. Look at pH (is it acidosis or alkalosis?)
pH = 7.58  alkalosis
2. Look at HCO3- (is it metabolic alkalosis?)
HCO3- = 18 mEq/L (normal 22-30)  not metabolic
alkalosis
3. Look at Pa.CO2 (is it respiratory alkalosis?)
Pa.CO2 = 23 mmHg (normal 35-45)  respiratory
alkalosis
4. See if appropriate compensation has occurred:
compensation for respiratory alkalosis is HCO3excretion
HCO3- = 18 mmHg (normal 22-30)
 partially compensated respiratory alkalosis
94
Case F2: acid base condition
pH 7.29; Pa.CO2 26 mm Hg; [HCO3-] 12 mEq/L
1. Look at pH (is it acidosis or alkalosis?)
pH = 7.29  acidosis
2. Look at HCO3- (is it metabolic acidosis?)
HCO3- = 12 mEq/L (normal 22-30)  metabolic
acidosis
3. Look at Pa.CO2 (is it respiratory acidosis?)
Pa.CO2 = 26 mmHg (normal 35-45)  not resp.
acidosis
4. See if appropriate compensation has occurred:
compensation for metabolic acidosis is
hyperventilation
Pa.CO2 = 26 mmHg (normal 35-45); partial
compensation
96
• In the patient described , which of the
following laboratory results would be
expected, compared with normal?
A) Increased renal excretion of HCO3B) Decreased urinary titratable acid
C) Increased urine pH
D) Increased renal excretion of NH4+
Mixed disturbances
Here several problems of acid-base
management are colliding at the same time.
It’s definitely not just a matter of the body
trying to compensate for one such disorder.
a. An example would be a DIABETIC with
KETOACIDOSIS, who also happens to
have C.O.P.D, or develops a bad PNEUMONIA
(and as a result develops a respiratory
acidosis.)
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Davenport diagram showing the relationships
among HCO3-, pH, and PCO2.
Normal buffer line BAC
Davenport diagram showing the relationships among HCO3,
pH, and PCO2. .
B shows the changes/compensation occurring in respiratory
and metabolic acidosis and alkalosis
109
Which point on the graph below would most likely represent the
systemic arterial blood of a mountain climber after several weeks
at high altitude?
Which arrow on the graph below could represent the
change in status of an individual with metabolic
acidosis who was then given an intravenous injection of
sodium bicarbonate?
Siggard Andersen Normogram
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