Download BIOL 2402 Acid Base Homeostasis

11/22/15
Dr. Chris Doumen
Collin College
BIOL 2402
Acid Base Homeostasis
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Acid -Base Balance and Regulation
Acid-Base Balance refers to the precise regulation of free
hydrogen ion concentration in the body fluids
Free hydrogen ions determine the acidity of the body fluids and
pH is used as a specific H+ indicator
pH = -log [H+]
H2O
H+ + OH-
where [H+] is the molar concentration
[H+] = 10-7 M
[H2O] = 55 M
pH = - log [10-7 ] = 7
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Acid -Base Balance and Regulation
In physiology, 7.4 is considered neutral because it reflects
the average blood pH ( concentration of H+ = 40 nM; compare
that with [Na+] = 140 mM … )
•  Acidosis or acidemia: A blood pH below 7.35
•  Alkalosis or alkalemia : A blood pH above 7.45
Death occurs within seconds when blood pH falls below
6.8 or above 8.0.
Regulation of blood pH, more specifically H+, is thus a
very important homeostatic factor in life.
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Acid -Base Balance and Regulation
Sources of H+ in the body
VOLATILE ACIDS : Carbon dioxide and carbonic acid
•  CO2 and H2CO3
FIXED ACIDS : Inorganic acids (non carbonic acids ) from diet
•  phosphoric acid, sulfuric acid, ammonia
•  ~ 1 - 1.5 mmoles of H+ /kg/day
ORGANIC ACIDS : resulting from metabolism
•  citric acid, lactic acid, pyruvic acid, ketone body acids
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Acid -Base Balance and Regulation
Quantities of H+ produced per day
Diet : 70 000 000 nmoles/day
Metabolism : 5 000 000 000 nmoles/day
But if we calculate the amount of free protons in the blood
stream ( = pH 7.4) is comes out to be 40 nmol /L ( = 10-7.4 M)
Our body is thus constantly challenged by an enormous
overload of protons. It requires different mechanism to keep
the proton levels under control. Small uncontrolled changes in
pH can have drastic effects on our metabolic sanity.
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Lines of defense against changes in pH
The body’s different mechanisms to deal with changes in
protons are 3 fold :
•  Chemical Buffers
•  Respiratory mechanism
•  Renal mechanism
Chemical buffers act immediately ; they bind excess protons
( or release protons) but do not eliminate H+ from the body.
They can only soak up extra H+ depending on the concentration
of the chemical buffers present. When capacity is full, the
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additional H+ needs to be removed from the body
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Lines of defense against changes in pH
The lungs and kidneys aid in the removal of acids from the
body. They act however more slowly, with the lungs being
faster compared to the kidneys.
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Lines of defense against changes in pH
1. Chemical Buffers
Chemical Buffers are composed out of compounds that
minimize pH changes when acids or bases are added.
The compounds come in pairs :
•  a weak acid (HY below) : releases H+
•  a weak base (Y- below): binds H+
HY
Weak acid
H+ + Y Weak base
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1. Chemical Buffers
HY
Weak acid
H+ + Y Weak base
These reactions have a certain equilibrium status and the direction
of the reactions are influenced by the presence of their compounds
on either side of the arrows : if there is a sudden increase in HY, the
reaction will proceed to the right and more H+ will be created
On the other hand, if more protons are added, the reaction will
proceed to the left and the extra prtons will be removed from a
solution by the weak base, Y9
Lines of defense against changes in pH
A. The most important buffer in ICF (inside cells) are the
proteins
Proteins contain both acid and basic groups and can thus
bind H+ and/or release protons quite easily.
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•  If pH climbs, the carboxyl group of amino acid acts
as a weak acid
•  If the pH drops, the amino group acts as a weak
base
Hemoglobin in RBC is important in binding H+ produced by
tissues and thus in buffering blood pH. It is the only intracellular
buffer system with an immediate effect on ECF pH .
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Lines of defense against changes in pH
B. The most important buffer in ECF is the bicarbonate buffer
CO2 + H2O
H2CO3
H+ + HCO3-
•  Has the following limitations:
•  It cannot protect the ECF from pH changes due
to increased or depressed CO2 levels
•  Only functions when respiratory system and
control centers are working normally
•  It is limited by availability of bicarbonate ions
(bicarbonate reserve)
•  Bicarbonate ion shortage is rare Due to large
reserve of sodium bicarbonate
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c. Phosphate buffer is an important intracellular and
urinary buffer system
Na2HPO4 + H+
NaH2PO4 + Na+
Humans consume more phosphate than needed. The excess is filtered
into the nephrons and is not re-absorbed by the kidney.
The phosphate helps to buffer urine pH in the nephron. It binds the
secreted protons and keeps the pH above 5. If it were not for this
buffer, urine pH would be extremely acidic very fast ( below 4.5) and
prevent the nephron from secreting H+ .
Since phosphates are used extensively in side the cell, this buffer
system also contributes to intracellular buffering of pH.
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Buffer Systems
occur in
Intracellular fluid (ICF)
Phosphate Buffer
System
The phosphate
buffer system has
an important role in
buffering the pH of
the ICF and of urine.
Extracellular fluid (ECF)
Protein Buffer Systems
Protein buffer systems contribute to the regulation of pH in the
ECF and ICF. These buffer systems interact extensively with the
other two buffer systems.
Hemoglobin buffer
system (RBCs only)
Amino acid buffers
(All proteins)
Carbonic Acid–
Bicarbonate Buffer System
The carbonic acid–
bicarbonate buffer
system is most
important in the ECF.
Plasma protein
buffers
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Lines of defense against changes in pH
•  Limitations of Chemical Buffer Systems
•  Provide only temporary solution to acid–base imbalance
•  They do not eliminate H+ ions
•  Supply of chemical buffer molecules is limited
•  Maintenance of Acid–Base Balance
•  For homeostasis to be preserved, captured H+ must:
•  Be permanently tied up in water molecules
•  Through CO2 removal at lungs
•  Be removed from body fluids
•  Through secretion at kidney
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2. Respiratory Mechanism of H+ regulation
Regulation occurs by CO2 removal and involves the
bicarbonate reaction
•  If not enough CO2 is expelled by the lungs, more CO2
stays behind in the blood
•  CO2 drives the bicarbonate reaction to the left and forms
more Bicarbonate and protons and pH drops
CO2 + H2O
H2CO3
H+ + HCO3-
•  The opposite occurs when too much CO2 is expelled
CO2 + H2O
H2CO3
H+ + HCO317
Lines of defense against changes in pH
This reaction is so important that it can provide direct information on the status of
the body by analyzing the components of this reaction in the blood stream.
Lets look at this reaction in a simplified version and think in acid-base terms.
H+ + HCO3-
CO2 + H2O
Weak acid
Weak base
Every equilibrium reaction has an equilibrium set point know as the Keq value .
The K value is the ratio of products over substrates at equilibrium ( the value
of water is ignored since it is huge compared to the others).
Keq =
[H+] x [HCO3-]
[H2O] x [CO2]
=
[H+] x [HCO3-]
[CO2]
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1) Let’s now take the log of this equation
log Keq = log
[H+] x [HCO3-]
[CO2]
:
= log [H+] + log [HCO3-]/[CO2]
2) Swap log Keq and log [H+] around and we get:
-log [H+] = -log Keq + log [HCO3-]/[CO2]
3) Since -log [H+] = pH and define pKa as -log Keq we obtain :
pH= pKa + log [HCO3-]/[CO2]
Or in general when dealing with weak acids and bases
pH = pKa + log{[Weak Base]/[ Weak Acid]}
Lines of defense against changes in pH
This is known as the Henderson-Hasselbalch equation for
weak acids/bases and it can be used to determine pH levels
pH = pKa + log {[Base]/[Acid]}
pH = pKa + log {[HCO3-]/[CO2]}
Since the CO2 levels in blood relate to the partial gas pressure for
CO2 in blood, the concentration of CO2 in blood at body temp. is
0.03 times the partial pressure for CO2.
pH = pKa + log {[HCO3-]/0.03 PCO2}
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pH = pKa + log {[HCO3-]/0.03 PCO2}
Healthy blood pH=7.4 and pKa for Bicarbonate reaction is a
constant equal to 6.1
If we substitute now we get 7.4 = 6.1 + log {[HCO3-]/0.03 PCO2}
Or
7.4 -6.1 = 1.3 = log {[HCO3-]/0.03 PCO2}
You can get rid of a log expression by putting it to the power 10.
10(pH - pKa) = 10(1.3) =[HCO3-]/0.03 PCO2 = 20
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So the ratio of [HCO3-] to {0.03 PCO2 } determines blood pH !
And since 101.3 = 20, the ratio of [HCO3-] to {0.03 PCO2 } in
healthy blood should be in a ration of 20 : 1 !
The respiratory system uses this to adjust pH by regulating CO2.
Changes can occur within minutes ; changing AVR by 2 ( or 1/2) will change
pH of the blood by 0.2 units
Anything that impairs respiratory system may thus affect the
acid-base balance of the body.
When a change in acid-base balance is due to a problem with the
respiratory system, it is referred to as Respiratory Acidosis or
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Respiratory Alkalosis.
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Figure 23-26a The Chemoreceptor Response to Changes in PCO2.
Increased
arterial PCO2
a
An increase in arterial PCO2
stimulates chemoreceptors
that accelerate breathing
cycles at the inspiratory center.
This change increases the
respiratory rate, encourages
CO2 loss at the lungs, and
decreases arterial PCO2.
Stimulation
of arterial
chemoreceptors
Stimulation of
respiratory muscles
Increased PCO2,
decreased pH
in CSF
Stimulation of CSF
chemoreceptors at
medulla oblongata
Increased respiratory
rate with increased
elimination of CO2 at
alveoli
HOMEOSTASIS
DISTURBED
Increased
arterial PCO2
(hypercapnia)
HOMEOSTASIS
RESTORED
HOMEOSTASIS
Normal
arterial PCO2
Start
Normal
arterial PCO2
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