Acid-base balance. pH homeostasis.
21
ACID-BASE BALANCE. PH HOMEOSTASIS.
INTRODUCTION
An important property of the internal environment (and also the blood) is its
degree of acidity or alkalinity. Acidity increases when the level of acidic
compounds in the body rises (through increased intake or production, or
decreased elimination), or when the level of basic (alkaline) compounds in the
body falls (through decreased intake or production, or increased elimination).
Alkalinity increases with the reverse of these processes. The body's balance
between acidity and alkalinity is referred to as acid-base balance or acid-base
homeostasis. The acidity or alkalinity of any solution, including blood, is
indicated on the pH scale.
Human homeostasis refers to the body's ability to physiologically
regulate its inner environment to ensure its stability in response to
fluctuations in the outside environment. Several organs help maintain
homeostasis including the liver, the lungs, the kidneys, the autonomic
nervous system and the endocrine system. An inability to maintain
homeostasis may lead to death or a disease, a condition known as
homeostatic imbalance.
The blood's acid-base balance is precisely (tightly) controlled, because even a
minor deviation from the normal range can severely affect many organs. The
body uses several different mechanisms to control the blood's acid-base balance.
CHEMICAL BACKGROUND
The hydrogen atom consists of only one proton and one electron. When
dissociated (i.e. oxidation of hydrogen) in the sense of removing its electron,
formally gives H+, containing no electrons and a nucleus composed of one
proton. That is why H+ is often called a proton.
Distilled water is water that has virtually all of its impurities removed through
distillation. Water, however pure, is not a simple collection of H2O molecules.
Even in "pure" water, sensitive equipment can detect a very slight electrical
conductivity. This must be due to the presence of ions H+ and OH- (or H3O+ and
OH-).
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Physiology laboratory exercises
A bare proton, H+, cannot exist in solution or crystals, because of its
attraction to other atoms or molecules with electrons. Except at the high
temperatures associated with plasma, such protons cannot be removed
from the electron clouds of atoms, and will remain attached to them.
However, the term "proton" is used loosely and metaphorically to refer to
positively charged hydrogen attached to other species (e.g. H2O and as
such is denoted "H+" without any implication that any single protons
exist freely as a species.
The dissociation constant (the ratio of dissociated/non-dissociated molecules) of
water is described by (1) and this can be determined by several experimental
methods. The water ionization constant (i.e. ionic product of water) is described
by (2) and can be exactly computed. The ratio of the two ions is represented by
(3). Because distilled water is made only from water molecules that dissociate to
H+ and OH- ions, their concentrations must be equal (3).
ka 
H  OH 


H 2 O
(1)
  

k w  H   OH   10 14 mol / l
(2)
H   1  H   OH 
OH 
(3)




Thus the concentration of H+ in distilled water will be 10-7 mol/l
The pH value is the negative logarithm (base 10) of the molar concentration of dissolved
hydrogen ions (4). Thus pH will have typical values between 0-14, 7.00 being
considered neutral (i.e. the concentration of H+ at 25°C is approximately
1.0×10−7 mol/l).
 
pH   log H 
The concept of pH was first introduced by Danish chemist Søren Peder
Lauritz Sørensen at the Carlsberg Laboratory in 1909. The p probably
stands for “Power” although this is debated.
Acid-base balance. pH homeostasis.
23
In order to maintain the body's acid-base balance buffer solutions are used. These
have the property that the pH of the solution changes very little when a small amount of
strong acid or base is added to it. Usually buffer solutions are aqueous solutions
consisting of a mixture of a weak acid and its conjugate base or a weak base and
its conjugate acid. In conclusion buffer solutions reversibly bind hydrogen ions
and impede any change in pH.
When added to water a strong acid will dissociate completely:
 
HCl  H   Cl   H    pH 
When added to water a strong base will dissociate completely:


NaOH  Na   OH   OH    pH 
Thus when adding a strong base or acid to a solution due to the complete
dissociation there will be massive changes in the pH.
If we mix an acid and a base we will obtain a salt:
HCl  NaOH  H 2 O  NaCl
When mixing an acid with the salt created by the reaction between a weak
acid (i.e. an acid that dissociates incompletely) and a strong base the
result will be a weak acid, with less dissociation and thus a smaller
modification of the pH. In the example below we obtained acetic acid with
a very low dissociation constant, and as such no major effect on pH.
CH3-COOH ↔ CH3-COO- + H+
CH 3COONa  HCl  CH 3COOH  NaCl
If we mix further this weak acid (acetic acid) with its salt (sodium
acetate) the buffer capacity increases (i.e. the pH will not be modified
significantly if we add strong acids or strong bases to the solution). Thus
buffer solutions are made up from weak acids and their salts (ex.
CH3COOH / CH3COONa).
CH 3 COOH  NaOH  CH 3COONa  H 2 O
CH 3 COONa  HCl  CH 3 COOH  NaCl
BUFFER SOLUTIONS OF THE HUMAN BODY
The human body has several buffer systems. Some of these are located in the
blood and extracellular space:
– bicarbonate (H2CO3/NaHCO3, carbonic acid/sodium-bicarbonate)
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Physiology laboratory exercises
– hemoglobin
Other buffer systems are found in the intracellular space:
– phosphate (Na2HPO4/NaH2PO4, disodium hydro-phosphate/sodium
dihydro-phosphate)
– proteins
From these buffers one of the most important is the bicarbonate buffer system as
it:
– reversibly binds hydrogen ions thus preventing pH changes
– can shift carbon dioxide through carbonic acid to hydrogen ions and
bicarbonate and vice-versa
– allows quick buffering (although low buffer capacity)
– acid-base imbalances that overcome this buffer system can be
compensated in the short term by:
– changing the rate of ventilation; this alters the concentration of
carbon dioxide in the blood, shifting the reaction (see below), which in
turn alters the pH (ex. if pH drops it can be compensated through
increased breathing, thus eliminating the CO2)
– excretion of excess acid or base; the kidneys are slower to
compensate, but there are several powerful renal mechanisms to control
pH

HCO3  H   H 2 CO3  CO2  H 2 O
The proteins have the highest buffering capacity of all buffer solutions of the
human body but this is a slow system.
The principle mechanisms to control pH are: dilution (i.e. diluting any acid or
base in the large volume of water in the organism), buffering (i.e. buffering using
buffer solutions with involvement of the respiratory and renal systems) and
restoration of normal state.
PH OF THE HUMAN INNER ENVIRONMENT
Probably the simplest way to calculate the pH is to use the HendersonHasselbalch equation and apply the discussed chemical background to the
bicarbonate buffer system.
Acid-base balance. pH homeostasis.
25
H  HCO 
k


H 2 CO3 
log k  log
3
H  HCO   logH   log HCO 
H CO 
H CO 


3
2
 


3
3
HCO   log k
2
3

 log H   log
 
HCO 

3
H 2 CO3 
 pH  pk  log
pH   log H  and  log k  pk
3
H 2 CO3 
which is the Henderson-Hasselbalch equation (where pk is the dissociation
constant of H2CO3)
Lawrence Joseph Henderson wrote an equation, in 1908, describing the
use of carbonic acid as a buffer solution. Karl Albert Hasselbalch later reexpressed that formula in logarithmic terms, resulting in the Henderson–
Hasselbalch equation.
Because the bicarbonate buffer is in open system through the lungs the
concentration of H2CO3 will be dependent on the elimination of CO2; in fact it
depends on the solubility coefficient of CO2 (a) and its partial pressure (pCO2 the partial pressure of a gas dissolved in a liquid is the partial pressure of that
gas, which would be generated in a gas phase in equilibrium with the liquid at
the same temperature). Thus the Henderson-Hasselbalch equation can be
modified to reflect this.
HCO   pk  log HCO 
pH  pk  log


3
H 2 CO3 
a  0.03
3
a  pCO2


mEq

; pCO2  40 mmHg ; HCO3  24 mEq / l ;
l  mmHg
pk  6.1
 pH  6.1  log 20  6.1  log 10  log 2  6.1  1  0.3  7.4
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Physiology laboratory exercises
There are some significant approximations implicit in the Henderson–
Hasselbalch equation. The most significant is the assumption that the
concentration of the acid and its conjugate base at equilibrium will
remain the same as the former concentration. This neglects the
dissociation of the acid and the hydrolysis of the base. The dissociation of
water itself is neglected as well. Also, the equation does not take into
effect the dilution factor of the acid and conjugate base in water.
ACID-BASE BALANCE PARAMETERS
► pH: is the measure of the acidity or alkalinity of a solution. In arterial
blood the normal range is 7.36-7.44. Blood pH values compatible with life
in mammals are limited to a pH range between 6.8 and 7.8.
► pCO2: partial pressure of carbon dioxide; its value is dependent on
ventilation (breathing). The brain regulates the amount of carbon dioxide
that is exhaled by controlling the speed and depth of breathing. The
amount of carbon dioxide exhaled, and consequently the pH of the blood,
increases as breathing becomes faster and deeper. By adjusting the speed
and depth of breathing, the brain and lungs are able to regulate the blood
pH minute by minute. In arterial blood the normal value is 40-44 mmHg.
► standard HCO3- (standard bicarbonate): concentration of bicarbonate
at a pCO2 of 40 mmHg, oxygenation 100% and t=37°C. Normal value in
arterial blood is 24 mEq/l.
► total CO2: in arterial blood the normal value is 25 mEq/l (= [HCO3] + 0.03 x pCO2)
► reserve alkalinity: the concentration of bases in plasma after the
neutralization of stronger acids than carbonic acid; normal range in
venous blood is 24-29 mmol/l.
► buffer base: is the sum of all buffering agents in the blood, normal
value 46 mEq/l.
Normal buffer base (NBB) is what the buffer base would have been at
normal pH (7.4), normal pCO2 (40 mmHg) and normal temperature
(37°C).
The content of hemoglobin in blood increases this value.
Acid-base balance. pH homeostasis.
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CO2  H 2O  H 2 CO3
H 2 CO3  Hb  HHb  HCO3

The correlation (empirical formula) is: NBB = 41.7+0.68×Hb, where Hb is
in mmol/l (the normal value hemoglobin is ~9.3 mmol/l).
► base excess: refers to an excess or deficit in the amount of base present
in the blood. The value is usually reported as a concentration in units of
mEq/l, with positive numbers indicating an excess of base and negative a
deficit. The reference range for base excess is −2 to +2 mEq/l.
Comparison of the base excess with the reference range assists in
determining whether an acid/base disturbance is caused by a respiratory,
metabolic, or mixed metabolic/respiratory problem. While carbon
dioxide defines the respiratory component of acid-base balance, base
excess defines the metabolic component. Accordingly, measurement of
base excess is defined under a standardized pressure of carbon dioxide,
by titrating back to a standardized blood pH of 7.40.
This parameter can be used to calculate the necessary amount of sodium
bicarbonate to neutralize the extracellular space (in acid build-up).
NaHCO3 = BEmeasured x 0.3 x body weight (kg) - for adults (water in
extracellular space is 1/3 of body weight)
NaHCO3 = BEmeasured x 0.4 x body weight (kg) - for children (water in
extracellular space is 40% of body weight)
► anion gap: it is calculated by subtracting the serum concentrations of
chloride and bicarbonate (anions) from the concentrations of sodium plus
potassium (cations).
Na+
140 mmol/l
Cl-
K+
4 mmol/l
144 mmol/l
HCO3-
110
mmol/l
24 mmol/l
134 mmol/l
If there are more unmeasured anions (compared to unmeasured cations)
in the serum the anion gap is higher (i.e. higher than the normal value of
10). The anion gap varies in response to changes in the concentrations of
the above mentioned serum components that contribute to the acid-base
balance and increases when fixed (non-volatile) acids accumulate in the
organism.
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Physiology laboratory exercises
DEFINITIONS
Acidosis – increased acidity due to excess acid or deficit of bases.
Acidemia means low blood pH (i.e. pH < 7.36); acidosis is used to
describe the processes leading to increased acidity. Incorrectly physicians
use these terms in an interchangeable fashion.
Alkalosis – reduced hydrogen concentration of arterial blood plasma
Metabolic – acid-base disturbance caused by deficit or build-up of acids
or bases in the extracellular space
Respiratory – acid-base disturbance caused by modification of ventilation
E.g. metabolic acidosis means: excessive blood acidity caused by an
overabundance of acid in the blood or a loss of bicarbonate from the blood.
E.g. respiratory acidosis means: buildup of carbon dioxide in the blood
that results from poor lung function or slow breathing
Simple acid-base disorder– one single etiological base, either metabolic
or respiratory disturbance (i.e. the presence of only one of the possible
derangements)
Mixed disorder– more than one cause
Compensated – the disturbance of one system (i.e. renal or respiratory)
altered the pH, this can be compensated by the opposite system, and thus
the pH can be corrected (the pH returns to normal). E.g. hypoventilation
causes the build-up of carbon dioxide and in turn the decrease of pH (i.e.
respiratory acidosis), this will be compensated by the increase of
bicarbonate to buffer the free H+. If the compensation is complete then the
pH is normal but both the pCO2 and bicarbonate concentrations are
altered.
Partially compensated – similar to the above, but the compensation is not
enough to return the pH to normal value.
Not compensated – there is no compensation of the acid-base
disturbance.
ACID-BASE DISTURBANCES
Respiratory acidosis - caused by diseases that increase the pCO2, which in term
will increase the hydrogen ion concentration (i.e. decrease the pH). There is no
primary modification of the concentration of bicarbonate.
Acid-base balance. pH homeostasis.
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Metabolic acidosis - caused by diseases that decrease the concentration of
bicarbonate, thus the concentration of hydrogen ions increase (i.e. decrease the
pH). There is no primary modification of the partial pressure of carbon dioxide.
Respiratory alkalosis - caused by diseases that decrease the pCO2, which in term
will decrease the hydrogen ion concentration (i.e. increase the pH). There is no
primary modification of the concentration of bicarbonate.
Metabolic alkalosis - caused by diseases that increase the concentration of
bicarbonate, thus the concentration of hydrogen ions decrease (i.e. increase the
pH). There is no primary modification of the partial pressure of carbon dioxide.
Use the Siggaard-Andersen nomogram and the given examples to
determine the acid-base balance parameters.
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Physiology laboratory exercises