UPDATES.ACID-BASE BALANCE.pptx

DR BILIAMINU S.A (FMCPath; Cert. in Clin. Embryology)
CONSULTANT CHEMICAL
PATHOLOGIST/CLINICAL EMBRYOLOGIST
DEPT. OF CHEMICAL PATHOLOGY/ART UNIT,
UITH; ILORIN.

OUTLINE
 Introduction
 Definitions
 Buffers
 Chief buffers in the blood
 Bicarbonate
 Determinants of pH
 Buffering
 Acid-Base Disturbances and its effects
 Anion gap
 Conclusion.

INTRODUCTION
Physicians generally agree that acid-base
balance is important but they struggle to
understand the science, pathology as well
as its application. This is one of the
reasons why acid-base chemistry has
occupied a special corner in clinical
medicine for more than 100 years now.

INTRODUCTION CONTD.
Undoubtedly, the body controls carefully
the relative concentrations of hydrogen
and hydroxyl ions in the intracellular and
extracellular spaces.
Alterations in this ‘’balance’’ distrupts
transcellular ion pumps leading to
significant cardiovascular problems
which may lead to death.

INTRODUCTION CONTD.
In fact it is not a gain-saying that no death
could hardly occur without acid-base
disturbance being involved.
Most acid-base abnormalities are easily
explained, but some remained proble-
matic.

INTRODUCTION CONTD.
Confusion exists regarding cause, effect
and treatment of acid-base abnormalities
since traditional teaching emphasizes data
interpretation rather than pathophy-
siology.

DEFINITIONS
 Acid-base balance can be defined as homeostasis of the
body fluids at a normal arterial blood pH ranging between
7.37 and 7.43.
 Acid: is a word from the Latin acidus/acēre meaning
sour is a substance which reacts with a base. Commonly,
acids can be identified as tasting sour, reacting with metals
such as calcium, and bases like sodium carbonate.
Aqueous acids have a pH of less than 7, where an acid of
lower pH is typically stronger. Chemicals or substances
having the property of an acid are said to be acidic.

DEFINITIONS CONTD.
 There are three common definitions for acids: the
Arrhenius definition, the Brønsted-Lowry definition, and
the Lewis definition.
 Arrhenius acids
 The Swedish chemist Svante Arrhenius attributed the
properties of acidity to hydrogen in 1884. An Arrhenius
acid is a substance that increases the concentration of the
hydronium ion, H3O+, when dissolved in water. This
definition stems from the equilibrium dissociation of
water into hydronium and hydroxide (OH−) ions:
 H2O(l) + H2O(l) H3O+(aq) + OH−(aq)

DEFINITIONS CONTD.
Brønsted-Lowry acids
 While the Arrhenius concept is useful for describing many
reactions, it is also quite limited in its scope. In 1923
chemists Johannes Nicolaus Brønsted and Thomas Martin
Lowry independently recognized that acid-base reactions
involve the transfer of a proton. A Brønsted-Lowry acid
(or simply Brønsted acid) is a species that donates a
proton to a Brønsted-Lowry base. Brønsted-Lowry acid-
base theory has several advantages over Arrhenius theory.

DEFINITIONS CONTD.
Lewis acids
 A third concept was proposed in 1923 by Gilbert N. Lewis
which includes reactions with acid-base characteristics
that do not involve a proton transfer. A Lewis acid is a
species that accepts a pair of electrons from another
species; in other words, it is an electron pair acceptor.

DEFINITIONS CONTD.
Base: In chemistry this is a substance that
can accept hydrogen ions (protons) or more
generally, donate electron pairs. A soluble
base is referred to as an alkali if it contains
and releases hydroxide ions (OH−)
quantitatively.
The Brønsted-Lowry theory defines bases as
proton (hydrogen ion) acceptors, while the
more general Lewis theory defines

DEFINITIONS CONTD.
bases as electron pair donors, allowing other
Lewis acids than protons to be included. The
oldest Arrhenius theory defines bases as
hydroxide anions, which is strictly applicable
only to alkali. In water, by altering the
autoionization equilibrium, bases give
solutions with a hydrogen ion activity lower
than that of pure water, i.e. a pH higher than
7.0 at standard conditions.

DEFINITIONS CONTD.
Arrhenius bases are water-soluble and
these solutions always have a pH greater
than 7 at standard conditions. An alkali is
a special example of a base, where in an
aqueous environment, hydroxide ions are
donated. There are other more generalized
and advanced definitions of acids and
bases.

DEFINITIONS CONTD.
 pH: The pH is the negative logarithm of the hydrogen ion
concentration. A complete definition requires that the
logarithm is defined as being to the base ten and the
concentration be measured as activity in moles per litre.
Confusion arises because, as the acidity increases, the pH
decreases. To avoid mistakes when discussing acid-base
balance, it is often safer to avoid "increase" and
"decrease" and use "more acid" and "more alkaline"
instead.
 Log (Base 10) of [ H+ ] in Mol/L

DEFINITIONS CONTD.
 When pH = 7, the H+ concentration is 10-7 or 1/107
 This is neutral because the H+ and OH- concentration is the
same.
 H2O <=> H+ + OH-
 When the pH = 1, the H+ concentration is 10-1 or 1/10. This is
a very strong acid.
 pH 7.00 = neutral
 pH >7 = alkaline
 pH <7 = acid
 pH 7.4 = physiological pH of extracellular fluid. (Range of
normal 7.35 - 7.45.)

DEFINITIONS CONTD.
 The starting point is the Henderson Equation:
 [H+] x [HCO3
-] = K x [CO2] x [H2O]
 Hasselbalch modified Henderson's elegant idea
by regarding the water concentration as
constant and taking logarithms of the
remaining components (pK is the negative
logarithm of "K"). This resulted in the
Henderson-Hasselbalch Equation:
 pH = pK + log ( [HCO3
-] / [CO2] )

DEFINITIONS CONTD.
 The consequence of using negative logarithms is that
"everything is upside down" and incomprehensible to
most physicians; it contains the same information as
Henderson's simple equilibrium equation. It could have
been so much easier; the conversion could have been
applied to the whole equation at once. The first step is to
write Henderson's equation in the right order with the
water concentration omitted as a constant.
 [H+] = K x [CO2] / [HCO3
-]

DEFINITIONS CONTD.
 The second is to take the negative logarithm: pH = -Log (
K x [CO2] / [HCO3
-] )
 The "K" is still "K" and the equation is still recognizable.
Why, then, have generations of medical students been
taught the Henderson-Hasselbalch version? Why, in fact,
were we taught it at all? Were our teachers so
mathematically naive that they failed to recognize that the
two equations were mathematically equivalent. If so, did
they succumb to the temptation to teach us - and therefore
test us - using the more complex version?

DEFINITIONS CONTD.
 Part of the reason lies outside medicine; chemists find
knowing the negative logarithm of "K" (pK) is a useful
shorthand way of writing a long number. In addition, the
same logarithmic version is in widespread use, although it
is known by other names in other places. At the Royal
Veterinary and Agricultural University of Copenhagen it is
known as the "Bjerrum equation" in honor of Professor
Bjerrum who worked there;

DEFINITIONS CONTD.
and, in the chemical world it is generally
known as the buffer equation.
There is no need in physiology for us to
use this equation. It is part of history's
legacy. The Modified Henderson
Equation is recommended.

DEFINITIONS CONTD.
Acidaemia: refers to a condition increasing
hydrogen ion concentration of arterial blood
plasma or clinical condition in which
bicarbonate concentration falls in the blood.
Acidosis: this refers to clinical manifestation
of acidosis. This is described as the condition
in which the pH falls below 7.35.

DEFINITIONS CONTD.
 Alkalaemia: refers to a condition reducing
hydrogen ion concentration of arterial blood
plasma or clinical condition in which
bicarbonate accumulates in the blood.
 Alkalosis: this refers to clinical manifestation
of alkalaemia.
 Generally, alkalosis is said to occur when pH
of the blood exceeds 7.45.

IMPORTANCE OF HYDROGEN
ION CONCENTRATION
Hydrogen ion concentration has a
widespread effect on the function of the
body's enzyme systems. The hydrogen ion
is highly reactive and will combine with
bases or negatively charged ions at very
low concentrations. Proteins contain
many negatively charged and basic
groups within their structure. Thus,

ION CONCENTRATION
a change in pH will alter the degree
ionization of a protein, which may in turn
affect its functioning. At more extreme
hydrogen ion concentrations a protein's
structure may be completely disrupted
(the protein is then said to be denatured).
Enzymes function optimally over a very
narrow range of hydrogen ion

ION CONCENTRATION
concentrations. For most enzymes this
optimum pH is close to the physiological
range for plasma (pH= 7.35 to 7.45, or [H+]=
35 to 45nmol/l). Figure 5 shows a typical
graph obtained when enzyme activity is
plotted against pH. Notice that the curve is a
narrow bell shape centred around
physiological pH.
Although most enzymes function optimally
around physiological pH it should be noted

ION CONCENTRATION
that a few enzymes function best at a much
higher hydrogen ion concentration (ie: at a
lower pH). The most notable of these
enzymes is pepsin, which works best in the
acid environment of the stomach - optimum
pH 1.5-3 or [H+]= 3-30 million nanomol/l.
 As enzymes have a huge number of
functions around the body, an abnormal pH
can result in disturbances in a wide range of

ION CONCENTRATION
body systems. Thus, disturbances in pH
may result in abnormal respiratory and
cardiac function, derangements in blood
clotting and drug metabolism, to name but
a few. From these few examples it is clear
that the anaesthetist should strive to
ensure that hydrogen ion concentration is
maintained within the normal range.

BUFFERS
Buffers are mixtures of chemicals which stabilize pH.
They can also be defined as mixtures of two
chemicals, weak acids, that resist pH changes:
 if the pH is too low one chemical will bind some of
the hydrogen ions and raise the pH:
 H +A -> HA
 Example: A = bicarbonate
 if the pH is too high the other chemical will donate
some hydrogen ions to lower the pH
 HA -> H + A
 Example: HA = carbonic acid

BUFFERS CONTD.
Blood pH Must be Kept Close to 7.4
 Hydrogen ion is extremely reactive and effects many
molecules which regulate physiological processes
 Blood pH is set at a slightly alkaline level of 7.4 (pH 7.0 is
neutral)
 A change of pH of 0.2 units in either direction is
considered serious
 Blood pHs below 6.9 or above7.9 are usually fatal if they
last for more than a short time

BUFFERS IN THE BODY
 The body has a very large buffer capacity.
 This can be illustrated by considering an old
experiment (see below) where dilute hydrochloric
acid was infused into a dog.
 Swan & Pitts Experiment
 In this experiment, dogs received an infusion of 14
mmols H+ per litre of body water. This caused a drop
in pH from 7.44 ([H+] = 36 nmoles/l) to a pH of 7.14
([H+] = 72 nmoles/l) That is, a rise in [H+] of only 36
nmoles/l.

BUFFERS IN THE BODY
CONTD.
 SO: If you just looked at the change in [H+] then you
would only notice an increase of 36 nmoles/l and you
would have to wonder what had happened to the other
13,999,964 nmoles/l that were infused.
Where did the missing H+ go?
 They were hidden on buffers and so these hydrogen
ions were hidden from view.

BUFFERS IN THE BODY
CONTD.
 Before we proceed, lets just make sure we appreciate
what this experiment reveals. The dogs were infused
with 14,000,000 nmoles/l of H+ but the plasma [H+]
only changed by a bit over 0.002%. By any analysis,
this is a system which powerfully resists change in
[H+]. Make no mistake: the body has:
 a HUGE buffering capacity, and
 this system is essentially IMMEDIATE in effect.
 For these 2 reasons, physicochemical buffering
provides a powerful first defence against acid-base
perturbations.

BUFFERS IN THE BODY
CONTD.
 Buffering hides from view the real change in H+ that
occurs.
 This huge buffer capacity has another not immediately obvious
implication for how we think about the severity of an acid-base
disorder. You would think that the magnitude of an acid-base
disturbance could be quantified merely by looking at the
change in [H+] - BUT this is not so.
 Because of the large buffering capacity, the actual change in
[H+] is so small it can be ignored in any quantitative
assessment, and instead, the magnitude of a disorder has to be
estimated indirectly from the decrease in the total concentration
of the anions involved in the buffering.

BUFFERS IN THE BODY
CONTD.
 The buffer anions, represented as A-, decrease because
they combine stoichiometrically with H+ to produce HA.
A decrease in A- by 1 mmol/l represents a 1,000,000 nano-
mol/l amount of H+ that is hidden from view and this is
several orders of magnitude higher than the visible few
nanomoles/l change in [H+] that is visible.) - As noted
above in the comments about the Swan & Pitts
experiment, 13,999,994 out of 14,000,000 nano-moles/l of
H+ were hidden on buffers and just to count the 36 that
were on view would give a false impression of the
magnitude of the disorder.

The Major Body Buffer Systems
ISF
 Bicarbonate: For metabolic acids
 Phosphate: Not important because concentration is too
low
 Protein: Not important because concentration is too low
Blood
 Bicarbonate: Important for metabolic acids
 Haemoglobin: Important for carbon dioxide
 Plasma protein: Minor buffer
 Phosphate: Concentration too low

The Major Body Buffer
Systems
ICF
 Proteins: Important buffer
 Phosphates: Important buffer
 Urine
 Phosphate: Responsible for most of 'Titratable Acidity'
 Ammonia: Important - formation of NH4
+
 Bone
 Ca carbonate: In prolonged metabolic acidosis

Systems
 Protein buffers in blood include haemoglobin (150g/l)
and plasma proteins (70g/l). Buffering is by the
imidazole group of the histidine residues which has a
pKa of about 6.8. This is suitable for effective
buffering at physiological pH. Haemoglobin is
quantitatively about 6 times more important then the
plasma proteins as it is present in about twice the
concentration and contains about three times the
number of histidine residues per molecule. For
example if blood pH changed from 7.5 to 6.5,

Systems
haemoglobin would buffer 27.5 mmol/l of H+ and
total plasma protein buffering would account for only
4.2 mmol/l of H+.
 Deoxyhaemoglobin is a more effective buffer than
oxyhaemoglobin and this change in buffer capacity
contributes about 30% of the Haldane effect. The
major factor accounting for the Haldane effect in CO2
transport is the much greater ability of
deoxyhaemoglobin to form carbamino compounds.

BUFFER SOLUTION
 This is an aqueous solution consisting of a mixture of a
weak acid and its conjugate base or a weak base and its
conjugate acid. It has the property that the pH of the
solution changes very little when a small amount of strong
acid or base is added to it. Buffer solutions are used as a
means of keeping pH at a nearly constant value in a wide
variety of chemical applications. Many life forms thrive
only in a relatively small pH range; an example of a buffer
solution is blood.

BUFFER SOLUTION CONTD.
Acidic buffer solutions
 An acidic buffer solution is simply one which has a pH
less than 7. Acidic buffer solutions are commonly made
from a weak acid and one of its salts - often a sodium salt.
 A common example would be a mixture of ethanoic acid
and sodium ethanoate in solution. In this case, if the
solution contained equal molar concentrations of both the
acid and the salt, it would have a pH of 4.76

BUFFER SOLUTION CONTD.
Alkaline buffer solutions
 An alkaline buffer solution has a pH greater than 7.
Alkaline buffer solutions are commonly made from a
weak base and one of its salts.
 A frequently used example is a mixture of ammonia
solution and ammonium chloride solution. If these were
mixed in equal molar proportions, the solution would have
a pH of 9.25. Again, it doesn't matter what concentrations
you choose as long as they are the same.

BUFFER BASE
Buffer base (BB) is the sum of all
buffering agents in the blood. It is
possible to state its value indirectly by
this equation: BB = (Na++K+)-(Cl-). It is
used in blood test to estimate acid-base
physiology in individuals.
Normal blood values are around
48 mmol/l

Normal Buffer Base
 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 Celsius) .
 The content of hemoglobin in blood increases this value.
The correlation is:
 NBB = 41.7+0.68×Hb, where Hb is in mmol/l. A normal
value of this is ~9.3 mmol/l, making the normal blood
value of NBB 48 mmol/l.

CHIEF BUFFER IN THE BLOOD
 The Chief Blood Buffer is a Mixture of Bicarbonate and
Carbon Dioxide
 All body fluids, inside or outside cells have buffers which
defend the body against pH changes
 The most important buffer in extracellular fluids, including
blood, is a mixture of carbon dioxide (CO2) and bicarbonate
anion (HCO3)
 CO2 acts as an acid (it forms carbonic acid when it dissolves in
water), donating hydrogen ions when they are needed
 HCO3 is a base, soaking up hydrogen ions when there are too
many of them

CHIEF BUFFER CONTD.
 The Bicarbonate Buffer System
 The major buffer system in the ECF is the CO2-
bicarbonate buffer system. This is responsible for
about 80% of extracellular buffering. It is the most
important ECF buffer for metabolic acids but it cannot
buffer respiratory acid-base disorders.
 The components are easily measured and are related to
each other by the Henderson-Hasselbalch equation.

CHIEF BUFFER CONTD.
 The HCO3/CO2 buffer system is extremely important
because it can be rapidly readjusted in alkalosis and
acidosis.
 There are also other buffers in blood, such as proteins and
phosphate.
 The ability to resist pH change is given by the buffer
capacity, which is a function of the concentration and
dissociation constant (pK) of the weak acid.
 If there is more than one buffer in the solution, the buffer
capacities add up.

THE BASE EXCESS
This is defined as the amount of acid (in
mmol) required to restore 1 litre of
blood to its normal pH, at a PCO2 of
5.3kPa (40mmHg). During the
calculation any change in pH due to the
PCO2 of the sample is eliminated,
therefore, the base excess reflects only the
metabolic component of any disturbance
of acid base balance.

THE BASE EXCESS CONTD.
If there is a metabolic alkalosis then acid
would have to be added to return the
blood pH to normal, therefore, the base
excess will be positive. However, if there
is a metabolic acidosis, acid would need
to be subtracted to return blood pH to
normal, therefore, the base excess is
negative.

STANDARD BICARBONATE
 This is defined as the calculated bicarbonate
concentration of the sample corrected to a PCO2 of
5.3kPa (40mmHg). Abnormal values for the standard
bicarbonate are only due the metabolic component of an
acid base disturbance. A raised standard bicarbonate
concentration indicates a metabolic alkalosis whilst a low
value indicates a metabolic acidosis.

ACTUAL BICARBONATE

This is the concentration of bicarbonate (hydrogen
carbonate) in the plasma of the sample. It is
calculated using the measured pH and pCO2 values.
The systematic symbol for arterial blood is cHCO3
-
(aP). The analyzer symbol may be cHCO3
- or
cHCO3
-(P).
 The actual bicarbonate is calculated by entering the
measured values of pH and pCO2 in the Henderson-
Hasselbalch equation.

ACTUAL BICARBONATE
CONTD.
An increased level of cHCO3
- may be due to a primary
metabolic alkalosis or a compensatory response to primary
respiratory acidosis. Decreased levels of cHCO3
- are seen
in metabolic acidosis and as a compensatory mechanism
to primary respiratory alkalosis.
 Reference ranges
cHCO3
-(aP) reference range (adult) [24]:
male: 24-31 mmol/L
female: 22-31 mmol/L

ACTUAL BICARBONATE
CONTD.
Clinical interpretation
The plasma bicarbonate level depends both on
“consumption” through titration of acids unrelated to
carbonic acid/pCO2 (the actual titration produces
water and carbon dioxide, of which the latter is
excreted by ventilation), on active regulation of renal
bicarbonate excretion, and on actual carbonic acid
concentration (i.e., pCO2 as per the Henderson-
Hasselbalch equation).
 A high bicarbonate concentration may thus be due to
either a true relative deficit of non-CO2-related acid

ACTUAL BICARBONATE
CONTD.
(e.g., in continuous gastric aspiration), to a
compensatory increase of renal proton excretion
(resulting in bicarbonate retention) in both instances
constituting an element of “metabolic alkalosis” - or
to a high pCO2 level.
 Conversely, a low bicarbonate concentration is
caused, either by titration of excess non-CO2-related
acid (e.g., in hyperlactatemia), by a compensatory
reduction in renal proton excretion, or by low pCO2.
 The evaluation of bicarbonate levels must therefore
always be done in conjunction with evaluation of the
pCO2 and the pH.

DETERMINANTS OF PH
 Water is the primary source of H+, and the
determinants of H+concentration are the determinants
of water dissociation. Fortunately, even for an
aqueous solution as complex as blood plasma, there
are but three independent variables that determine H+
concentration. These three variables are
mathematically independent determinants of the H+
concentration.

DETERMINANTS OF PH
CONTD.
 Thus, these variables are causally related to the H+
concentration rather than being merely correlated. Only by the
careful analysis of causal variables can mechanisms be
determined. For blood plasma, these three variables are :
 (i) pCO2,
 (ii) SID, and
 (iii) the total weak acid concentration (ATOT).

DETERMINANTS OF PH
CONTD.
Carbon dioxide
 CO2 is an independent determinant of pH and is
produced by cellular metabolism or by the titration of
HCO3 - by metabolic acids. Normally, alveolar
ventilation is adjusted to maintain the arterial pCO2
between 35 and 45 mmHg. When alveolar ventilation is
increased or decreased out of proportion to pCO2
production, a respiratory acid–base disorder exists. CO2
production by the body (at 220 ml/min) is equal to
15000 mmol/day of carbonic acid. This compares with
less than 500 mmol/day for all nonrespiratory acids.

DETERMINANTS OF PH
CONTD.
 The respiratory center, in response to signals from
pCO2, pH, and partial oxygen tension, as well as
some from exercise, anxiety, wakefulness, and
others, controls alveolar ventilation. A precise
match of alveolar ventilation to metabolic CO2
production attains the normal arterial pCO2 of 40
mmHg. Arterial pCO2 is adjusted by the
respiratory center in response to altered arterial pH
produced by metabolic acidosis or alkalosis in
predictable ways.

DETERMINANTS OF PH
CONTD.
 When CO2 elimination is inadequate relative to the
rate of tissue production, pCO2 will increase to a new
steadystate that is determined by the new relationship
between alveolar ventilation and CO2 production.
Acutely, this increase in pCO2 will increase both the
H+ and the HCO3
- concentrations according to the
Henderson-Hasselbach equation. Thus, this change in
HCO3
- concentration is mediated by chemical
equilibrium, and not by any systemic adaptation.

DETERMINANTS OF PH
CONTD.
 Similarly, this increased HCO3
- concentration
does not 'buffer' H+ concentration. There is no
change in the SBE. Tissue acidosis always
occurs in respiratory acidosis, because CO2
diffuses into the tissues. If the pCO2 remains
increased the body will attempt to compensate
by altering another independent determinant of
pH, namely the SID.

DETERMINANTS OF PH
CONTD.
Electrolytes (strong ions)
 Blood plasma contains numerous ions. These ions can be
classified both by charge, positive 'cations' and negative
'anions', as well as by their tendency to dissociate in aqueous
solutions. Some ions are completely dissociated in water, for
example, Na+, K+, Ca2+, Mg2+, and Cl-. These ions are called
'strong ions' to distinguish them from 'weak ions' (eg albumin,
phosphate and HCO3
-), which can exist both as charged
(dissociated) and uncharged forms. Certain ions such as lactate
are so nearly completely dissociated that they may be
considered strong ions under physiologic conditions.

DETERMINANTS OF PH
CONTD.
 In a neutral salt solution containing only water and NaCl, the
sum of strong cations (Na+) minus the sum of strong anions
(Cl-) is zero (ie Na+ = Cl-). In blood plasma, however, strong
cations (mainly Na+) outnumber strong anions (mainly Cl-).
The difference between the sum of all strong cations and all
strong anions is known as the SID. SID has a powerful
electrochemical effect on water dissociation, and hence on H+
concentration. As SID becomes more positive, H+, a 'weak'
cation, decreases (and pH increases) in order to maintain
electrical neutrality.
 In healthy humans, the plasma SID is between 40 and 42
mmol/l,

DETERMINANTS OF PH
CONTD.
Weak acids
 The third and final determinant of H+ concentration is ATOT.
The weak acids are mostly proteins (predominantlyalbumin)
and phosphates, and they contribute the remaining charges to
satisfy the principle of electroneutrality, such that SID–(CO2
+A-)=0. However, A- is not an independent variable because it
changes with alterations in SID and pCO2. Rather, ATOT (AH +
A-) is the independent variable, because its value is not
determined by any other. The identification of ATOT as the third
independent acid–base variable has lead some authors to
suggest that a third 'kind' of acid–base disorder exists.

DETERMINANTS OF PH
CONTD.
 Thus, along with respiratory and metabolic acidosis and
alkalosis, we would also have acidosis and alkalosis due
to abnormalities in ATOT. However, mathematical, and
therefore chemical independence does not necessarily
imply physiologic independence. Although the loss of
weak acid (ATOT) from the plasma space is an alkalinizing
process,there is no evidence that the body regulates ATOT
to maintain acid–base balance. Furthermore, there is no
evidence that we as clinicians should treat
hypoalbuminemia as an acid–base disorder.

DETERMINANTS OF PH
CONTD.
 Critically ill patients frequently have hypoalbuminemia
andas such their ATOT is reduced. These patients are not
often alkalemic and their SID is also reduced, however.
When these patients have a normal pH and a normal SBE
and HCO3
- concentration, it would seem most appropriate
to consider this to be physiologic compensation for a
decreased ATOT, rather than classifying this condition as a
complex acid–base disorder with a mixed metabolic
acidosis/hypoalbuminemic alkalosis. Thus, it seems far
more likely that this 'disorder' is in fact the normal
physiologic response to a decreased ATOT. Furthermore,
because changes in ATOT generally occur slowly,

DETERMINANTS OF PH
CONTD.
 the development of alkalaemia would require the kidney
to continue to excrete Cl- despite an evolving alkalosis. I
would consider such a scenario to be renal-mediated
hypochloremic metabolic alkalosis, the treatment for
which would include fluids and/or chloride, depending on
the clinical conditions. Stewart's designation of a 'normal'
SID of approximately 40 mmol/l was based on a 'normal'
CO2 and ATOT. The 'normal' SID for a patient with an
albumin of 2g/dl would be much lower (eg approximately
32 mmol/l).

CARBON DIOXIDE CARRIAGE
Carbon dioxide is carried in 3 forms which are:
(i) Dissolved CO2
(ii) Bicarbonate
(iii) Carbamino compounds
1. Dissolved CO2
 Follows Henry's Law
=> amount dissolved is proportional to partial pressure
 It is 20-25 times more soluble than O2
 At 37degrees, solubility = 0.0308 mmol/L/mmHg
 (?Roughly, 0.0835mL of CO2 is dissolved in 100mL of blood
per mmHg of PCO2)

CARBON DIOXIDE CARRIAGE CONTD.
2. Bicarbonate
 CO2 + H2O <-CA-> H2CO3 <-> H + HCO3
 The second step occurs rapidly (pKa'=6.1)
 Carbonic acid (H2CO3) forms a very small
percentage (<1%)
 The first step occurs very slowly in plasma,
BUT occurs rapidly in RBC due to presence of
carbonic anhydrase (CA) (zinc containing
enzyme)

Chloride shift
 In RBC, as H2CO3 is formed (with aid of CA), it breaks
down to H + HCO3
 => HCO3 diffuse out easily
 => H+ doesn't really diffuse out because RBC membrane
is relatively impermeable to cations
 => As [H+] builds up, chloride diffuse into RBC to
maintain electrical neurality (in accordance to Gibbs-
Donnan equilibrium)
aka "chloride shift"

 Exchange of HCO3 and Cl (Hamburger effect) occurs at the
transporter "capnophorin", which is a band 3 protein
 H+ are buffered by histidine
 Osmolarity increase, and water enters RBC as a result
 => slight increase in RBC volume as CO2 is taken up.
 3. Carbamino compound
 Formed by CO2 binding to terminal amino groups in blood
protein.
 To much lesser extent amino groups in the side chains of
arginine and lysine.
 Globin in haemoglobin is the most significant.

 => as with bicarbonate, RBC is the main stage.
 Hb-NH2 + CO2 <-> Hb-NH-COOH <-> Hb-NH-
COO + H
=> at normal pH, almost completely dissociates
 Hb-NH-COOH is called "carbamino haemoglobin"
 Responsible for 30% of the CO2 eliminated in lung
 Not really affected by PCO2

 Imidazole group of histidine
 Imidazole group of the amino acid histidine is the only real
effective buffer in the normal range of pH (pKa 6.8)
 The buffering power of plasma protein is more or less
proportional to their histidine content.
 Haemoglobin has 38 histidine residues, far more than plasma
proteins
 Buffering
 Both carbonic acid and carbamino haemogoblin almost
completely dissociate
 => H+ are produced
 => buffered by histidine

 As haemoglobin becomes reduced (i.e. deoxyHb)
 => it becomes less acid / better base
 => it becomes a better buffer
 Haldane effect
 Haldane effect refers to the increased ability of blood to carry CO2
when haemoglobin gives up oxygen.
 Haldane effect is due to:
 DeoxyHb is 3.5 times more effective than oxyHb in forming
carbamino compounds.
=> accounts for 70% of the Haldane effect
 DeoxyHb is a better buffer than oxyHb, thus improving CO2
carriage as bicarbonates
=> accounts for 30% of the Haldane effect

 NB: Bohr effect refers to increased unloading of O2 from
Hb when carbon dioxide is loaded
(causing pH drop, and thus right shift in oxygen
dissociation curve)
 CO2 dissociation curve
 more linear than O2 dissociation curve
 more steep than O2 dissociation curve
=> smaller change in PCO2 when content changes
 as PaO2 drops, affinity for CO2 increases
=> CO2 dissociation curve moves to the LEFT (Haldane
effect)

 Carriage in arterial blood vs venous blood
 Arterial blood contains 48 mLs of CO2 per 100mL
blood.
 Venous blood contains 52 mLs of CO2 per 100mL
blood.
 Total carriage
 5% - Dissolved CO2
 90% - Bicarbonate
 5% - Carbamino compounds

 % Contribution to A-V difference
 (i.e. % of CO2 eliminated at lung)
 10% - Dissolved CO2
 60% - Bicarbonate
 30% - Carbamino compounds
 Due to storage of CO2 as bicarbonate, changes in
PaCO2 due to changes in ventilation takes a little
longer to equilibrate than PaO2.

 Factors affecting PCO2 in steady state
 Alveolar CO2 conc. x alveolar ventilation = CO2
output
 Alveolar ventilation
 Concentration effect
=> when inert gases are taken up rapidly, PACO2 is
increased due to concentration effect
 CO2 output
 Inspired FICO2

RENAL ACID EXCRETION
 The net quantity of H+ ions excreted in the urine
is equal to the amount of H+ excreted as titratable
acidity and NH4+ minus any H+ added to the
body because of urinary HCO3- loss.
 Net acid excretion(NAE) = titratable acidity +
NH4
+ - urinary HCO3
-
 Note that normally there is no urinary HCO3- and
therefore:

 Net acid excretion(NAE) = titratable acidity
+ NH4
+
Titratable acidity is dependent on the dietary
intake of phosphate and cannot be regulated to
increase acid excretion. The kidney 's main
response to an increased acid load is to
increase ammonium production and excretion.

 A very important feature of titrable acidity and
ammonium excretion is the regeneration of
bicarbonate ions. The kidney must reabsorb all
filtered HCO3- in order to maintain acid base
balance. Hydrogen ion secretion in the
collecting tubule is very important in
maximally acidifying the urine.

 In states of acidosis, maximal acidification of
the urine in the collecting tubule must occur for
adequate ammonium excretion. Ammonium
excretion is increased by increasing ammonium
production and increased hydrogen ion
secretion in the collecting duct.

 Aldosterone stimulates secretion of hydrogen
ion in the collecting duct. Although the
extracellular pH is the primary physiologic
regulator of net acid excretion, in
pathophysiologic states, the effective
circulating volume, Aldosterone, and the
plasma K+ concentration all can affect acid
excretion, independent of the systemic pH.

PRINCIPLE OF BUFFERING
 Buffer solutions achieve their resistance to pH
change because of the presence of an
equilibrium between the acid HA and its
conjugate base A-.
 HA H+ + A- When some strong acid is added to
an equilibrium mixture of the weak acid and its
conjugate base, the equilibrium is shifted to the
left, in accordance with Le Chatelier's
principle.

 Because of this, the hydrogen ion concentration
increases by less than the amount expected for
the quantity of strong acid added. Similarly, if
strong alkali is added to the mixture the
hydrogen ion concentration decreases by less
than the amount expected for the quantity of
alkali added.

 The effect is illustrated by the simulated titration of a
weak acid with pKa = 4.7. The relative concentration
of undissociated acid is shown in blue and of its
conjugate base in red. The pH changes relatively
slowly in the buffer region, pH = pKa ± 1, centered at
pH = 4.7 where [HA] = [A-], but once the acid is more
than 95% deprotonated the pH rises much more
rapidly. Blood pH is determined by a balance between
bicarbonate and CO2 ad shown by these diagrams:

 Too Much CO2 or Too Little HCO3 Will Cause
Acidosis
 The balance will swing toward a low pH, producing
acidosis, if CO2 is raised or HCO3 lowered
 CO2 can be raised by hypoventilation (pneumonia,
emphysema)
 Metabolic conditions such as ketoacidosis caused by
excess fat metabolism (diabetes mellitus) will lower
bicarbonate

COMPENSATORY
MECHANISMS
 Compensation for acidosis (rebalances the pH to 7.4):
Add HCO3
 Remove CO2: occurs first because lungs work faster
than kidneys

COMPENSATORY
MECHANISMS
 Too Much HCO3 or Too Little CO2 Will Cause
Alkalosis
 The balance will swing the other way, producing alkalosis,
if CO2 is lowered or HCO3 raised
 CO2 can be lowered by hyperventilation
 Vomiting removes stomach acid and raises bicarbonate
 Alkalosis is less common than acidosis

COMPENSATORY
MECHANISMS
 Compensation for alkalosis (rebalances the pH
to 7.4): Remove HCO3
 Add CO2: occurs first because lungs work
faster than kidneys

COMPENSATORY
MECHANISMS
 Blood pH is Chiefly Regulated by the Lungs and Kidneys
 Normal metabolism produces large amounts of CO2
continuously (about 14 moles/day)
 If this CO2 were not removed we would rapidly develop fatal
acidosis
 Almost all of the CO2 is removed, as a gas, from the lungs
 If blood pH is low respiration is stimulated so that more CO2 is
removed, raising the pH to the normal level
 Bicarbonate is adjusted in the kidney
 Most filtered bicarbonate is reabsorbed in the proximal tubule

COMPENSATORY
MECHANISMS
The kidneys also dispose of non-volatile acids
produced in metabolism Additional processes
are used by the kidney to regulate pH:
Secretion of H ions
 Occurs in the proximal tubule and distal
tubules
 Secretion into blood lowers the pH
 Secretion into the tubule raises the pH

COMPENSATORY
MECHANISMS
 Production of new bicarbonate in distal tubule:
 The distal tubule has fine control over
bicarbonate
 Secreted into the blood raises the pH
 Secretion into tubule lowers the pH indirectly
 Production of ammonia (NH3) in proximal tubule
cells during acidosis
 Helps to remove excess H by forming
ammonium ion (NH4+) in the tubule

ACID BASE DISTURBANCES
 Acid–base imbalance is an abnormality of the
human body's normal balance of acids and
bases that causes the plasma pH to deviate out
of the normal range (7.35 to 7.45). An acid
base disorder is a change in the normal value
of extracellular pH that may result when renal
or respiratory function is abnormal or when an
acid or base load overwhelms excretory
capacity.

ACID BASE DISTURBANCES
CONTD.
In the fetus, the normal range differs based on
which umbilical vessel is sampled (umbilical
vein pH is normally 7.25 to 7.45; umbilical
artery pH is normally 7.18 to 7.38).[1] It can
exist in varying levels of severity, some life-
threatening. It can also be defined as a change
in the normal value of extracellular pH that
may result when renal or respiratory function
is abnormal or when an acid or base load
overwhelms excretory capacity.

ACID BASE DISTURBANCES CONTD.
Metabolic Acidosis
Respiratory Acidosis
Metabolic Alkalosis
Respiratory Alkalosis
Mixed acid-base disorders

ACID BASE DISTURBANCES CONTD.
 The presence of only one of the above derangements is
called a simple acid–base disorder. In a mixed disorder
more than one is occurring at the same time. Mixed
disorders may feature an acidosis and alkosis at the
same time that partially counteract each other, or there
can be two different conditions effecting the pH in the
same direction. The phrase "mixed acidosis", for
example, refers to metabolic acidosis in conjunction
with respiratory acidosis. Any combination is possible,
except concurrent respiratory acidosis and respiratory
alkalosis, since a person cannot breathe too fast and
too slow at the same time.

RESPIRATORY ALKALOSIS
 Blood PCO2 less than 35 and pH greater than 7.45
Etiology:
1. low carbon dioxide due to increased losses via the lungs
2. Causes may be salicylate poisoning, hypoxia, gram-
negative sepsis, liver failure, hyperventilation, mechanical
ventilation, inflammation or tumor of the thorax, primary
CNS disorder (tumor, infection, trauma or CVA), but also
anxiety (hyperventilation) or pregnancy.

RESPIRATORY ALKALOSIS
CONTD.
 Manifestations:
1. CNS disorders may be expressed as anxiety or
alterations in the state of consciousness.
2. Acute alkemia may cause symptoms of tetany and
needs to be differentiated from hypocalcemia.
Principles of management:
1. Correct the underlying disorder.
2 If pH is greater than 7.6, controlled ventilation may
be called for.

METABOLIC ALKALOSIS
 Blood pH greater than 7.45 and bicarbonate greater than 28
Etiology:
1. Increased production of bicarbonate is seen with vomiting
(due to loss of acid) and with the rapid correction of
hypercapnia from respiratory acidosis.
2 Excess bicarbonate from an exogenous source is also a
problem.
3 Primary or secondary hyperaldosteronism stimulate hydrogen
ion secretion in the distal tubule, raising bicarbonate levels.
Diuretics have the same effect.

METABOLIC ALKALOSIS
CONTD.
 Manifestations:
of the underlying problem. May be symptoms of
tetany
Diagnosis:
1. Blood analysis shows increased bicarbonate and
decreased chloride, often with hypokalemia.
2 Urine chloride is increased in normovolemic states
and very low with hypovolemia.
3 Arterial blood gases show increased bicarbonate and
increased pCO2 and may also show hypoxia.

METABOLIC ALKALOSIS
CONTD.
 Principles of management:
1 Correct underlying disease state.
2 In cases where there is also hypokalemia, the best
treatment is potassium chloride.
3 Metabolic alkalosis that occurs after hypercapnia
correction may require acetazolamide as well.

RESPIRATORY ACIDOSIS
 Definition:
Blood pCO2 greater than 40 mm Hg and pH less than
7.35
Etiology:
1. Diminished capacity of the lungs to clear CO2
2. May be due to primary lung disease, primary CNS
dysfunction, neuromuscular disease or drugs causing
hypoventilation.

 Manifestations:
1. Respiratory acidosis stimulates blood flow to
the brain that may cause cerebrospinal fluid
pressure to increase above the normal limit.
This may lead to CNS depression.
2 Acidemia decreases cardiac output and
increases pulmonary hypertension, thereby
decreasing blood flow to tissues.

 Principles of management:
1. The underlying disorder needs to be corrected, if
possible.
2 If blood pCO2 is more than 60 mm Hg, assisted
respiration may be necessary.

METABOLIC ACIDOSIS
This is defined as pH less than 7.35 and bicarbonate less than 21
Etiology:
1. Is due to either loss of bicarbonate or accumulation of another acid
(such as lactic acid)
2. Can be divided into metabolic acidosis with and without an anion
gap. The anion gap refers to the anions actually present in the serum
but usually not measured – mainly albumin, phosphates, sulfates and
organic acids. It is calculated by the following formula (cations
minus measured anions):
Anion gap = [Na+] – ([Cl-] + [HCO3-])
The normal value is 6-14 mEq/ml.

METABOLIC ACIDOSIS
CONTD.
 (A) Metabolic acidosis with a normal anion gap
(hyperchloremic metabolic acidosis) is due to loss of
bicarbonate via the gastrointestinal tract (diarrhea,
pancreatic fistula, ureterosigmoidostomy) or via the
kidney. Renal loss may be due to proximal tubular
acidosis (nephrotic syndrome, cystinosis, multiple
myeloma, Wilson’s disease, heavy metal poisoning),
distal tubular acidosis (SLE, Sjogren’s syndrome,

METABOLIC ACIDOSIS
CONTD.
obstructive uropathy, amphotericin B toxicity)
or hyperkalemic renal tubular acidosis.
 Carbonic anhydrase inhibitors such as
acetazolamide or mafenide inhibit bicarbonate
reabsorption in the proximal tubule. Moderate
renal failure, with GFR of 15-40 ml/min,
shows a decline in ammonium excretion due to
decreased renal mass.

METABOLIC ACIDOSIS
CONTD.
 (B) Metabolic acidosis with positive anion gap
is seen with ketoacidosis (diabetes, starvation,
alcohol abuse), lactic acidosis (shock, sepsis),
drug intoxications (salicylate, methanol,
ethylene glycol) and renal failure.
Winter’s formula – With pure metabolic
acidosis, pCO2 is 1.5 times the bicarbonate
concentration plus 6-10 mm Hg.

METABOLIC ACIDOSIS
CONTD.
 An actual pCO2 less than that predicted by this
formula suggests primary respiratory alkalosis is also
involved; if actual pCO2 is higher than predicted, the
complication is a disorder of pulmonary function and
CO2 retention.
Manifestations:
1. Decreased cardiac output may be seen if pH is less
than 7.2.
2 Kussmaul breathing (deep and rhythmic) appears as
the lungs increase ventilation rate to compensate.

METABOLIC ACIDOSIS
CONTD.
Principles of management:
Bicarbonate treatment is necessary to raise pH
to at least 7.2. The amount of bicarbonate to
be administered is based on bicarbonate
occupying half of the body weight, according
to the formula:
(Desired bicarbonate level – actual
bicarbonate level) x 0.5 x body weight

Mixed acidosis
 PaCO2 increased and HCO3
- decreased
 This is very dangerous and may occur in severe
diseases such as septic shock, multiple organ
dysfunction, cardiac arrest.

Compensatory Mechanisms
 The body's acid–base balance is tightly regulated. Several
buffering agents exist which reversibly bind hydrogen
ions and impede any change in pH. Extracellular buffers
include bicarbonate and ammonia, while proteins and
phosphate act as intracellular buffers. The bicarbonate
buffering system is especially key, as carbon dioxide
(CO2) can be shifted through carbonic acid (H2CO3) to
hydrogen ions and bicarbonate (HCO3
- ) as shown below.

Compensatory Mechanisms contd.
Acid–base imbalances that overcome the
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 above reaction
according to Le Chatelier's principle, which in
turn alters the pH. For instance, if the blood
pH drops too low (acidemia), the

Compensatory Mechanisms Contd.
body will compensate by increasing breathing,
expelling CO2, and shifting the following
reaction to the right such that less hydrogen
ions are free - thus the pH will rise back to
normal. For alkalemia, the opposite occurs.
 The kidneys are slower to compensate, but
renal physiology has several powerful
mechanisms to control pH by the excretion of
excess acid or base. In responses to

Compensatory Mechanisms contd.
acidosis, tubular cells reabsorb more bicarbonate
from the tubular fluid, collecting duct cells secrete
more hydrogen and generate more bicarbonate,
and ammoniagenesis leads to increased formation
of the NH3 buffer. In responses to alkalosis, the
kidney may excrete more bicarbonate by
decreasing hydrogen ion secretion from the
tubular epithelial cells, and lowering rates of
glutamine metabolism and ammonia excretion.

EFFECTS OF ACID-BASE
DISTURBANCES
 Acid–base homeostasis exerts a major influence on
protein function, thereby critically affecting tissue and
organ performance. Deviations of systemic acidity in
either direction can have adverse consequences and, when
severe, can be life-threatening. Yet it is the nature of the
condition responsible for severe acidemia or alkalemia
that largely determines the patient's status and prognosis.
Whereas a blood pH of 7.10 can be of little consequence
when caused by a transient or easily

DISTURBANCES
reversible condition, such as an isolated
seizure, it forecasts an ominous outcome if it is
the result of methanol intoxication. Similarly, a
blood pH of 7.60 seldom has serious
consequences when caused by the anxiety–
hyperventilation syndrome, but it imparts a
major risk to a patient with cardiomyopathy
treated with digitalis and diuretics

DISTURBANCES
. Consequently, the management of serious
acid–base disorders always demands precise
diagnosis and treatment of the underlying
disease, and in certain circumstances, it
requires steps to combat the deviation in
systemic acidity itself. In this article, we
address general concepts and some specific
aspects of the management of life-threatening
acid–base disorders.

DISTURBANCES
 The major adverse consequences of severe acidemia
occurs at blood pH of <7.20 Consequences of Severe
Acidemia can occur independently of whether the
acidemia is of metabolic, respiratory, or mixed origin.
The effects on the cardiovascular system are particularly
pernicious and can include decreased cardiac output,
decreased arterial blood pressure, decreased hepatic and
renal blood flow, and centralization of blood volume.1,2
Reentrant arrhythmias and a reduction in the threshold
for ventricular fibrillation can occur, while the

DISTURBANCES
defibrillation threshold remains unaltered.
Acidemia triggers a sympathetic discharge but
also progressively attenuates the effects of
catecholamines on the heart and the vasculature;
thus, at pH values below 7.20, the direct effects of
acidemia become dominant.
 Although metabolic demands may be augmented
by the associated sympathetic surge, acidemia
decreases the uptake of glucose in the tissues by

DISTURBANCES
inducing insulin resistance and inhibits
anaerobic glycolysis by depressing 6-
phosphofructokinase activity. This effect can
have grave consequences during hypoxia, since
glycolysis becomes the main source of energy
for the organism. The uptake of lactate by the
liver is curtailed, and the liver can be converted
from the premier consumer of lactate to a net
producer.1 Acidemia causes potassium to

DISTURBANCES
leave the cells, resulting in hyperkalemia, an
effect that is more prominent in nonorganic
acidoses than in organic and respiratory
acidoses. Increased net protein breakdown and
development of a catabolic state also occur in
patients with acidosis. Brain metabolism and
the regulation of its volume are impaired by
severe acidemia, resulting in progressive
obtundation and coma.

ANION GAP
The Anion Gap is defined as the diffe-
rence between the sum of the major
anions and the major cations:
Gap = Na+ + K+ - Cl- - HCO3
-
It can also be defined as the concentration
difference between the cations, sodium
and potassium,

ANION GAP CONTD.
and the measured anions, chloride and
bicarbonate.
Anion Gap(K+) = cNa+ + cK+ - cCl– -
cHCO-
3.
The systematic symbol is Anion Gap(K+).
The analyzer symbol may be Anion
Gap(K+).

ANION GAP CONTD.
 Of what use is Anion Gap(K+)?
Anion Gap(K+) is a reflection of the unmeasured
anions in the plasma, e.g., proteins, organic acids,
sulfates, and phosphates (although changes in plasma
calcium and magnesium also affect the Anion
Gap(K+)).
Anion Gap(K+) may be an aid in the differential
diagnosis of metabolic acidosis.

ANION GAP CONTD.
 Metabolic acidosis can be classified in two groups:
 Those with an increase in Anion Gap(K+), thus implying
the presence of increased amounts of organic acid.
 Those with normal Anion Gap(K+), due to loss of
bicarbonate.
 Reference ranges
Anion Gap(K+) reference range (adult): 10-20 mmol/L

ANION GAP CONTD.
 Clinical interpretation
A. Decreased Anion Gap(K+) can be caused by:
 Decrease in plasma proteins
 Hyponatraemia
 Increase in unmeasured cations
B. Increased Anion Gap(K+) can be caused by:
 Ketoacidosis
 Lactoacidosis
 Renal failure
 Intoxication with: salicylate, methanol, and ethylene glycol

ANION GAP CONTD.
C. Metabolic acidosis with a normal Anion
Gap(K+):
 Diarrhea
 Uremic acidosis of recent onset
 Renal tubular acidosis

THE END
THANKS FOR YOUR RAPT ATTENTION!!!

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