Transcript acid-base balance

49 – Perioperative Acid-Base Balance
Patrick J. Neligan,
Clifford S. Deutschman
Life is a struggle,
not against sin, not against Money Power . . but
against hydrogen ions.
--H.L. Mencken
[H] = 24 × PCO2/HCO3
In the arterial blood plasma of a normal person at rest, the hydrogen ion
concentration usually lies in the range 35 to 45 nanomoles/litre, with an
average of 40.
This chapter first explores:
the physical chemistry of water,
then examines
the evolution of acid-base medicine,
and finally applies:
clinical approaches to acid-base
conundrums
+
H
OH
Acid-base abnormalities should be seen as
resulting from other biochemical
changes in the extracellular environment
strong ions (Na , Cl , K , SO4 , Mg , Ca )
weak acids(albumin , phosphate )
+
-
+
2-
2+
2+
carbon dioxide
To maintain electrical neutrality
PHYSICAL CHEMISTRY OF WATER
H2O
105o
+
H
O
O
+
OH
fundamental
to the
existence of
life
PHYSICAL CHEMISTRY OF WATER
1.0 × 10−7 mmol/L
H2O
+
H
+
OH
25o C
Keq H2 O = [H+ ][OH− ]
Keq H2 O = Keq (55.5) = Kw = [H+ ][OH− ]
A solution is considered acidic if :
([H+ ] > 1.0 × 10−7 mmol/L, [OH− ] < 1.0 × 10−7 mmol/L).
A solution is considered alkaline if
([H+ ] < 1.0 × 10−7 mmol/L, [OH− ] > 1.0 × 10−7 mmol/L).
ACIDS AND BASES
• Svante Arrhenius in 1903 established the foundations
of acid-base chemistry.
• In an aqueous solution,
an Arrhenius acid is any substance that delivers a
hydrogen ion into the solution. (HCl)
A base is any substance that delivers a hydroxyl ion
into the solution. (NaOH)
• In 1909, L.J. Henderson coined the term
acid-base balance.
• Hasselbalch (1916)
H2O + CO2 → H2CO3 → [H+] + [HCO3-]
• pH = pKa + log [HCO3 − ] / [H2 CO3 ]
• pH = 6.1 + log [HCO3− ] / PCO2 × 0.03
• The degree of dissociation of substances in water determines
whether they are strong acids or strong bases.
• Lactic acid, pKa of 3.4, is a strong acid.
• Carbonic acid, pKa of 6.4, is a weak acid.
• Similarly, ions such as sodium, potassium, and chloride, which
do not easily bind other molecules, are considered
strong ions; they exist free in solution.
• Strong cations (Na+ , K+ , Ca2+ , Mg2+ ) act as Arrhenius bases
• Strong anions (Cl- , LA- [lactate], ketones, sulfate, formate) act
as Arrhenius acids.
• One problem with the Arrhenius theory:
ammonia (NH3 ), sodium carbonate (Na2 CO3 ), and sodium
bicarbonate (NaHCO3 )
• In 1923, Brønsted and Lowry
They defined acids as proton donors and bases as proton
acceptors.
• NH3 + H2 O ⇌ NH4 + + OH−
In this situation, water is the proton donor, the Brønsted-Lowry
acid, and ammonia the proteon acceptor, the Brønsted-Lowry
base.
• HCl + H2 O → H3 O+ + Cl−
In the previous reaction, hydrogen chloride acts as a
Brønsted-Lowry acid and water as a Brønsted-Lowry base.
• CO2 + H2 O ⇌ H2 CO3 ⇌ H+ + HCO3 −
In this reaction, carbon dioxide is hydrated to carbonic acid, a
Brønsted-Lowry acid, which subsequently dissociates to
hydrogen (H+ ) and bicarbonate (HCO3 - ) ions.
The Lewis Definitions of Acids and Bases
-------------------------------------------------------------------------------*In 1923 G. N. Lewis suggested another way of looking at the reaction between
H+ and OH- ions. In the Brnsted model, the OH- ion is the active species in this
reaction it accepts an H+ ion to form a covalent bond. In the Lewis model, the H+
ion is the active speciesit accepts a pair of electrons from the OH- ion to form a
covalent bond
*In the Lewis theory of acid-base reactions, bases donate pairs of electrons
and acids accept pairs of electrons. A Lewis acid is therefore any
substance, such as the H+ ion, that can accept a pair of nonbonding
electrons. In other words, a Lewis acid is an electron-pair acceptor. A Lewis
base is any substance, such as the OH- ion, that can donate a pair of
nonbonding electrons. A Lewis base is therefore an electron-pair donor.
STEWART APPROACH TO
ACID-BASE BALANCE
• What Determines the Acidity or Alkalinity of a
solution?
• The molar concentration of hydrogen and hydroxide must
be used to reflect the relative acidity and alkalinity of a
solution.
• The pH scale, developed by Sorenson in the 1920s.
• Neutral pH for pure water is 7.0 (1.0 × 10−7 mmol/L).
• Physiological pH for the ECF is 7.4, which is alkaline.
• The pH of the intracellular space is 6.8 to 7.0
STEWART APPROACH TO
ACID-BASE BALANCE
• 1. Electrical neutrality. In aqueous solutions in any
compartment, the sum of all of the positive charged ions must
equal the sum of all of the negative charged ions.
• 2. Dissociation equilibria. The dissociation equilibria of all
incompletely dissociated substances, as derived from the law of
mass action, must be satisified at all times.
• 3. Mass conservation. The amount of a substance remains
constant unless it is added, removed, generated, or destroyed.
• The total concentration of an incompletely dissociated
substance ATOT is the sum of concentrations of its dissociated
and undissociated forms.
Strong Ions
• The most abundant strong ions in the
extracellular space are Na+ and Cl- .
• Other important strong ions include
K+, SO4 2- , Mg2+, and Ca2+.
• If NaOH and HCl, added to solution
([Na+ ] − [Cl− ]) + ([H+ ] − [OH− ]) = 0
• In this system, ([Na+ − [Cl−]) must determine
[H+] and [OH−].
•In any solution, the sum total of the charges imparted by strong
cations minus the charges from strong anions represents the SID.
•The SID independently influences hydrogen ion concentration.
•In human ECF, the SID is positive.
•SID is an independent variable and [H+] and [OH-] are dependent,
meaning that the addition of hydrogen ions alone (without
strong corresponding anions) cannot influence the pH of the
solution
A
B
c
???
D
In the arterial blood plasma of a normal person at rest,
the hydrogen ion concentration usually lies in the range
35 to 45 nanomoles/litre, with an average of 40..
Weak Acid Buffer Solutions
• These are partially dissociated compounds
whose degree of dissociation is determined by the
prevailing temperature and pH.
• The predominant molecules in this group are
albumin and phosphate.
• Stewart used the term ATOT to represent the total
concentration of weak ions that influenced acid-base
balance
Weak Acid Buffer Solutions
• KA [HA] = KA [H+] [A−]
• [HA] + [A− ] = [ATOT ] or [HA] [A− ] = [ATOT ]
• [H+]× [OH−] = Kw (water dissociation)
• [SID] + [H+] − [A−] − [OH−] = 0
(electrical neutrality)
• SID and ATOT are independent variables
• Kw and KA are constants
• [HA], [H+], [OH-], and [A-] are dependent variables.
Carbon Dioxide
• 1- carbon dioxide, denoted CO2 (d);
2- carbonic acid (H2 CO3)
3- bicarbonate ions (HCO3- )
4- carbonate ions (CO32- )
• [CO2 (d)] = [SCO2] × PCO2
• [CO2 (d)] × [OH− ] = K1 × [HCO3− ]
• [H+ ] × [HCO3− ] = Kc × PCO2
• [H+ ] × [CO32− ] = K3 × [HCO3− ]
Factors Independently Influencing Water Dissociation
• Water dissociation equilibrium:
[H+ ] × [OH- ] = KW
• Weak acid dissociation equilibrium:
[H+ ] × [A- ] = KA × [HA]
• Conservation of mass for weak acids:
[HA] + [A- ] = [ATOT ]
• Bicarbonate ion formation equilibrium:
[H+ ] × [HCO3 - ] = KC × PCO2
• Carbonate ion formation equilibrium:
[H+ ] × [CO3 2- ] = K3 × [HCO3 - ]
• Electrical neutrality:
[SID] + [H+ ] - [HCO3 - ] - [A- ] - [CO3 2- ] - [OH- ] = 0
• [H+ ]4 + ([SID] + KA ) × [H+ ]3 + (KA × ([SID] − [ATOT ]) −
Kw − Kc × PCO2 ) × [H+ ]2 − (KA × (Kw + Kc × PCO2 ) −
K3 × Kc × PCO2 ) × [H+ ] − KA × K3 × Kc × PCO2 = 0
• [H+ ] is a function of SID, ATOT , PCO2 ,
and several constants.
• All other variables, most notably [A- ]
[H+], [OH-], and [HCO3 -], are dependent
and cannot independently influence the
acid-base balance
ACID-BASE ABNORMALITIES
• Stewart approach : SID, ATOT, PCO2
Traditional approach
• Alterations in (PaCO2) tension :
respiratory acidosis or alkalosis
• Alterations in blood chemistry : ( HCO3- , BE )
metabolic acidosis, or alkalosis
although respiratory or metabolic abnormalities rarely occur independently
of one another
ACID-BASE ABNORMALITIES
•
Respiratory Acid-Base Abnormalities
Respiratory acidosis → acute rise in PaCO2
principally because of respiratory failure
• Clinically, : (signs of CO2 retention)
Cyanosis, vasodilatation, and narcosis.
• Respiratory alkalosis → acute decrease in PaCO2
(caused by hyperventilation.)
• Clinically :
Vasoconstriction: light-headedness, visual disturbances,
dizziness, and perhaps hypocalcemia
Respiratory acidosis
Causes a rapid increase in [H+].
Compensation for hypercarbia is slow
Increased urinary excretion of chloride
There is a concomitant increase in the serum
bicarbonate, reflecting a higher total CO2 load,
rather than compensation.
• The acuity of respiratory failure can be deduced by
looking at the relative ratio of CO2 to HCO3–
• Many investigators have suggested that respiratory
acidosis may not necessarily be harmful.
• There has been extensive clinical experience with
"permissive hypercapnia" for acute respiratory
failure, which appears to be well tolerated.
•
•
•
•
_
TABLE 41-1 -- Changes in PaCO2 and [HCO3 ] in response to acute
and chronic acid-base disturbances
Disturbances
[HCO3 - ] vs. PaCO2
Acute respiratory acidosis
ΔHCO3 - = 0.2 ΔPaCO2
Acute respiratory alkalosis
ΔHCO3 - = 0.2 ΔPaCO2
Chronic respiratory acidosis
ΔHCO3 - = 0.5 ΔPaCO2
Metabolic acidosis
ΔPaCO2 = 1.3 ΔHCO3 -
Metabolic alkalosis
ΔPaCO2 = 0.75 ΔHCO3 -
Δ, change in value; [HCO3-], concentration of bicarbonate ion; PaCO2,
partial pressure of arterial carbon dioxide.
ACID-BASE ABNORMALITIES
Metabolic Acid-Base Disturbances
Metabolic acid-base abnormalities → SID or ATOT, or both
•
• An increase in the SID causes alkalemia
• A decrease in the SID causes acidemia
(e.g., hyperchloremia, lacticemia, dilutional acidosis )
• Metabolic acidosis is of clinical significance for two
reasons:
1. pathologies arising from the acidosis itself
2. pathologies arising from the cause of the acidosis
(Increased ionized calcium → vasodilation, diminished
muscular performance (particularly myocardial), and
arrhythmias )
TABLE 41-2 -- Classification of primary acid-base abnormalities
Abnormalities
Respiratory
Acidosis
Increased PCO2
Alkalosis
Decreased PCO2
Metabolic
Abnormal SID
Caused by water excess or deficit
Water excess = dilutional
↓ SID + ↓ [Na+ ]
Water deficit = contraction
↑ SID ↑ [Na+ ]
Caused by electrolytes
Chloride excess
Chloride deficit
Chloride (measured)
↓ SID ↑ [Cl- ]
↑ SID + ↓ [Cl- ]
Other (unmeasured) anions, such as
lactate and keto acids
↓ SID ↑ [UMA- ]
Abnormal ATOT
Albumin [Alb]
↑ [Alb] (rare)
Phosphate [Pi]
↑ [Pi]
↓ [Alb]
[Alb], concentration of serum albumin; ATOT , to represent the total concentration of weak ions; [Cl- ],
concentration of chloride ions; [Na+ ], concentration of sodium ions; PCO2 , partial pressure of carbon
dioxide; [Pi], concentration of inorganic phosphate; SID, strong ion difference; [UMA - ], unmeasured
anions; ↑, increased; ↓, decreased.
REGULATION OF ACID-BASE BALANCE
• A buffer is a solution of two or more chemicals that
minimizes changes in pH in response to the addition
of an acid or base.
• Most buffers are weak acids. Ideally, a buffer has a
pKa that is equal to the pH, and an ideal body buffer
has a pKa between 6.8 and 7.2.
• The major source of acid in the body is CO2 , from
which is produced 12,500 mEq of H+ each day.
• The metabolic compensation for respiratory
acidosis is increased SID by removal of chloride.
• Volatile acid is principally buffered by hemoglobin.
• Chloride shift:
Erythrocyte Buffering System
And Chloride Shift
REGULATION OF ACID-BASE BALANCE
• Metabolic acid is buffered principally by increased
alveolar ventilation, producing respiratory
alkalosis and extracellular weak acids.
• These weak acids include :
plasma proteins, phosphate, and bicarbonate.
• The bicarbonate buffering system (92% of plasma
buffering and 13% overall) is probably the most
important extracellular buffer.
REGULATION OF ACID-BASE BALANCE
• The major effect of the kidney on acid-base balance is related to renal
handling of sodium and chloride ions.
• In metabolic acidosis, chloride is preferentially excreted by the
kidney.
• In metabolic alkalosis, chloride is retained, and sodium and
potassium are excreted.
• In renal tubular acidosis, there is an inability to excrete Cl- in
proportion to Na+ .
• The diagnosis can be made by observing a hyperchloremic
metabolic acidosis with inappropriately low levels of Cl- in the urine;
the urinary SID is positive.
• If the urinary SID is negative, the process is not renal.
• Gastrointestinal losses (diarrhea, small bowel or pancreatic
drainage), parenteral nutrition, excessive administration of saline;
and the use of carbonic anhydrase inhibitors.
90 percent
of filtered
HCO3–
Ammoniogenesis
actively
synthesize
–
HCO3 in
addition
HPO (pK 6.8)
excretion in the urine is
another mechanism for H elimination
to insecreting
H is trapped
the urine as the acid
H2PO +
H.
Distal
Pretubular Cell
NH4+
HCO3-
1mMol/Kg/day
H+
GLUTAMINE
NH3
2–
4
+
H+
NH3
+
K+
GFR
–
4
HEPATIC HCO3– “PRODUCTION” AND CONSUMPTION
1.
2.
3.
4.
5.
6.
7.
The liver is the principal organ that clears lactic acid produced by
different tissues of the body
Each mole of lactic acid is accompanied by a mole of H+ .
Lactic acid taken up can be metabolized by two pathways; either
oxidation to CO2 , or gluconeogenesis to form glucose and glycogen
Removal of free H+ during lactate metabolism in effect increases the
available HCO3– pool by diminishing its consumption.
Decreased ECF pH stimulates hepatic lactate uptake unless the liver
itself is ischemic or hypoxic.
Countering H+consumption during lactate metabolism is HCO3–
consumption during synthesis of urea from protein and amino acid
catabolism. Urea synthesis, which occurs only in the liver, can be
written empirically as
Each mole of urea synthesis consumes two moles of HCO3– .
Urea produced by the liver is excreted in the urine. A normal daily
excretion of 30g urea in the urine translates to the equivalent of 1,000
mmol of HCO3–