Acid-Base Balance · ASSESSING ACID-BASE BALANCE An indication of the acid-base status of a patient...
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Acid-Base Balance PS Ogunro
Transcript of Acid-Base Balance · ASSESSING ACID-BASE BALANCE An indication of the acid-base status of a patient...
AIMS OF THE LECTURES
1. Understand the basic biochemistry and physiology of acid-base balance.
2. Understand human acid-base balance and interpret clinical acid-base data,
3. Understand the diseases that cause acid-base disturbance.
HYDROGEN ION CONCENTRATION and CONCEPT OF pH
Blood hydrogen ion concentration (abbreviated [H+]) is maintained within tight limits in health, with the normal concentration being between 35 - 45 nmol/l.
Concentrations below 20 nmol/l or above 120 nmol/l are generally incompatible with life.
Blood hydrogen ion concentration is often expressed as pH.
The [H+] when expressed in mol/l is 3.5 - 4.5 x 10-⁸ mol/l, and such negative exponential numbers are difficult to work with, therefore SORENSON formulated a term, pH, which describes the free H+ concentration. The definition of pH is :
pH = -log [H+]
when [H+] = 4.0 x 10-8 mol/l
then pH = (-log 4.0) + (-log 10-8)
= -0.6 + 8
= 7.4
Note the relative sizes of [H+] and pH :
[H+] = 1 x 10-6 [H+] = 1 x 10-7 [H+] = 1 x 10-8
pH = 6 pH = 7 pH = 8
i.e., for every 10 fold increase in [H+]
pH decreases by 1.
SOURCES OF HYDROGEN IONS
Hydrogen ions are produced in the body as a result of metabolism.
1. The oxidation of proteins, nucleic acids and phospholipids produces phosphoric and sulphuric acids, while the incomplete (anaerobic) metabolism of fat and carbohydrates produces organic acids such as lactic, acetoacetic and β-hydroxybutyric acids.
2. Complete (aerobic) metabolism of fat and carbohydrates produces CO2. In solution, CO2forms a weak acid (carbonic acid) which therefore has the potential to affect [H+] and pH.
ACID = Substance that dissociates to produce H+ ions,
HA ↔ H+ + A- e.g., H3PO4 ↔ H+ + H2PO4-.
Acids dissociate in water to varying degrees, depending on their strength.
BASE = Substance that accepts H+ ions,
e.g., H2PO4- + H+ ↔ H3PO4.
One mechanism of accepting H+ ions, is to produce OH- ions, which with H+ ions forms water,
e.g., NaOH + H+ ↔ Na+ + H2O.
Bases dissociate in water to varying degrees, depending on their strength
SALT An ionic compound, where the positive ion (cation) is anything except H+, and the negative ion (anion) is anything except OH-. Salts dissociate completely in water.
STRENGTH OF ACIDS The strength of an acid is defined by its tendency to
dissociate, thereby producing free hydrogen ions
A strong acid dissociates completely even in acidic solutions e.g., H2SO4 → H+ + HSO4
-
A weak acid only dissociates partially in acidic solutions, reaching a state of equilibrium between the acid HA and its conjugate base A-e.g., H3PO4 ↔ H+ + H2PO4-
H2CO3 ↔ H+ + HCO3-NH4
+ ↔ H+ + NH3
The strength of an acid is measured by its dissociation constant K :
HENDERSON-HASSELBALCH EQUATION
The pH of a solution containing a conjugate acid-base pair is described by the HENDERSON-HASSELBALCH equation. Since pH depends on the free [H+], it therefore depends on the degree of dissociation of the acid i.e., it depends on the pK of the acid
[H+] [A-]
K = --------------
[HA]
K [HA]
[H+] = ---------
[A-]
[HA]
-log [H+] = -log K + (-log ------ )
[A-]
[A-] base
pH = pK + log ------
[HA] acid
the basic HENDERSON - HASSELBALCH equation
ASSESSING ACID-BASE BALANCE
An indication of the acid-base status of a patient can be obtained by measuring the components of the bicarbonate buffer system. Because of the sites of excretion and regulation, the pCO2 is called the respiratory component, and the HCO3
- the metabolic component of the bicarbonate buffer system.
NORMAL VALUES
pH 7.35 - 7.45
pCO2 4.5 - 6.1 kPa
[HCO3-] 22 - 26 mmol/l
Buffer Systems and Their Role in Regulating the pH of Body Fluids.
A buffer is a mixture of a weak acid and a salt of its conjugate base that resists changes in pH when a strong acid or base is added to the solution.
A buffer is a solution containing a conjugate acid-base pair, made up of a weak acid and its salt, which minimises changes in pH.
The body contains a number of buffers. Proteins can act as buffers by binding hydrogen ions, and haemoglobin in red blood cells, in particular, has a high capacity for binding hydrogen ions.
proteins Pr- + H+ ↔ PrH
haemoglobin Hb- + H+ ↔ HbH
In extracellular fluid the most important buffer system is however the bicarbonate system. In this buffer system, the base bicarbonate (HCO3
-) combines with hydrogen ions to form the weak acid carbonic acid (H2CO3).
HCO3- + H+ ↔ H2CO3
This buffer system is unique because of two factors:
1. H2CO3 dissociates to H2O and CO2
2. HCO3- is retained and regenerated by the kidneys.
1. Bicarbonate/Carbonic Acid Buffer System
As mentioned previously, eventually all hydrogen ions produced must be excreted from the body. Buffering of hydrogen ions is a vital short-term solution to the problem of maintaining a constant and normal pH, and adequate amounts of bicarbonate must therefore be maintained to preserve the blood’s buffering capacity.
The bicarbonate–carbonic acid system has low buffering capacity, it still is an important buffer for
three reasons:
(1) H2CO3 dissociates into CO2 and H2O, allowing CO2
to be eliminated by the lungs and H as water;
(2) changes in CO2 modify the ventilation (respiratory) rate; and
(3) HCO3- concentration can be altered
by the kidneys
3. Plasma Protein Buffer System and Plasma Base Excess
Plasma protein, especially the imidazole groups of histidine, also forms an important buffer system in plasma. Most circulating proteins have a net negative charge and are capable of binding H.
4.Haemoglobin Buffer System and Whole Blood Base Excess
CO2, produced by complete (aerobic) metabolism of fat and carbohydrates, diffuses out of cells into the ECF. In the ECF a small amount combines with water to form carbonic acid, thereby increasing the [H+] and decreasing the pH of the ECF.
In red blood cells metabolism is anaerobic and no CO2 is formed. CO2 therefore diffuses into red blood cells down a concentration gradient
In the red blood cells the majority of the CO2 combines with water to form carbonic acid, due to the presence of carbonic anhydrase. The carbonic acid dissociates to form hydrogen ions and bicarbonate ions, and the hydrogen ions are bound by the haemoglobin. Deoxygenated haemoglobinbinds hydrogen ions more strongly than oxygenated haemoglobin, and in fact the binding of hydrogen ions to haemoglobin facilitates the release of oxygen (the Bohr effect). The overall effect of this process is that CO2 is converted to bicarbonate in red blood cells.
ROLE OF THE KIDNEY IN HANDLING OF BICARBONATE AND HYDROGEN
As mentioned previously, eventually all hydrogen ions produced must be excreted from the body. Buffering of hydrogen ions is a vital short-term solution to the problem of maintaining a constant and normal pH, and adequate amounts of bicarbonate must therefore be maintained to preserve the blood’s buffering capacity.
ASSESSING ACID-BASE BALANCE
An indication of the acid-base status of a patient can be obtained by measuring the components of the bicarbonate buffer system. Because of the sites of excretion and regulation, the pCO2 is called the respiratory component, and the HCO3
- the metabolic component of the bicarbonate buffer system.
NORMAL VALUES
- pH 7.35 - 7.45
- pCO2 4.5 - 6.1 kPa
- [HCO3-] 22 - 26 mmol/l
size of buffer pools bicarbonate 24 mmol/l
haemoglobin 10 mmol/l
plasma protein 10 mmol/l
size of free [H+] pool 40 nmol/l (40 x 10-6 mmol/l) i.e., a million fold smaller
