Transcript Slajd 1 - KiZF UMP

ACID
BASE
BALANCE
Dorota Nowak MD, PhD
Internal Medicine Specialist
Dept. of Physiology
University of Medical Sciences
E-mail: [email protected]
Blood H+ concentration is normally maintained within tight limits around
normal value of about
0,00000004 Eq/L = 0,00004 mEq/L = 40 nEq/L
Normal variations are only ± 5 nEq/L
H+ concentration is expressed on logarithm scale, using pH units
ph= - log ( H+ concentration ) = - log ( 0,00000004 ) = 7,4
7,4
8,0
8
In clinical laboratories ph of blood is not estimated using this equation
ph= - log ( H+ concentration )
In clinical laboratories ph of blood is calculated using following method:
1. Concentration of H+ is not estimated at the blood , but is calculated
according to equation
H = K’ x
H2 CO3
HCO3 –
All acids are inoized to some extent (different for every acid).
All ions of strong acid (HCl) dissociate in the solution ( HCl realises all H+ ).
Weak acid (H2CO3) does not realise all H+ into the solution,
it does not completely dissociate.
The concentration of the acid in solution relative to the concentration of
the dissociated ions of the acid is defined by the dissociation constant K’
.
K’ =
H+ x HCO3H2CO3
When the equation is rearange
H+ = K’ x
H2CO3
HCO3-
This equation indicates that in an H2CO3 solution concentration of free H+
can be calculated if concentration of acid (H2CO3) and base (HCO3 – )
are known
In clinical laboratories ph of blood is calculated using following method:
1. Concentration of H+ is not estimated at the blood , but is calculated
according to equation
H+ =
K’ x
H2 CO3
HCO3 –
Concentration of undissociated H2CO3 cannot be measured in solution , because
it rapidly dissociates into CO2 and H2O or to H+ and HCO3CO2 dissolved in the blood is directly proportional to the amount of undissociated
H2CO3 , therefore above equation can be rewritten as
CO2
H+ = K’ x
HCO3Most clinical labolatories measure the blood CO2 tension ( PCO )
2
rather than the actual amount of CO2
In physiological condition 0,03 milimole of CO2 is present in the blood for each
milimeter of mercury PCO
2
measured
H+ = K’
x
( 0,03 x PCO )
2
HCO3
H+ concentration is expressed in pH units
ph= - log ( H+ concentration )
The dissociation constant can be expressed in a similar manner
pK= - log K
In clinical laboratories ph is estimated as (Henderson-Hasselbach equation)
pH = pK + log
HCO3(0,03 x Pco2 )
pK=6,1
for
bicarbonate buffer
system
When body is in acid – base homeostasis the pH of arterial blood is
7,35 – 7,45
The limit of pH at which a person can live more then few hours is
(the lower limit)
6,8
8,0 (the upper limit)
short period of time below
6,8
death
short period of time above
8,0
death
Precise H+ regulation is essential, because the activities of almost all
enzyme systems in the body are influenced by H+ concentration
Enzymes cannot function at altered pH
Enzymes have active site with specific shape, if the shape changes the
enzyme would no longer be able to function.
Change in hydrogen concentration alter all cells and body function !!!
Acids which are produced in the body are the sources of H+
Two types of acids are produced in the body: volatile and nonvolatile acids
Volatile acid
is CO2
is produced from the aerobic metabolism of the cells
CO2 combines with H2O to form the weak acid H2CO3 which dissociates into
H+ and HCO3 - by following reactions
carbonic anhydrase
CO2 + H2O
H2CO3
H+ + HCO3-
Carbonic anhydrase, which is present in most cells, catalyzes reaction between
CO2 and H2O
Two types of acids are produced in the body: volatile and nonvolatile acids
Nonvolatile acids
are also called fixed acids
include sulfuric acid
(a product of protein catabolism)
and posphoric acid (a product of phospholipid catabolism)
are normally produced at a rate of 40 - 60 mmol / day
Other fixed acids, that may be overproduced
ketoacids (in patients with diabetes mellitus)
lactic acid (during physical effort)
Other fixed acids, that may be ingested
salicylic acid
The body has 3 ways of maintaining normal pH range
- chemical buffer system
- respiratory controls
- renal mechanisms
First will be described chemical buffer system, which acts within seconds
to help to maintain ph at constant level
Buffers are composed of weak acid and weak base pairs
Acid - chemical substance that donate hydrogen ions
Strong acids realise all H+ into the solution (all ions dissociate in solution).
Hydrogen acid (HCl) is the only one strong acid which is present in our body
HCl
HCl
H+
H+
Cl-
solution pH = 1
Clwater
As strong acid hits water it completly dissociates into H+ and Cl-
Hydrogen acid is present in the stomach
Weak acid does not realise all H+ into the solution, it does not completely
dissociate.
Carbonic acid is an example of weak acid
H2CO3
H2CO3
H2CO3
H+
HCO3-
solution pH = 4,5
H2CO3
H2CO3
water
Lactic acid release more H+ than H2CO3 and it is stronger weak acid than H2CO3
Lactic acid
Lactic acid
Lactic acid
Lactate-
H+
Lactate-
water
H+
Lactic acid
Solution pH = 2,5
All acids are inoized to some extent (different for every acid).
All ions of strong acid (HCl) dissociate in the solution ( HCl realises all H+ ).
Weak acid (H2CO3) does not realise all H+ into the solution,
it does not completely dissociate.
The concentration of the acid in solution relative to the concentration of
the dissociated ions of the acid is defined by the dissociation constant K’
.
K’ =
H+ x HCO3H2CO3
When the equation is rearange
H+ = K’ x
H2CO3
HCO3-
This equation indicates that in an H2CO3 solution concentration of free H+ can be
calculated if concentration of acid and HCO3 – are known
Base – chemical substance that accept hydrogen ion (takes up H+ from solution)
Only weak bases are found in our body
HCO3- (bicarbonate ions) function as weak base. A small persentage of
HCO3- ions take up H+ from the water. When water molecules lose H+
remains, which increases pH ( solution becomes more alkaline )
( OH- )
HCO3HCO3-
( H2CO3 )
water
HCO3- + H2O
H2CO3 + OH-
solution pH = 8,0
the
OH-
A neutral solution has pH = 7 , because H+ and OH- concentration are equal
H+ = OH-
The acidic solution has greater H+ concentration than OH-
H+ > OH-
pH < 7
The basic (alkaline) solution has H+ concentration less than OHH+ < OH-
pH > 7
Some electrolytes normally found in plasma can act as weak bases :
-
HCO3- (bicarbonate ions)
-
HPO4-2
-
SO4-
-
ions of organic acides (lactate)
-
proteins have basic and acidic side group, may function as weak
(hydrogen phosphate)
(sulfate ions)
bases
or
weak acides ( can accept H+ and also donate H+ )
Some electrolytes normally found in plasma can act as weak acides :
-
H2CO3 (carbonic acid)
-
H2PO4-
-
proteins have basic and acidic side group, may function as weak bases or
weak acides ( can accept H+ or donate H+ )
(dihydrogen phosphate)
Large drop in pH is observed when acid is add to solution without buffer
H+
H+
H2O
H+
H+
H2O
pH = 7
pH = 4
Large increase in pH is observed when base is add to solution without buffer
OHOH-
OH-
H20
pH = 7
OH-
H2O
pH = 9
Chemical buffer system acts within seconds to help to
.
maintain pH
at constant level
Buffer are composed of
pairs
weak acid
and
weak base
The important buffer systems in the body are:
H2CO3 / HCO3- (carbonic acid/bicarbonate) buffer system
H2PO4- / HPO4-2 (dihydrogen phosphate/hydrogen phosphate) buffer system
proteins buffer system
Buffer systems do not eliminate H+ from the body or add them to the body, but
keep pH at the constant level as long as it is possible
A buffer is any substance that can reversibly bind H+
The general form of the buffering reaction is
Buffer + H+
H Buffer
This reaction is reversible
The presence of the buffer system in the body is very important.
About 80 mEq of H+ is ingested and produced by metabolism each day,
whereas
H+ concentration of the body fluids is only about 0,00004 mEq/L
Without buffering, the daily production or ingestion of acids would cause
large changes in body fluid pH
The bicarbonate buffer is the most important extracellular fluid buffer system
The bicarbonate buffer system consist of a water solution that contains two
ingridients : weak acid
..
H2CO3
weak base – bicarbonate salt such as NaHCO3 (sodium
bicarbonate)
Bicarbonate buffer
H2CO3 is formed in the body by the reaction of CO2 with H2O
CO2 + H2O
carbonic anhydrase
H2CO3
Without enzyme ( carbonic anhydrase ) reaction is very slow
This enzyme is especially abundant
- in the walls of the lung alveoli, where CO2 is release
- in the epithelial cells of renal tubules, where CO2 reacts with H20 to form
H2CO3
When a strong acid is added to the solution with buffer, the increased H+
released from the acid is buffered by HCO3H+ + HCO3-
H2 CO3
CO2 + H2O
When a strong base , such as sodium hydroxide is added to the buffer,
OH- combines with H2CO3
NaOH + H2CO3
NaHCO3 + H2O
Thus weak base NaHCO3 replaces the strong base NaOH
At the same time , the concentration of H2CO3 decreases
causing more CO2 to combine with H2O to replace the H2CO3
As the net result CO2 levels in the blood decreases, but the decreased CO2 in
the blood inhibits respiration and decreases the rate of CO2 expiration.
The rise in blood HCO3- that accurs is compensated by increased renal
excretion of HCO3-
Tritation curve for bicarbonate buffer system
When the two components of the buffer system are equal,
the pH of the solution is the same as the pK (6,1)
pH= 6,1 + log
HCO30,03 x Pco2
When acid is added, it is buffered by HCO3 , which is then
converted into dissolved CO2 , decreasing the ratio of HCO3
to CO2 and decreasing the pH of the extracellular fluid
When base is added part of the CO2 is converted into HCO3causing an increase in the ratio of HCO3 to CO2 and
increasing the pH
The „buffer power” of buffer system
The buffer system is the most effective in the central part of the curve, where pH is near pK
of system and is still reasonably effective for 1,0 pH unit on either side of the pK (pH of 5,1 to 7,1)
Beyond these limits, the buffering power rapidly diminishes (above 7,1 to buffer bases;
below 5,1 to buffer acides)
The „buffer power” is determined also by the concentration of the buffer
Bicarbonate buffer system is the most important extracellular buffer
although the concentration of the two elements of the bicarbonate system CO2 and HCO3are not great.
Bicarbonate buffer system is the most important extracellular buffer, because concentration
of the two elements of the buffer is regulated by the kidneys and lungs.
The pH of the extracellular fluid can be precisely controlled by
- the rate of removal or addition of HCO3- by the kidneys
- and the rate of removal of CO2 by the lungs.
Phosphate buffer system
It plays major role in buffering renal tubular fluid and intracellular fluids.
Phosphate buffer is composed of weak acid H2PO4- (dihidrogen phosphate) and
weak base HPO4-2 (hidrogen phosphate)
When strong acid is added to the buffer H+ is accepted by the base HPO4-2
H+ + HPO4-2
H2PO4-
When strong base is added to the buffer OH- is beffered by acid H2PO4OH- + H2PO4-
HPO4-2 + H2O
The phosphate buffer system has pK=6,8 which is not far from the normal pH of 7,4 in body
fluids. This allows system to operate near its maximum buffering power.
Concentration of phosphate buffer in extracellular fluids is only 8 per cent of the
concentration of bicarbonate buffering system. Total buffering power of the phosphate
buffer in the extracellular fluid is much less than that of the bicarbonate buffering system.
Protein buffer system
It is very important buffer system, because of high protein concentration, especially within
cells. pK of protein buffer system is fairly close to 7,4.
The buffer system within cells help prevent changes in the pH of extracellular fluid but may
take several hours to become maximally effective. H+ and HCO3- require several hours to
come to equilibrium within the extracellular fluid, because they diffuse slowly through the
cell membrane. CO2 can rapidly diffuse through cell membranes.
Rapid equilibrium for H+ , HCO3+ and CO2 between intra and extracellular fluid occurs in
the red blood cells.
Hemoglobin is an important buffer in red blood cells.
Approximately 60 – 70% of the total chemical buffering of the body fluids is inside the cells.
Respiratory regulation of acid-base balance
The rate of ventilation controls concentration of CO2 in extracellular fluid.
An increase in ventilation eliminates CO2 from extracellular fluid, which reduces
the H+ concentration.
A decrease in ventilation decreases elimination of CO2 from extracellular fluid, which
increases the H+ and H2CO3 concentration
CO2 is formed continually in the body by intracellular metabolic processes
Fedback control of hydrogen ion concentration by respiratory system
Increased H+ stimulates respiration, increased alveolar ventilation deceases the H+ concentration
Respiratory system acts as a negative feedback controller of H+ concentration
H+
Alveolar ventilation
Pco2
Respiratory control cannot return the H+ concentration
all the way back to normal when a disturbance outside
the respiratory system has altered pH
Respiratory mechanism for controlling H+ concentration
has an effectivenees between 50 -75 %
When pH falls from 7,4 to 7,0 the respiratory system
can return the pH to a value of about 7,2 to 7,3
This response occurs within 3 to 12 minutes.
Impairement of lung function can cause respiratory acidosis
Severe emphysema decreases the ability of the lungs to eliminate CO2 which causes
increase of CO2 concentration in extracellular fluid. It leads to respiratory acidosis
Renal control of acid-base balance
The kidneys control acid-base balance by excreting either an acidic or basic urine and
by generating of HCO3- by the tubular cells.
The mechanism by which the kidneys excrete acidic or basic urine is as follows:
- large amounts of HCO3- are filtered continously into the tubules, if they are excreted
with urine it removes bases from the blood
- H+ are filtered and also secreted into the tubular fluid. If more H+ than HCO3- is
excreted with urine the net loss of acid from extracellular fluid occurs
Bicarbonate reabsorption and hydrogen ion secretion occur in all parts of the tubules
except the descending and ascending thin limbs of the loop of Henle
Typically 85% of filtered HCO3- is reabsorbed in the PCT
HCO3- cannot diffuse through luminal side of PCT cell membrane
HCO3- combines with H+ in the fitered fluid to form H2CO3
Carbonic anhydrase which may be atached to the microvilli of
PCT cell catalizes break of H2CO3 to CO2 and H2O
CO2 diffiuses into PCT cell
Carbonic anhydrase, which is present inside the cell catalizes
formation of H2CO3
Next H2CO3 splits to HCO3- and H+
HCO3- difuses to intestitium
H+ are secreted to the tubular fluid
The net result is that for every H+ secreted into the tubular lumen an HCO3- enters the blood
Under normal conditions the rate of tubular H+ secretion is about 4400 mEq/day, the rate of
HCO3- filtration is 4320 mEq/day. Slight excess of H+ ( 80 mEq/day ) rids the body fluids of
nonvolatile acids produced by metabolism. Most of H+ is not excreted as free H+ but in
combination with other urinary buffers, especially phosphate and amonia
1. CO2 diffuses from the blood into intercalated cell
2. CO2 combines with H2O to form H2CO3 _
carbonic
anhydrase
3. H2CO3 splits to HCO3- and H+
.
4. H+ are secreted by primary active
transport to tubular fluid
5. HCO3 are reabsorbed to the blood
For each H+ secreted an HCO3 is reabsorbed in late distal and collecting tubules
Hydrogen phosphate (weak base)
Dihydrogen phosphate (weak acid)
comes mainly from the
metabolism of amino
acids in the liver
2NH3+ 2H+
Amonia (weak base)
Amonium (weak acid)
This process occurs in severe acidosis
Renal acid-base excretion can be measured
Bicarbonate excretion is calculated as urine flow rate multiplied by urinary bicarbonate
concentration. It indicates how rapidly the kidneys are removing HCO3- from the blood
( which is the same as adding an H+ to the blood )
H+ excretion is calculated by
- measuring NH4- excretion ( urine flow rate multiplied by urinary NH4- concentration )
- determing the value known as titratable acid ( it is measure by titrating the urine with
a strong base to return the urinary pH to 7,4 which equals mEq of H+ that combined with
phosphate).
Net acid excretion by the kidneys can be measure as:
Net acid excretion= NH4+ excretion + urinary titratable acid – bicarbonate excretion
Regulation of potassium ions distribution
Acidosis ( H+)
Alkalosis ( H +)
pH=7,4
K+
cell
Extracellular fluids hyperkaliemia
K+
cell
Extracellular fluids hypokaliemia
At the kidneys
At the kidneys
decreased H secretion
increased H secretion
and decreased HCO3
reabsorption
and increased HCO3 reabsorption
Extracellular fluid potassium concentration normally is regulated precisely at about
4,2 ± 0,3 mEq/L
Many cells function is very sensitive to changes in extracellular fluid potassium
concentration. A decrease in plasma potassium concentration leads to ventricular
arrhythmias and ventricular fibfillation. Also an increase in potassium concentration can
lead to cardiac arrhythmias and cardiac arrest.
Redistribution of potasium
98% (3950 mEq) of total body potassium is contained in the cells,
only 2% (59 mEq) in the extracellular fluid.
The potassium contained in a single meal is ofen as high as 50 mEq, the daily intake usually
ranges between 50 -200 mEq. Failure to rapidly rid the extracellular fluid of the ingested
potassium could cause life-threatening hyperkalemia.
A small loss of potassium from the extracellular fluid could cause severe hypokaliemia
Maintenance of potassium balance depends on:
.
- excretion by the kidneys. Kidneys must be able to control potassium
..
excretion rapidly and precisely to wide variations in intake
..
- control of potassium distribution between the extracellular and intracellular
components
Regulation of potassium ions distribution
Most of the ingested potassium rapidly moves into the cells until the kidneys can eliminate
the excess
Factors that shift potassium into cells:
- insulin
- aldosterone (hyperkaliemia stimulates aldosteron secretion)
- catecholamines (by β-adrenergic stimulation)
- alkalosis
Factors that shift potassium out of cells:
- insulin deficiency (diabetes mellitus)
- aldosterone deficiency (Adisson’s disease)
- cell lysis (severe muscle injury, red blood cell lysis)
- strenuous exercise (during prolonged exercise potassium is released from
skeletal muscle)
- increased extracellular fluid osmolarity
- acidosis (decreases sodium-potasium ATPase activity)
Renal potassium excretion is determined by
- the rate of potassium filtration (in healthy person is constant)
- the rate of potassium reabsorption is constant in PCT- 65% and loop of Henle – 27%,
potasssium can be also reabsorbed by intercalated cells by hydrogen-potassium ATPase
transport mechanism (only during extracellular fluid potassium depletion)
- the rate of potassium secretion by principal cells is the main process, which controls
potassium excretion.
Potassium secretion is stimulated by
1. Increased extracellular fluid potassium concentration
2. Increased aldosterne secretion
3. Increased tubular flow rate
4. Alkalosis
Potassium secretion is inhibited by
1. Acidosis, which decreases activity of sodium-potassium ATPase pump, but
chronic acidosis leads to loss of potassium
Clinical causes of acid-base disorders (acidosis)
pH = pK + log
HCO30,03 x Pco2
According to Henderson-Hasselbach equation acidosis occurs when the ratio of
HCO3-
decreases in extracellular fluid, decreasing pH.
CO2
If this ratio decreases because of a fall in HCO3 , the acidosis is referred to as metabolic
acidosis.
If this ratio decreases because of an increase in CO2 , the acidosis is referred to as
respiratory acidosis.
The ratio of HCO3- / CO2 is also decreased in renal tubular fluid in patients with both types
of acidosis. As a result, there is an excess of H+ in renal tubules.
In physiological condition about 80 mEq of H+ / day (acids formed by metabolism) must be
excreted with urine.
With severe chronic acidosis , as much as 500 mEq/day of H+ can be excreted in the urine,
mainly in the form of NH4+ , which causes formation of 500 mEq/day of new HCO3- that is
added to the blood.
Respiratory acidosis may be caused by
- decreased alveolar ventilation
.
a. demage of respiratory center in medulla oblongata
b. obstruction of the bronchi
- decreased pulmonary membrane surface area
a. pneumonia
- decreased permeability of pulmonary membrane ( factors that interferes with the
. exchange of gases between the blood and alveolar air)
a. odema of lungs
Metabolic acidosis may be caused by
- failure of the kidneys to excrete metabolic acids normally formed in the body, which is
caused by 2 types of disorders
a. impairment of renal tubular HCO3- reabsorption
b. impairment of renal tubular H+ secretion
- loss of base from body fluids
a. diarrhea
b. vomiting of intestinal contents
- formation of excess quantities of metabolic acids
a. diabetes mellitus – in diabetic patients fats, which are split into acetoacetic acid
. become the source of energy. With severe diabetes mellitus blood acetoacetic acid levels
. can rise very high causing acidosis.
-ingestion of acids
- increased potassium concentration in extracellular fluid
Chronic renal failure decreases glomerular filtration rate, which decreases the amount of
filtered phosphates and NH3+ , it reduces the amount of bicarbonate added back to the body
fluids.
Chronic renal failure can be associated with severe metabolic acidosis
Clinical causes of acid-base disorders (alkalosis)
pH = pK + log
HCO30,03 x Pco2
According to Henderson-Hasselbach equation alkalosis occurs when the ratio
of HCO3- / CO2 increases in extracellular fluid, increasing pH.
If this ratio increases because of an increase in HCO3- the alkalosis is referred to as
metabolic alkalosis.
If this ratio increases because of an decrease in CO2 the alkalosis is referred to as
respiratory alkalosis.
Respiratory alkalosis is caused by hypoxia
( when concetration of the oxygen in inspeare air is decreased )
Such condition occurs e.g. in high mountains.
Metabolic alkalosis is caused by
- retention of HCO3-
a. ingestion of alkaline drugs, such as sodium bicarbonate for the treatment of gastric
or peptic ulcer
- loss of H+
a. excess aldosterone, which promotes reabsorption of sodium ions which is coupled
with secretion of H+ in distal tubule and collecting duct
b. vomiting of gastric contents
Quiz
1.
Which parameters are necessary to calculate plasma pH?
1. plasma concentration of H+
2. plasma concentration of H2CO3
3. plasma concentration of HCO3
4. plasma CO2 tension
A. 1
B. 1,2
C. 1,2,3,4
D. 2,3,4
E. 3,4
2. What is the most important extracellular buffer system?
A. H2CO3 / HCO3- buffer system
B. H2PO4 / HPO4-2 buffer system
C. Protein buffer system
D. Amonia / amonium buffer system
3. In acidosis, most of the H+ secreted by the proximal tubule are associated with which of the following
processes?
A. Excretion of NH4+
B. Reabsorption of HCO3C. Reabsorption of phosphate ions
D. Reabsorption of potassium ions
4. Which of the following tends to decrease potassium secretion by the cortical collecting tubule?
A. Increased plasma potassium concentration
B. A diuretic that decreases proximal tubule sodium reabsorption
C. A diuretics that inhibits the action of aldosterone (e.g. spironolacton)
D. Acute alkalosis
E. High sodium intake
5. The clinical laboratory returned the following values for arterial blood taken from a patient:
plasma pH= 7,28
plasma HCO3- = 32 mEq/L
plasma p CO2 = 70 mmHg
What is the patient’s acid-base disorder?
A. Metabolic alkalosis
B. Metabolic acidosis
C. Respiratory alkalosis
D. Respiratory acidosis
Correct answers
1.E
2.A
3.B
4.C
5.D
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