Acid-base balance. pH homeostasis.

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ACID-BASE BALANCE. PH HOMEOSTASIS. INTRODUCTION An important property of the internal environment (and also the blood) is its degree of acidity or alkalinity. Acidity increases when the level of acidic compounds in the body rises (through increased intake or production, or decreased elimination), or when the level of basic (alkaline) compounds in the body falls (through decreased intake or production, or increased elimination). Alkalinity increases with the reverse of these processes. The body's balance between acidity and alkalinity is referred to as acid-base balance or acid-base homeostasis. The acidity or alkalinity of any solution, including blood, is indicated on the pH scale. Human homeostasis refers to the body's ability to physiologically regulate its inner environment to ensure its stability in response to fluctuations in the outside environment. Several organs help maintain homeostasis including the liver, the lungs, the kidneys, the autonomic nervous system and the endocrine system. An inability to maintain homeostasis may lead to death or a disease, a condition known as homeostatic imbalance. The blood's acid-base balance is precisely (tightly) controlled, because even a minor deviation from the normal range can severely affect many organs. The body uses several different mechanisms to control the blood's acid-base balance.

CHEMICAL BACKGROUND The hydrogen atom consists of only one proton and one electron. When dissociated (i.e. oxidation of hydrogen) in the sense of removing its electron, formally gives H+, containing no electrons and a nucleus composed of one proton. That is why H+ is often called a proton. Distilled water is water that has virtually all of its impurities removed through distillation. Water, however pure, is not a simple collection of H2O molecules. Even in "pure" water, sensitive equipment can detect a very slight electrical conductivity. This must be due to the presence of ions H+ and OH- (or H3O+ and OH-).

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A bare proton, H+, cannot exist in solution or crystals, because of its attraction to other atoms or molecules with electrons. Except at the high temperatures associated with plasma, such protons cannot be removed from the electron clouds of atoms, and will remain attached to them. However, the term "proton" is used loosely and metaphorically to refer to positively charged hydrogen attached to other species (e.g. H2O and as such is denoted "H+" without any implication that any single protons exist freely as a species. The dissociation constant (the ratio of dissociated/non-dissociated molecules) of water is described by (1) and this can be determined by several experimental methods. The water ionization constant (i.e. ionic product of water) is described by (2) and can be exactly computed. The ratio of the two ions is represented by (3). Because distilled water is made only from water molecules that dissociate to H+ and OH- ions, their concentrations must be equal (3).

ka 

H  OH  



H 2 O

(1)

  



k w  H   OH   10 14 mol / l

(2)

H   1  H   OH  OH 

(3)









Thus the concentration of H+ in distilled water will be 10-7 mol/l The pH value is the negative logarithm (base 10) of the molar concentration of dissolved hydrogen ions (4). Thus pH will have typical values between 0-14, 7.00 being considered neutral (i.e. the concentration of H+ at 25°C is approximately 1.0×10−7 mol/l).

 

pH   log H 

The concept of pH was first introduced by Danish chemist Søren Peder Lauritz Sørensen at the Carlsberg Laboratory in 1909. The p probably stands for “Power” although this is debated.

Acid-base balance. pH homeostasis.

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In order to maintain the body's acid-base balance buffer solutions are used. These have the property that the pH of the solution changes very little when a small amount of strong acid or base is added to it. Usually buffer solutions are aqueous solutions consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid. In conclusion buffer solutions reversibly bind hydrogen ions and impede any change in pH. When added to water a strong acid will dissociate completely:

 

HCl  H   Cl   H    pH  When added to water a strong base will dissociate completely:





NaOH  Na   OH   OH    pH  Thus when adding a strong base or acid to a solution due to the complete dissociation there will be massive changes in the pH. If we mix an acid and a base we will obtain a salt:

HCl  NaOH  H 2 O  NaCl

When mixing an acid with the salt created by the reaction between a weak acid (i.e. an acid that dissociates incompletely) and a strong base the result will be a weak acid, with less dissociation and thus a smaller modification of the pH. In the example below we obtained acetic acid with a very low dissociation constant, and as such no major effect on pH. CH3-COOH ↔ CH3-COO- + H+

CH 3COONa  HCl  CH 3COOH  NaCl

If we mix further this weak acid (acetic acid) with its salt (sodium acetate) the buffer capacity increases (i.e. the pH will not be modified significantly if we add strong acids or strong bases to the solution). Thus buffer solutions are made up from weak acids and their salts (ex. CH3COOH / CH3COONa).

CH 3 COOH  NaOH  CH 3COONa  H 2 O CH 3 COONa  HCl  CH 3 COOH  NaCl

BUFFER SOLUTIONS OF THE HUMAN BODY The human body has several buffer systems. Some of these are located in the blood and extracellular space: – bicarbonate (H2CO3/NaHCO3, carbonic acid/sodium-bicarbonate)

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– hemoglobin Other buffer systems are found in the intracellular space: – phosphate (Na2HPO4/NaH2PO4, disodium hydro-phosphate/sodium dihydro-phosphate) – proteins From these buffers one of the most important is the bicarbonate buffer system as it: – reversibly binds hydrogen ions thus preventing pH changes – can shift carbon dioxide through carbonic acid to hydrogen ions and bicarbonate and vice-versa – allows quick buffering (although low buffer capacity) – acid-base imbalances that overcome this buffer system can be compensated in the short term by: – changing the rate of ventilation; this alters the concentration of carbon dioxide in the blood, shifting the reaction (see below), which in turn alters the pH (ex. if pH drops it can be compensated through increased breathing, thus eliminating the CO2) – excretion of excess acid or base; the kidneys are slower to compensate, but there are several powerful renal mechanisms to control pH 

HCO3  H   H 2 CO3  CO2  H 2 O The proteins have the highest buffering capacity of all buffer solutions of the human body but this is a slow system. The principle mechanisms to control pH are: dilution (i.e. diluting any acid or base in the large volume of water in the organism), buffering (i.e. buffering using buffer solutions with involvement of the respiratory and renal systems) and restoration of normal state.

PH OF THE HUMAN INNER ENVIRONMENT

Probably the simplest way to calculate the pH is to use the HendersonHasselbalch equation and apply the discussed chemical background to the bicarbonate buffer system.

Acid-base balance. pH homeostasis.

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H  HCO  k 



H 2 CO3 

log k  log

3

H  HCO   logH   log HCO  H CO  H CO  



3

2

 





3

3

HCO   log k

2

3



 log H   log

 

HCO  

3

H 2 CO3 

 pH  pk  log

pH   log H  and  log k  pk

3

H 2 CO3 

which is the Henderson-Hasselbalch equation (where pk is the dissociation constant of H2CO3) Lawrence Joseph Henderson wrote an equation, in 1908, describing the use of carbonic acid as a buffer solution. Karl Albert Hasselbalch later reexpressed that formula in logarithmic terms, resulting in the Henderson– Hasselbalch equation. Because the bicarbonate buffer is in open system through the lungs the concentration of H2CO3 will be dependent on the elimination of CO2; in fact it depends on the solubility coefficient of CO2 (a) and its partial pressure (pCO2 the partial pressure of a gas dissolved in a liquid is the partial pressure of that gas, which would be generated in a gas phase in equilibrium with the liquid at the same temperature). Thus the Henderson-Hasselbalch equation can be modified to reflect this.

HCO   pk  log HCO  pH  pk  log 



3

H 2 CO3 

a  0.03

3

a  pCO2





mEq  ; pCO2  40 mmHg ; HCO3  24 mEq / l ; l  mmHg

pk  6.1  pH  6.1  log 20  6.1  log 10  log 2  6.1  1  0.3  7.4

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There are some significant approximations implicit in the Henderson– Hasselbalch equation. The most significant is the assumption that the concentration of the acid and its conjugate base at equilibrium will remain the same as the former concentration. This neglects the dissociation of the acid and the hydrolysis of the base. The dissociation of water itself is neglected as well. Also, the equation does not take into effect the dilution factor of the acid and conjugate base in water.

ACID-BASE BALANCE PARAMETERS ► pH: is the measure of the acidity or alkalinity of a solution. In arterial blood the normal range is 7.36-7.44. Blood pH values compatible with life in mammals are limited to a pH range between 6.8 and 7.8. ► pCO2: partial pressure of carbon dioxide; its value is dependent on ventilation (breathing). The brain regulates the amount of carbon dioxide that is exhaled by controlling the speed and depth of breathing. The amount of carbon dioxide exhaled, and consequently the pH of the blood, increases as breathing becomes faster and deeper. By adjusting the speed and depth of breathing, the brain and lungs are able to regulate the blood pH minute by minute. In arterial blood the normal value is 40-44 mmHg. ► standard HCO3- (standard bicarbonate): concentration of bicarbonate at a pCO2 of 40 mmHg, oxygenation 100% and t=37°C. Normal value in arterial blood is 24 mEq/l. ► total CO2: in arterial blood the normal value is 25 mEq/l (= [HCO3] + 0.03 x pCO2) ► reserve alkalinity: the concentration of bases in plasma after the neutralization of stronger acids than carbonic acid; normal range in venous blood is 24-29 mmol/l. ► buffer base: is the sum of all buffering agents in the blood, normal value 46 mEq/l. Normal buffer base (NBB) is what the buffer base would have been at normal pH (7.4), normal pCO2 (40 mmHg) and normal temperature (37°C). The content of hemoglobin in blood increases this value.

Acid-base balance. pH homeostasis.

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CO2  H 2O  H 2 CO3 H 2 CO3  Hb  HHb  HCO3



The correlation (empirical formula) is: NBB = 41.7+0.68×Hb, where Hb is in mmol/l (the normal value hemoglobin is ~9.3 mmol/l). ► base excess: refers to an excess or deficit in the amount of base present in the blood. The value is usually reported as a concentration in units of mEq/l, with positive numbers indicating an excess of base and negative a deficit. The reference range for base excess is −2 to +2 mEq/l. Comparison of the base excess with the reference range assists in determining whether an acid/base disturbance is caused by a respiratory, metabolic, or mixed metabolic/respiratory problem. While carbon dioxide defines the respiratory component of acid-base balance, base excess defines the metabolic component. Accordingly, measurement of base excess is defined under a standardized pressure of carbon dioxide, by titrating back to a standardized blood pH of 7.40. This parameter can be used to calculate the necessary amount of sodium bicarbonate to neutralize the extracellular space (in acid build-up). NaHCO3 = BEmeasured x 0.3 x body weight (kg) - for adults (water in extracellular space is 1/3 of body weight) NaHCO3 = BEmeasured x 0.4 x body weight (kg) - for children (water in extracellular space is 40% of body weight) ► anion gap: it is calculated by subtracting the serum concentrations of chloride and bicarbonate (anions) from the concentrations of sodium plus potassium (cations). Na+

140 mmol/l

Cl-

K+

4 mmol/l 144 mmol/l

HCO3-

110 mmol/l 24 mmol/l 134 mmol/l

If there are more unmeasured anions (compared to unmeasured cations) in the serum the anion gap is higher (i.e. higher than the normal value of 10). The anion gap varies in response to changes in the concentrations of the above mentioned serum components that contribute to the acid-base balance and increases when fixed (non-volatile) acids accumulate in the organism.

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DEFINITIONS Acidosis – increased acidity due to excess acid or deficit of bases. Acidemia means low blood pH (i.e. pH < 7.36); acidosis is used to describe the processes leading to increased acidity. Incorrectly physicians use these terms in an interchangeable fashion. Alkalosis – reduced hydrogen concentration of arterial blood plasma Metabolic – acid-base disturbance caused by deficit or build-up of acids or bases in the extracellular space Respiratory – acid-base disturbance caused by modification of ventilation E.g. metabolic acidosis means: excessive blood acidity caused by an overabundance of acid in the blood or a loss of bicarbonate from the blood. E.g. respiratory acidosis means: buildup of carbon dioxide in the blood that results from poor lung function or slow breathing Simple acid-base disorder– one single etiological base, either metabolic or respiratory disturbance (i.e. the presence of only one of the possible derangements) Mixed disorder– more than one cause Compensated – the disturbance of one system (i.e. renal or respiratory) altered the pH, this can be compensated by the opposite system, and thus the pH can be corrected (the pH returns to normal). E.g. hypoventilation causes the build-up of carbon dioxide and in turn the decrease of pH (i.e. respiratory acidosis), this will be compensated by the increase of bicarbonate to buffer the free H+. If the compensation is complete then the pH is normal but both the pCO2 and bicarbonate concentrations are altered. Partially compensated – similar to the above, but the compensation is not enough to return the pH to normal value. Not compensated – there is no compensation of the acid-base disturbance.

ACID-BASE DISTURBANCES Respiratory acidosis - caused by diseases that increase the pCO2, which in term will increase the hydrogen ion concentration (i.e. decrease the pH). There is no primary modification of the concentration of bicarbonate.

Acid-base balance. pH homeostasis.

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Metabolic acidosis - caused by diseases that decrease the concentration of bicarbonate, thus the concentration of hydrogen ions increase (i.e. decrease the pH). There is no primary modification of the partial pressure of carbon dioxide. Respiratory alkalosis - caused by diseases that decrease the pCO2, which in term will decrease the hydrogen ion concentration (i.e. increase the pH). There is no primary modification of the concentration of bicarbonate. Metabolic alkalosis - caused by diseases that increase the concentration of bicarbonate, thus the concentration of hydrogen ions decrease (i.e. increase the pH). There is no primary modification of the partial pressure of carbon dioxide. Use the Siggaard-Andersen nomogram and the given examples to determine the acid-base balance parameters.

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