UNIT  4 --  BLOOD, HEMOGLOBIN, AND ACID - BASE CONTROL

All objectives listed are in the cognitive domain unless otherwise noted. The student, at the end of the instructional period, is responsible for meeting these objectives by achieving a cumulative score of 70% or better on all problem sets, case studies, major exams, "pop" quizzes, and library assignments.

Objectives are indicated by a number followed by a purposeful question or statement. The student, upon completion of the classroom component of clinical biochemistry, will be responsible to successfully:

01    DISCUSS pH AND ITS IMPORTANCE TO HOMEOSTASIS

Maintenance of pH is important to proper physiological functioning of cells and tissues. Any changes in pH can alter enzyme activity, cellular uptake, incorporation and use of minerals and metabolites, uptake and release of oxygen, and the formation of biological structural components. Normal plasma pH = 7.40 (±0.05). The pH range that is compatible with life is from 6.8 to 7.8. The body can comfortably tolerate a shift in pH of about 0.04. Most cells of the body have a pH = 7.0, but RBC's boast a pH of 7.2. pH of the body affects its acid-base balance and the pH of blood has the greatest effect. The most common sources for pH disturbances are the body's production of organic acids (acetic, acetoacetate, propionic, butyric, lactic, etc.), which are the major sources of hydrogen ion. The body has to keep hydrogen ion concentrations within acceptable limits.

02     LIST THE MAJOR BUFFER SYSTEMS OF THE BODY

(1) HCO3/CO2 (bicarbonate/carbon dioxide), (2) HPO42―/H2PO4

― (phosphate), (3) Organic Phosphate Esters, and (4) Proteins. Proteins with side chains that contain more carboxyl terminal groups than amino terminal groups promote an acidic environment. Proteins with side chains that contain more amino terminal groups than carboxyl terminal groups promote an alkaline environment. Protein with side chains containing equal numbers of amino and carboxyl side groups are neutral, not affecting the pH.

03    DESCRIBE AN ANHYDRIDE AND GIVE AN EXAMPLE.

An anhydride is a molecule that is formed from another with the removal of a water molecule. Note the following reaction that is catalyzed by carbonic anhydrase.Carbonic acid     ——————>             water   +   carbon dioxide(H2CO3)         carbonic anhydrase            (H2O)             (CO2)

In this example, carbon dioxide is an anhydride. When the anhydride is added to water, then an acid is formed. The above reaction is reversible.

04    DISCUSS THE PRINCIPLES OF A TYPICAL BUFFER SYSTEM

Any buffer system must consist of a weak acid and its conjugate base. The use of the word "conjugate" infers a paired relationship. In chemistry, a conjugate base (such as  —NH3, CH3 — COO—, and HCO3) can accept proton(s). The corresponding conjugate acids (NH4, CH3 - COOH, and H2CO3) can donate proton(s). If you pour a strong acid into a buffer, the added H+ combines with the conjugate base causing it to diminish. The amount of conjugate acid increases causing a change in the concentration of conjugate acids and bases. There is a corresponding change in the pH.

05    EXPLAIN HOW THE BICARBONATE / CARBON DIOXIDE BUFFER SYSTEM WORKS

Examine first the exchange of carbon dioxide in the lungs. CO2 moves across the alveolar membrane and enters the plasma as dissolved CO2. Approximately 5% of CO2 is transported in it free form. Note that water makes up about 90% of the volume of blood. Erythrocytes contain carbonic anhydrase (CA) and the following reaction occurs:

        CO2 + H2O   ———>    H2CO3

In this reaction, CO2 is an acid anhydride and H2CO3 is a weak acid. Carbonic acid is unstable and will rapidly ionize to H+ (hydrogen ion) and HCO3

― (bicarbonate ion). H2CO3 exists in the plasma and cells at a concentration of .02% of the dissolved CO2. Note: If a chemical equation is contained between [     ], this infers the presence of a concentration. The above equation can be written like this:

CO2   +   H2O  ———>   [ H2CO3 ]   ———>     HCO3―   +   H+

Carbonic acid is quantitatively insignificant and is dropped from the equation. Because of the dibasic nature of H2CO3, it can dissociate to either CO2 or HCO3

―, dependent upon the pCO2, pO2, temperature, and pH.

The bicarbonate / carbon dioxide system is an open system in which the [CO2] can change to keep its buffers and pH constant. The plasma pH in a living person generally will not exceed 7.8 nor fall below 7.0 except in very ill situations. In the body, if an acid is added, increasing the H+, the conjugate base takes up the H+, removing it from the system. The conjugate acid remains relatively constant and the pH changes very little. The HCO3

―/ CO2 operates this way... as H+ are added, the [HCO3-]

diminishes as H2CO3 forms. Because of its lack of stability, H2CO3

dissociates into CO2 + H2O. The excess CO2 is blown off/exhaled in the lungs. There is little/insignificant change in the HCO3

― to pO2 ratio. Examine the following graph (see next page) for an illustrated example of how this buffer system functions under differing pH and [HCO3

―] conditions.       

In the clinical laboratory, it is relatively easy to test for the pH and pCO2. If you know these two values, then you can use the above graph to plot the [HCO3

―]. This graph shows only three isobars, but plots for all of the pCO2 isobars are commercially available. Consider the following examples for plotting the [HCO3

―]:a.   If the pH = 7.4, pCO2 = 40, then plotting the graph for [HCO3

―] = 24 (normal)b.   If the pH = 7.5, pCO2 = 20, then plotting the graph for [HCO3

―] = 14 (abnormal)

Now lets take one more look at how this buffer system works when the amount of CO2 is increasing. This increase means the body is acidifying.  Each CO2 molecule will react with water to form carbonic acid. The H2CO3 dissociates to HCO3

― and H+. The H+ that forms will be buffered out by the conjugate base forms of proteins and phosphates, allowing the buffer's conjugate base (HCO3

―) to increase, maintaining the physiological level of bicarbonate ion to CO2.  If the pCO2 should be low, then inconsequential amounts of H+ is formed. There will be a corresponding decrease in the [HCO3

―].  If the blood pH is being affected by the loss of the bicarbonate / carbon dioxide buffering system, the other buffering systems will dissociate to adjust the pH accordingly.

06    DEVELOP THE HENDERSON-HASSELBALCH EQUASION USING THE HCO3

― / CO2 BUFFER SYSTEM AS A MODEL AND EXPLAIN ITS PURPOSE

(1)   Begin with the equation, CO2  +  H2O    ———> HCO3― +  H+

                                                                                    [H+] [HCO3―]

(2)   Write the equation to this format:         Ka = ------------------------                                                                                            [CO2](3)   The next step is to rearrange the equation by solving for H+ and the Equation         rearranges to                                     [CO2]

                H+ = Ka -----------------                                    [HCO3

―]

(4)   More mathematical hocus-pocus is required. Take the negative logarithm of         both sides so that substitutions can be made. A method for substitution in pH        is needful and also pKa, which is a measurement of the strength of an acid. "K"        in the previous Equation is a dissociation constant for acidity and it needs to         substituted out. The equation now becomes:                                                            [CO2]            ― log [H+] = - log Ka ― log --------------                                                            [HCO3

―]

(5)    Now it is time to substitute in "pH" = - log [H+] and "pKa" for ―log Ka.            Now the Equation can now be rewritten as:                                               [HCO3

―]                pH = pKa + log ----------------                                                [CO2]

Oops!   Almost forgot, the signs were changed to obtain the above Equation, which is the Henderson-Hasselbalch Equation.

The purpose of the Henderson-Hasselbalch equation is to relate the pKa to the pH and compare the ratio of protons acceptors (base) to the proton donors (acid) of a conjugate base pair.

07    WHEN GIVEN DATA, THE STUDENT WILL BE ABLE TO USE THE HENDERSON-HASSELBALCH EQUASION

To use this equation, there is a small adjustment to be made. The part of the equasion

           [HCO3―]

+ log -------------- must be rewritten to introduce a conversion factor (α) that             [CO2]represents the solubility coefficient for carbon dioxide in normal plasma at 37 0C. The solubility coefficient for carbon dioxide = 0.03. The new and slightly modified equation now looks as follows:                             [HCO3

―]pH = pKa + log ----------------                               α • pCO2

It's okay to rearrange the equation to a new form: pH = pKa ―log [HCO3―]

― log [α • CO2]

a.   Lets make a few assumptions: (1) the pKa for this buffer system is 6.1,

      (2) pH = 7.2, (3) pCO2 = 37.5 mm Hg

b.   Substituting in: 7.2 = 6.1 - log [HCO3―] ― log [.03 × 37.5]

c.   α • pCO2 = dissolved CO2 concentration in this equation . . . .          0.03 × 37.5 = 1.13 mmol/L.

d.   Rearranging and solving for [HCO3―], the answer is 23.97 mmol/L

e.   To find the total CO2 concentration:

Total [CO2] = dissolved [CO2] + [HCO3―]

Total [CO2] = 1.13 mmol/L + 23.97 mmol/L

Total [CO2] = 25.1 mmol/L

08   EXPLAIN THE CONCEPT OF THE BUFFERING LINE OF BLOOD

This concept of the buffering line of blood arose by experimentation with the buffering properties of blood. It was found that pCO2 was altered in the presence of HCO3-. All this started out by trying to identify normal values. Examine the following graph for the buffering properties of blood.

    This graph shows changes occurring in the acid/base properties of blood as changes in pH and HCO3

― occurred. Notice the line sloping downward toward the right. This is the buffering line of blood and was obtained from the following data on the premise that normal pCO2 was at 40 mm Hg and a pH of 7.4:a.   If pCO2 increases to 60 mm Hg, then HCO3

― would increase by 2 meq/L―1       to 26 meq/L―1 and the pH would adjust to 7.3.

b.   If pCO2 increases to 80 mm Hg, then HCO3― would increase↑ by 4

meq/L―1        to 28 meq/L―1 and the pH would adjust to 7.2.c.   If pCO2 decreases to 30 mm Hg, then HCO3

― would decrease by 2 meq/L―1        to 22 meq/L―1 and the pH would adjust to 7.5.d.   If pCO2 decrease to 20 mm Hg, then HCO3

― would decrease by 2 meq/L―1        to 18 meq/L―1 and the pH would adjust to 7.6.

The intent of this "line" is to show how the bicarbonate system behaves in the presence of other buffers.  What happens in the blood to prevent big changes in pH and [H+] is the other buffering systems will capture the hydrogen ions formed by the creation of bicarbonate ion.  This buffering line of blood is a model or concept to predict the buffering behavior of bicarbonate ion when pCO2 is changed.

Let take a look at this buffering system when [HCO3―] increase to 29

meq/L―1, a 5 meq/L―1 increase.  This means that H2CO3 increased by 5 meq/L―1.  This H2CO3 will immediately ionize to form 5 meq/L―1 of H+ and HCO3

―.  This will result in the pH dropping down to about 7.2. What happened to the H+ that affected this degree of change in the ionization of H2CO3?  They were taken up into the phosphate, hemoglobin and plasma protein buffer system.

Look again at what happens when the pCO2 decreases to 20 mm Hg. [HCO3

―] decreases  to approximately 18 meq/L―1, a 6 meq/L―1 decrease.  The plasma level of [HCO3

―] has dropped and the pH adjusts itself upward to 7.6.

By now, it is seen that as the pCO2 is altered, the [HCO3―] varies according

to the amount of H+ generated. Other buffers will interact to pick up the excess hydrogen ions.  The above graph is characterized by a line that is generated with a distinct non-zero slope. It is this slope that is called the buffering line of blood.

09    DESCRIBE HOW HEMOGLOBIN FUNCTIONS AS A PART OF THE BUFFERING SYSTEM

Hemoglobin (Hgb) can fluctuate widely in a disease state. Heme can take up H+ to form deoxyhemoglobin, also called "reduced hemoglobin", a weaker acid than oxyhemoglobin. Because of its buffering properties, Hgb gives up CO2 and H+ in the lungs and in the tissues, O2 is released. If hemoglobin falls in a disease state, its buffering power is diminished. Hemoglobin determines it own buffering line of blood. assume that the Hgb level is 20 mg/dL, the pH = 7.4 (normal),and the pCO2 = 40 mm Hg (normal), the [HCO3

―] will = 24 meq/L―1.  If the pCO2 is increased to 80 mm Hg, the [HCO3

―] will rise to 31 meq/L―1 (up from 24 meq/L―1, a rise of 7 meq/L―1). This represents an increase of 7 meq/L―1 due to the ionization

of H2CO3 to HCO3― and H+. Hemoglobin buffers the blood, holding the pH

to 7.23. Refer to the following graph illustrating the buffering power of blood.

When contrasted against a normal Hgb of 15 g/dL to 20 g/dL, the following occurs as hemoglobin drops to 10 g/dL and buffering power is compromised as the pCO2 = 80 mm Hg: (1) there will be a lowering in the [HCO3

―] to 28 meq/L―1, (2) loss of control over pH as it adjusts downward to around 7.2.  Decrease in hemoglobin results in an overall drop in pH of blood.

10    DESCRIBE WHAT HAPPENS WHEN AN ACID (HCl OR CH3-COOH) OR BASE (NaOH) IS ADDED TO THE HCO3

― / CO2 BUFFERING SYSTEM

Assume that normal blood values are present: pCO2 = 40 mm Hg, pH = 7.4, [HCO3

―] = 24 meq/L―1. a.    If HCl is added to the system.... it will react with all of the buffers in the blood causing a decrease in the conjugate bases (proton acceptors). Because HCO3

― is the major buffer system, it will be affected the most. Using only a pCO2 system that is limited to 40 mm Hg and cannot fluctuate, lets assume that the amount of acid added will cause the pH to decrease to 7.2. This means that the [HCO3

―] will fall to 16 meq/L―1. This is a decrease of eight meq/L―1. b.    If NaOH is added to the system.... it will react with all of the buffers in the blood causing a decrease in the conjugate acids (proton donors). Again in this setting, HCO3

― is the major buffer system and HCO3―-

increases with addition of base causing a drop in the conjugate acids. If the amount of base added causes the pH to adjust upward to 7.6, the [HCO3

―] will increase to 40 meq/L―1. This is an increase of 16 meq/L―1.

Refer to the following graph:

11    BRIEFLY DISCUSS THE IMPORTANCE OF MAINTIANING AN ACID-BASE BALANCE IN THE BODY

There are a number of mechanisms in the body to control excesses in acids or alkali. If acidosis is present, there is an increase in acid or a decrease in alkali. If alkalosis is present, then there is a decrease in acid or an increase in alkali. When there is an imbalance, the body must compensate. The plasma pH must be controlled to maintain compatibility with life and the cause of the imbalance removed, re-establishing the normal acid-base balance. Note: The body can tolerate the pH extremes of 7.0 to 7.8.

In the body, CO2 is a normal metabolite that must be dealt with. It is the major acid, the body producing from 12,500 meq to 50,000 meq of CO2 daily. The average young adult male produces about 22,000 meq/day. CO2

is a volatile acid and most is blown off out of the lungs daily. Why is CO2 an acid?    Go back to Objective #05 in this section. If the lungs fail to perform this task, the other body systems cannot handle this increased load and acidosis will result.

12    EXPLAIN RESPIRATORY ACIDOSIS.

This is a base deficient disorder caused by hypoventilation of the lungs.  Remember that healthy alveoli are the key to good ventilation.  Hypoventilation can occur with diminished rate and/or depth of respiration. There is a resultant accumulation of CO2 (hypercapnia) in the body with a decrease in HCO3

―. An increase in conjugate acids in the blood can be demonstrated. If the pCO2 increases to 70 mmHg then

confusion occurs and the patient becomes obtund (insensitive to pain). Respiratory acidosis can develop rapidly if conditions are favorable. Causes for this disorder include:(1)    mechanical causes: airway obstruction, pneumonia, and pulmonary fibrosis         (Note: Lung diseases can impair gas exchange contribution to this disorder).(2)    neuromuscular disorders (3)    CNS disease (polio, encephalitis) causes respiratory depression(4)    emphysema, which gives rise to a chronic form of respiratory acidosis.(5)    decreased circulation as seen in cardiac disease.(6)    barbituates and similar drugs

The pulmonary system tries to compensate by an increased rate of respiration (due to the stimulation of the respiratory center by increasing pCO2 to breathe out and eliminate excess CO2.  Medical intervention recommends the infusion of alkali.  If the patient has a cardiopulmonary disorder this may not be a recommended procedure. Oxygen therapy, if employed, should be done so with caution.

The kidney's play a role in this disorder and this is what happens. There is an increase in the reabsorption of HCO3

― along with an increase in the secretion of H+ producing a more acid urine. Since there will be more CO2 in the kidney's, they simply generate additional HCO3

― for it. The kidney's will increase their absorption of Na+ in an attempt to eliminate more H+. Some of the excess H+ will be combined with NH3 to form NH4 , increasing NH3 formation. This renal response is called "compensatory metabolic alkalosis. As [HCO3

―] is increased, Cl―is correspondingly decreased to maintain the electric neutrality of the cells and tissues.

Laboratory testing will demonstrate:(1)    increased CO2

(2)    increased pCO2 and HCO3―

(3)    decreased pH(4)    decreased ratio of HCO3

― / H2CO3

13    EXPLAIN RESPIRATORY ALKALOSIS

This is a base excess disorder. It arises when there in a decrease in alveolar pCO2 by stimulated ventilation (hyperventilation).  Causes for this problem may be any of the following:(1)    CNS injury to the respiratory center(2)    early salicylate poisoning or other drugs(3)    pyrexia (causes hyperventilation)(4)    artificial ventilation. (5)    anxiety or hysteria causing hyperventilation (if this type of alkalosis occurs,         then the kidney will take the major role of trying to compensate.

There is a type of chronic respiratory alkalosis that can result from living or being at a high altitude. In respiratory alkalosis, there is a loss of CO2, but only a minimal change in the HCO3

― concentration. There is a relative excess of HCO3

―, due to the decrease in CO2. The plasma pH will increase.

Medical intervention may be as simple as breathing into a paper bag to rebreathe CO2 if symptoms are present. If not treated, the patient may present with enhanced neuromuscular excitement (which may culminate in tetany) or experience confusion and then become unconscious. Symptoms include light-headedness, paresthesias (numbness or prickling over body), circumoral (around the mouth) numbness, and tingling of the extremities. Some patients may have an enhancement of glycolysis which will result in increased lactate and pyruvate lowering the serum HCO3

―.

The kidney's respond to this disorder by (1) decreasing reabsorption of HCO3

― and Na+ and (2) increased reabsorption of Cl― to maintain an anion environment.  If hyperventilation stops, the situation reverses itself.

Laboratory testing will demonstrate:(1)    decreased Cl―, pCO2, and HCO3

―

(2)    increased ratio of HCO3― / H2CO3 and pH

14    EXPLAIN METABOLIC ACIDOSIS

This is a base deficient disorder (decrease in HCO3―) caused by a

metabolism that produces varying amounts of organic and non-organic acids. This condition is usually a chronic problem, not an acute one. Examples of organic and non-organic acids include:(1)    lactic acid (if this is the cause, then it is known as lactate acidosis)(2)    acetoacetic acid and β-hydroxybutyric acid (if this is the cause, then it is         known as diabetic keto-acidosis)(3)    H+ and SO4

2― (equivalents of sulfuric acid) and hydrolysis of phosphate         esters (equivalents of phosphoric acid). If these types of acids accumulate, the        pH will decrease. causing metabolic acidosis. Causes of metabolic acidosis include:    a.     diarrhea:  HCO3

― is lost in the stool causing a change in the             HCO3

― / CO2 ratio, creating a relative increase in H+.     b.     administering ammonium chloride (NH4Cl) to create an acid urine to             prevent kidney stone formation. NH4Cl is converted to urea and HCl is             synthesized. The same thing occurs if arginine•HCl or lysine•HCl is             administered.    c.     ingestion of methanol, salicylates, or ethylene glycol results in the production

            of strong acids.    d.     increases in organic acids other than H2CO3.   In uncontrolled diabetes,             acetoacetic acid and β-hydroxybutyric acid accumulate.     e.     strenuous exercise and systemic infections both result in the increased             production of lactic acid.    f.     medications (as sulfonamide or acetazolamide [anticonvulsant]) inhibits             carbonic anhydrase disrupting the formation of HCO3

―.    g.     acute renal failure results in the loss of the ability to excrete acids but there is            an increase in anions, creating a higher anion gap.    h.     diabetes mellitus or starvation results in the production of ketoacids.

In metabolic acidosis, sodium lactate is administered as a remedy. It is converted to NaHCO3 as follows:Na+  +  lactate  +  3 O2  <———>  Na+  +  HCO3

―  +  2 CO2  +  2 H2O

The body responds by hyperventilating to overcome acidosis. CO2 is exhaled through the lungs and the pCO2 decreases. A physiological overall effect occurs in which there is an increase of HCO3

― to CO2 ratio. The plasma pH will return toward normal. This process is called compensatory respiratory acidosis, but the response is incomplete and the mild acidosis usually remains.

The kidney's respond to the acidosis by retaining Na+ and excreting H+ and increasing the secretion of NH3 and forming NH4.  If the patient has an acute renal failure, the kidney cannot excrete H+. This inability to secrete means that there will be no [HCO3

―] to replace lost to buffering the acids produced by the body.

Laboratory testing will demonstrate:(1)   decreased HCO3

―, pCO2, and H2CO3, CO2, pH, K+, and acidic urine         (resulting in an increased anion gap)(2)   increased Cl― (occurs with diarrhea and related HCO3

― loss)

15    EXPLAIN METABOLIC ALKALOSIS

This is a base excess disorder. This may be the result of any of the following: a.    an excessive intake of alkali as bicarbonate of soda,b.    an abnormal loss of acid through prolonged vomiting due to the loss of HCl.        The loss of chloride ion is called hypochloremic alkalosis and the kidney's will          reabsorb more HCO3

―to compensate for Na+ reabsorption.c.    prolonged administration of certain diuretics such as Diamon, Lasix,

and        Dichlorphenamide.  These medications increase Na+ loss and also increase        the secretion and loss of K+.d.    nasogastric suction (removal and loss of H+ rich fluid).e.    If there is a loss of fluid from the body that is low or void in bicarbonate ion, the       remaining body fluid will be low resulting in a transitory concentration

       of [HCO3―].

This disorder usually develops as a chronic disorder, very seldom as an acute disorder. If the patient develops severe symptoms, then the patient will become apathetic with stupor and confusion. If there is too much calcium ion loss, then tetany may result.

Medical intervention is seldom required for mild to moderate metabolic alkalosis. If the problem is gastric induced, then infusion of saline is usually sufficient for correction of the alkalosis.

The body tries to compensate with respiratory center depression and subsequent retention of CO2. There should be a corresponding increase in HCO3

―and normalization of the HCO3― / H2CO3 ratio and the pH will

approach closer normal.The kidney's will increase the rate of elimination of HCO3

― in the urine by decreasing it reabsorption. There will be a corresponding decrease in the formation of NH3 and decrease in the exchange of Na+ and H+.

Laboratory testing will demonstrate an increased HCO3―, CO2, p CO2, and

pH.

16    EXPLAIN AND/OR ILLUSTRATE ION EXCHANGE BY THE RENAL TUBULE IN THE KIDNEY’S ROLE TO REGULATE ACID-BASE BALANCE THROUGH INORGANIC ACID EXCHANGE

   

17    EXPLAIN AND / OR ILLUSTRATE ION EXCHANGE BY THE RENAL TUBULE IN THE KIDNEY'S ROLE TO REGULATE ACID - BASE BALANCE THROUGH THE REABSORPTION OF BICARBONATE ION

     18    EXPLAIN AND/OR ILLUSTRATE ION EXCHANGE BY THE RENAL TUBULE IN THE KIDNEY’S ROLE TO REGULATE ACID-BASE BALANCE THROUGH THE EXCRETION OF AMMONIA

    19    CREATE A TABLE LISTING THE CHANGES THAT OCCURS IN BLOOD GASS PARAMTERS IN ACIDOSIS STATES AND USE THIS INFORMATION TO INTERPRET ANALYTICAL DATA

20    DEFINE THE FOLLOWING TERMS:

a.    Compensated: T he acid-base imbalance has been in effect. Then internal         mechanisms come into action and readjust the acid-base balance back to the        normal pH.b.    Uncompensated:  The acid-base balance is "out-of-whack". There is some        abnormality present that is preventing compensation and adjsutment back to

       the normal pH.c.    Partially compensated:  The acid-base adjustments are incomplete or partial.        There may be some factors that limit full recovery.d.    Total Plasma CO2:  The sum of the bicarbonate and dissolved CO2. This         value will be higher than the bicarbonate concentation.

21    EXPLAIN THE CONCEPT OF BASE EXCESS

This is a mathematical calculation to access the acid-base status of the patient. If the value calculated is for a positive base excess, then the excess is HCO3

―. If the value is for a negative base excess, there there is a deficit of HCO3

―. Base excess (positive) is defined as the amount of acid needed to be added to blood to titrate it to a pH of 7.4 with pCO2 2― 40 mm Hg at 37OC. a.    If blood is acidic, alkali will have to be added and the base excess calculation         would have been negative.b.    If blood is basic, acid will have to be added and the base excess calculation         would have been positive.

Review the following graph  to further define "base excess".

   

22    DISCUSS THE SIGNIFICANCE OF SODIUM AND CHLORIDE IONS IN ACID-BASE BALANCE

This concept is best approached by examining the concept of the anion gap and how it can help evaluate acid-base disorders. The anion gap concept requires having the lab data for sodium, chloride, postassium, and bicarbonate ions at hand. The anion gap is calculated as follows:anion gap = ( Na+  +  K+ ) ― ( Cl―  +  HCO3

― )

This is a lab measurement that estimates the net number of anions in serum. The net number of anions are not directly measured. The concept of the anion gap is a tool to alert the physician to a potential disorder that may be developing and might alter the electrolyte balance. The normal range for the anion gap is from 8 to 18 meq/L. The most common application of the anion gap is to diagnose metabolic acidosis. If metabolic acidosis is the problem, then it will be due to one of the following causes: (1) there is an increase in hydrogen and chloride ions or (2) there is a decrease in sodium and bicarbonate ions.

An increase in the anion gap may be due to:a.    hypocalcemia:  a loss of calcium ions (these are cations) and it will be

       characterized by a slight increase in unmeasured anions. (NOTE: The same         effect occurs with hypomagnesemia, hypokalemia, and other cations).b.   salicylate poisoning:  results in an increase in acidic metabolites.c.    renal failure:  there is an increase in sulfate (SO4

2― ) and phosphate ions        (PO4

―).  The kidney can no longer eliminate these in the urine.d.    diabetes mellitus:  an increased number of ketones (acetoacetic acid).

A decrease in the anion gap may be due to:a.    hypoalbuminemia:  Albumin is a negatively charged protein.b.    polyconial gamma globulins:  These type of globulins are proteins with a        net positive charge (cation effect). Their increase would have the effect of        diluting out (decreasing) the anions.c.    increase in the amount of plasma water:   Increases in water content has        the effect of diluting out the anions and cations.d.    hypercalcemia:  An increase in calcium ions (cations) that add more positive        charges to the electrolyte environment. This has the effect of decreasing the net         number of negative charges. (NOTE: The same effect occurs with

        hypermagnesemia, hyperkalemia, and other cations).

The anion gap is a non-specific test producing the general diagnostic data as does the ESR and CRP.

                                                                NOTEElectrolyte analysis is generally confined to measuring Na+, K+, Cl―, and HCO3

―. These four ions are referred to as the measured ions. In normal, healthy individuals, the anion gap will be positive. Look at a sample problem for a normal individual:( Na+ [142 meq/L]  +  K+ [4 meq/L])  ―  ( Cl― [103 meq/L]  +  HCO3

― [27 meq/L)                       140    +    4    =    144     minus        103    +    27   =    130         Answer   =   +14 meq/L

Electrolytes NOT measured for electrolyte studies are Ca2+ [5 meq/L], Mg2+ [2 meq/L], other cations [1 meq/L], HPO4

2- [2 meq/L], SO42- [1 meq/L],

Organic anions [5 meq/L], and proteins with net negative charge [16 meq/L].  If you calculate the anion gap with these values....5 + 2 + 1  =   2 + 1 + 5 + 16   =   8  ―  24   =  ―16meq/L  (cations)          (anions)

If you calculate the anion gap using both measured and non-measured ions, the following is obtained:142 + 4 + 5 + 2 + 1  ( ― )  103 + 27 + 2 + 1 + 5 + 16                            154  ( ―)   154   =   0The electrolyte system is in balance, with an equal number of positive and negative ions.Note the normal range for the anion gap is 8 to 18 mEq/L.

23    DESCRIBE WHY HEMOGLOBIN CAN FUNCTION AS A BUFFER

Hemoglobin contains ionizing groups. For example, there are 38 histidine molecules per tetramer.  There are also four N-terminal amino acids with buffering potential.  Hemoglobin has the ability to "give-up" oxygen and "take-up" hydrogen ion.  This ability to do this is called the isohydric carriage of carbon dioxide. The following reaction describes the phenomenon:

Hgb • O2  +  CO2  +  H2O  <———>   Hgb • H  +  HCO3―  +  O2

Oxyhemoglobin is a weak acid. Organic acids and H+ (called the acidity of metabolism) are present in the tissues causing a lowered pH with favors the release of oxygen from the hemoglobin molecule.

24    ILLUSTRATE THE BUFFERING ACTION OF HEMOGLOBIN IN THE LUNGS

   

25    ILLUSTRATE THE BUFFERING ACTION OF HEMOGLOBIN IN THE TISSUES

   

26    EXPLAIN HOW HEMOGLOBIN CAN TRANSPORT CARBON DIOXIDE

Carbon dioxide can combine with the amino groups of proteins and hemoglobin has exposed amino groups. The N-terminal group of valine is

the site of the CO2 bond. Note the following reaction:

Hgb - VAL -NH2  +  CO2  <———>   Hgb - VAL - NH - COO―  +  H+

This is a rapid and readily reversible reaction.  There is no evidence that an enzyme is required for this reaction.  The combination of carbon dioxide to hemoglobin forms carbaminohemoglobin and about 13% of CO2

is transported in this form.  Up to 9% of CO2 is transported in its free or dissolved form. There are textbook references that limit this amount to 5%.  Approximately 78% of CO2 is transported in the from of HCO3

―.

27     EXPLAIN HOW CARBON MONOXIDE EXERTS ITS TOXIC EFFECTS

Carbon monoxide (CO) is an odorless gas that binds to the sixth position of the heme iron (just as does oxygen). If CO and oxygen (O2) are present at the heme structure at the same time, CO will preferentially bind to the heme iron. CO acts as an allosteric activator and once it is bound to hemoglobin, O2 will neither bind to nor be released from the heme iron because the heme structure is shifted to the R-state or forms a 6-coordinate iron so that the iron molecule is stabilized in such a way that O2 cannot be released or taken up.  In the deoxygenated state, the iron molecule is in the 5-coordinate form or T-state and can readily take up an oxygen molecule to initiate the biding of oxygen to the remaining heme subunits.  Normal blood levels of carboxyhemoglobin in nonsmokers will range up to 1.5% of the total hemoglobin whereas in smokers, it will range between 4 - 5%. In ‘heavy’ smokers, the range may be up to 8%. Breathing air that contains 1.0% CO will be fatal in 7 minutes.  Automobile exhaust contains about 7% CO. A person who has a blood concentration of 20% carboxyhemoglobin will experience mild to throbbing headaches.  When the concentration goes up to 30 to 50%, then other symptoms of irritability, confusion, dizziness, visual disturbances, nausea, vomiting, and/or fainting will appear.  When the blood concentration is more than 50%, the victim will become comatose, experience convulsions, respiratory failure, and death will then onset.  The most characteristic sign of CO poisoning are a cherry red (flush) color in the skin, nail beds, and buccal membranes.  Treatment consists of therapy with 90% oxygen and 5% carbon dioxide which will eliminate the CO within 90 minutes.  In severe cases of CO toxicity, blood transfusions may be a therapeutic strategy.

                                For Your Information.1.   Types of electrolytes:         A.     Major Electrolytes are Na+, K+, Cl―, and HCO3

―.        B.     Anions are Cl―, HCO3

―, HPO42―, and HPO4

―.        C.     Cations are Na+, K+, Ca++, Mg++, Fe++, Fe+++, Cu+, Cu++.2.   Chloride is the major extracellular anion affecting water distribution and        osmotic pressure.  Its normal extracellular concentration ranges from

98 to       106 mMol/L.3.   Potassium is the major intracellular cation, its intracellular concentration about      23 times that of the outside concentration. Normal extracellular concentration       ranges from 3.5 to 5.5 mMol/L.  Each 0.5% hemolysis of RBC’s will increase       the extracellular K+ concentration by about 0.5 mM/L. Item of interest . . . .       DO NOT over centrifuge blood.  Extended centrifugation may cause the       gravitational force exerted by the spinning centrifuge to cause K+ to be extruded       from the cell, increasing extracellular concentration.4.   Sodium is the major extracellular cation, having a major role in maintenance of        the body’s hydration and osmotic pressure.  Its extracellular concentration is        controlled by the renin-angiotensin-aldosterone mechanism.  Normal Na+        concentration ranges from 135 to 145 mMol/L.5.   Magnesium is the second most abundant intracellular cation.  Its normal        concentration ranges from 1.6 to 2.1 mEq/L