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Introduction to Acid-Base Physiology

I. Definitions

A. Acid - can donate a hydrogen ion

B. Base - can accept a hydrogen ion

C. Strong acid - completely or almost completely dissociates into a hydrogen ion and its conjugate base in aqueous solution; a weak acid is only slightly ionized in aqueous solution. A strong acid usually has a weak conjugate base and a weak acid usually has a strong conjugate base. Do not confuse the strength of an acid or base with its concentration.

D. Buffer - mixture of substances in aqueous solution, usually a combination of a weak acid and its conjugate base, that can resist changes in [H+] when strong acids or bases are added.

II. Quantification of Acidity

A. pH = -log [H+]

B. pH of 7.10 is [H+] of 80 nM/L

C. pH of 7.30 is [H+] of 50 nM/L

    pH of 7.40 is [H+] of 40 nM/L

    pH of 7.50 is [H+] of 32 nM/L

    pH of 7.70 is [H+] of 20 nM/L

D. Note that as [H+] increases, pH decreases. An increase of 0.3 pH units indicates that [H+] is cut in half; a decrease of 0.3 pH units indicates that [H+] is doubled.

E. By convention:

1. Arterial pH< 7.35 is acidemia

2. Arterial pH >7.45 is alkalemia

3. Acidosis and alkalosis are underlying disorders

III. Sources of Acids in the Body

A. CO2 + H2srlarrow H2CO3 ( a volatile acid)
    Lungs remove about 15,000 - 25,000 mmol of carbon dioxide per day.

B. Fixed acids produced as a result of metabolism. About 70 mEq of H+'s is removed mainly by the kidneys (a minor portion is removed by the G.I. tract) each day.       Fixed acids normally represent only about 0.2% of the total body acid production. May be much higher - e.g. in diabetic ketoacidosis.

IV. Buffer systems of the human body - Isohydric principle - all buffer pairs in a homogenous solution are in equilibrium with the same hydrogen ions.

A. Bicarbonate:


1. Open system: CO2 removed in lungs

2. CO2 dissolved in plasma (in mM/L) is equal to 0.03 x Pco2

3. CO2 + H2Osrlarrow​ H2CO3; but in equilibrium, so both dissolved CO2 and H2CO3 are considered part of [HA]. The ratio of dissolved CO2 to carbonic acid is about 1,000 to 1.

4. Note that "total CO2" = dissolved CO2 + H2CO3 + HCO-3

5. Henderson-Hasselbalch equation:

              pH = pK' + log hco3

    But the H2CO-3 is negligible,

              so:pH = pK' + loghco2

    The pK' of this system at physiological pH's and at 38°C is 6.1.

    Therefore, at pHa of 7.40 and PaCO2 of 40 torr: 

             7.40 = 6.1 + log hco1

6. The [HCO-3] is therefore normally about 24 mM/L because the log of 20 is about 1.3.

7. The buffer value of plasma in the presence of Hb is 4-5 times that of plasma separated from erythrocytes. (Buffer value = H+'s in mEq/L that can be added to or removed from a solution with a resultant change of one pH unit).

B. Phosphate:

          h2po4srlarrowhpo4  ; pK is about 6.8. Other organic phosphates have pK's near 7.0: Glucose - 1-P, ATP, etc.

C. Proteins

     Imidazole groups of histidine residues have pK's ranging from 5.5 to 8.5

1. Main protein: Hb. DeoxyHb is a weaker acid than is oxyHb. This allows CO2 loading in the tissues as deoxyHb can accept H+’s from the dissociation of H2CO3 and from carbamino compounds (see Levitzky, Chapter 7); and unloading in the lungs.

2. Other plasma proteins may also act as buffers.

D. Buffers of the interstitial fluid:

1. Mainly bicarbonate; some phosphate.

2. Note that the volume of interstitial fluid is much larger than that of the plasma, so the ISF may play an important role in buffering.

E. Bone:

       Much calcium and phosphate in bone. Chronic acidosis may lead to bone demineralization.

F. Intracellular buffering: Intracellular proteins and organic phosphates. Note that Hb is an intracellular buffer.

V. Acidosis and Alkalosis: use of the Davenport Diagram (Levitzky Fig. 8-1)

A. Respiratory acidosis:

             hcopaco4rarrowdarrow pHa

                     Causes: Impairment of vdotA (Levitzky Table 8-1)

B. Respiratory Alkalosis:


                   Causes: Hyperventilation (Levitzky Table 8-2)

C. Metabolic Acidosis

              hcopaco2rarrowdarrow pHa

              Causes: ingestion, infusion or production of fixed acids; decreased renal H+excretion; loss of HCO-3 etc. e.g. diarrhea, diabetic ketoacidosis, lactic acidosis (Levitzky Table 8-3)

D. Metabolic Alkalosis

             hcopacorarrowuparrow pHa

               Causes: Loss of fixed acids; ingestion, infusion, or excessive reabsorption of bases. e.g. vomiting; excess antacids (Levitzky Table 8-4)

VI. Compensatory Mechanisms

A. Respiratory: H+’s stimulate arterial chemoreceptors; compensation takes minutes

B. Renal: excretion of fixed acids; reabsorption of base; compensation takes hours or days. Mechanisms:

1. Secretion of H+ by tubular cells into lumen (inverse relationship with K+secretion) and reabsorption of filtered HCO-3. 90% of all HCO-3 is "reabsorbed" either directly or by carbonic acid dissociation in the proximal tubule.

2. Phosphate

3. Ammonia

4. In alkalosis the kidney decreases H+ secretion and decreases HCO-3 reabsorption. Kidney tends to reabsorb almost all filtered HCO-3 until [HCO-3]p reaches about 27-28 mEq/L.

C. Summary of Acid-Base Disorders (Levitzky Table 8-5):

               pHa = a constant +pha

VII. Base excess (or Deficit): The number of mEq of acid or base needed to titrate one liter of blood to a pHa of 7.40 at 37°C, if the Paco2 were held constant at 40 torr.

VIII. Anion gap = [Na+] - ([Cl-] + [HCO-3]); should be 12+ 4 mEq/L High anion gap (>16 mEq/L ) in metabolic acidosis suggests presence of anions that are not usually measured. Causes of high Anion Gap: (Levitzky Table 8-6)

IX. Disorders that can cause of tissue hypoxia (Levitzky Table 8-7)


Copyright 2000 M. G. LEVITZKY

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