Quantitative Acid-Base Balance Charles L. Webber, Jr., Ph.D.

Learning Objectives: You should be able to:

• Explain how hydrogen ions interact with enzymatic proteins, altering reaction rates within intracellular spaces.
• Give the general form of the Henderson-Hasselbalch equation and explain why a buffer is most effective when pH equals pK.
• Give the specific form of the Henderson-Hasselbalch equation for the bicarbonate buffer system, identifying the respiratory and renal variables.
• Compare and contrast the phosphate/protein buffer systems (closed) and the bicarbonate buffer system (open) in regulating extracellular [H+].

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Introduction to Acid-Base Chemistry

1. Perspective (Fig. 13)
   • acid-base balance refers to the regulation of [H+] in body fluids
   • this is a sizable task since H+ and CO2 are continually being produced by metabolic reactions

     Fig. 13

    

2. Concept of pH (Fig. 14)
   • wide ranges in [H+] are conveniently represented by the pH scale]
     • þ pH = -log[H+]
     • þ [H+] = 10-pH
   • non-living, in vitro systems
     • þ [H+] can vary from 1 mol/L (pH = 0) to 10-14 mol/L (pH = 14)
     • þ neutral point at pH = 7.00 separates acidic from alkaline solutions
   • living, in vivo systems
     • þ [H+] can vary from 10-6 mol/L (pH = 6) to 10-8 mol/L (pH = 8)
     • þ normal point at pH = 7.40 separates acidotic from alkalotic fluids
   • quick conversions from pH to [H+]
     • þ from pH = 7.28 to 7.45, [H+] = 80 - (pH - 7) * 100
     • þ over a wider range from pH = 7.0 to 8.0 the expedient looses accuracy

     Fig. 14

    

3. Importance of [H+] in Body Fluids (Fig. 15)
   • H+ are highly reactive cations that change the charge distribution on proteins þ H+ + COO- þ COOH and H+ + NH2 þ NH3+
   • altered charge distributions lead to protein conformational changes and modified reaction rates þ protein enzymes have optimal pH ranges
   • it is the regulation of intracellular [H+] which is of crucial importance þ protein enzymes are located mostly intracellularly
   • it is the extracellular [H+] in the plasma that is operated on by the kidneys and lungs þ the extracellular fluid bathes intracellular spaces
   • the regulation of extracellular [H+] helps regulate intracellular [H+]
     • þ proton transfer from intracellular (IC) space to interstitial (IS) space is slow
     • þ proton transfer from IS space to plasma (P) space is fast
     • þ proton transfer from P space to urine is slow
     • þ proton transfer (in form of CO2) from P space to alveolar air is fast

       Fig. 15

    

4. Three Major Systems Responsible for Maintaining Arterial Plasma [H+]
   • chemical buffering (detailed below)
     • þ phosphate buffer system
     • þ protein buffers system
     • þ bicarbonate buffer system
   • renal system (slowly responding system)
     • þ kidney excretes 50 mmol H+ per day as H+, NH4+ and H2PO4- (urine is acidic)
     • þ kidney reabsorbs 5500 mmol HCO3- per day ([HCO3-] is at renal plasma threshold)
   • respiratory system (rapidly responding system)
     • þ lung ventilates off 13000 mmol CO2 per day (potential H+ via CO2 hydration reaction)
     • þ this is 150 times the capacity of the kidney (13000 mmol - 5500 mmol) / 50 mmol

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Acids, Bases and Buffers

1. Henderson-Hasselbalch Equation (Fig. 16)
   • Brþnsted-Lowry concept of acids and bases
     • þ acid (HA): proton donor
     • þ base (A-): proton acceptor
   • dissociation constant K
     • þ small K: weak acid (weak proton donor)
     • þ large K: strong acid (strong proton donor)
     • þ pK = - log(K)
   • computation of pH
     • þ pH = pK + log(base/acid)
     • þ any  (or ) in base [A-] can be balanced by a corresponding  (or ) in acid [HA]
     • þ it is the base/acid ratio that is critical, not the absolute concentrations of acid or base

       Fig. 16

    

2. Principles of Titration (Fig. 17)
   • titration of 10 mmol acetic acid in 1 L water at pH = 4.70 (point A)
     • þ acetic acid is a weak acid with a pK = 4.70
     • þ CH3COOH þ H+ + CH3COO-
   • since pH = pK, half the acid is dissociated, half is undissociated
     • þ [CH3COOH] = 5 mM
     • þ [CH3COO-] = 5 mM
     • þ [H+] = 0.02 mM
   • if 2 mmol H+ is added, pH falls to 4.33 (point B), not 2.69 as expected
     • þ pH = -log(0.02 mM + 2 mM) = 2.69 does not happen
   • mass action shift to the left buffers the excess H+
     • þ  [CH3COOH] > 5 mM
     • þ  [CH3COO-] < 5 mM
     • þ  [H+] to only 0.05 mM, not 2.02 mM if no buffering had occurred
     • þ this works only because acetic acid is a weak acid

       Fig. 17

    

3. Principles of Buffer Action
   • buffer pair
     • þ buffer system is most resistant to changes in [H+] when operating at pH = pK
   • buffer range
     • þ buffer system still has buffering ability one pH unit on either side of the pK
   • buffer power
     • þ buffer system strength is directly related to the concentration of the buffer pair components

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Non-Bicarbonate Buffer Systems

1. Phosphate Buffer System (weak)
   • H2PO4- þ H+ + HPO4- -
     • þ dihydrogen phosphate: acid (proton donor)
     • þ monohydrogen phosphate: base (proton acceptor)
     • þ pK = 6.8 (within physiological range)
   • weak system in the plasma since the buffer power is low (low concentration of components)
   • stronger system in the kidney where the environment is more acidic and [phosphate] is higher

    

2. Protein Buffer System (strong)
   • a large fraction of all chemical buffering power is attributed to intracellular proteins
   • hemoglobin is the most important blood borne protein buffer
     • þ hemoglobin concentration is the highest for any protein in blood
     • þ hemoglobin is classified as extracellular despite its intracellular location within red blood cells
   • the imidazole group of histidine residues is most significant group of proteins for buffering H+
     • þ pK of different imidazole groups vary from 5.3 to 8.3
     • þ many imidazole groups have pKs with the physiological range (pKs = 7.40 ñ 1.0)

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The Bicarbonate Buffer System

1. CO2 Hydration Reaction

    

   • three variables (analogous to the Henderson-Hasselbalch derivation)
     • þ CO2: acid (proton donor)
     • þ HCO3-: base (proton acceptor)
     • þ H+: free proton responsible for setting the pH
   • two constants
     • þ 0.03 constant converts PCO2 from mm Hg to mM CO2 (not to be confused with CO2 solubility of 0.06 mL CO2/dLblood per mm Hg)
     • þ pK = 6.1 (apparently out of physiological range)
   • pH = 6.1 + log(kidney/lung)

    

2. Importance of the Bicarbonate Buffer System
   • the bicarbonate buffer system is an open system
     • þ CO2 is directly linked to the environment via the lungs (ventilation)
     • þ H+ is directly linked to the environment via the kidneys (urination)
   • mass action shifts in such an open system can continue throughout life in one direction
     • þ CO2 and H+ are continuously removed from the body at rates that exactly match production
     • þ normally [CO2] and [H+] cannot build up
     • þ pathologically [CO2] and [H+] can build up (respiratory or metabolic acidosis)
   • the open nature of the bicarbonate buffer system more than compensates for its low pK

     Fig. 18

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