3. pH of Body Fluids
• normal range of arterial pH is 7.37 to 7.42
• Acidemia: arterial pH is less than 7.37
• Alkalemia: Arterial pH is greater than 7.42
• Conc. of H⁺ is responsible for acid-base status
• In arterial blood
– the H+
concentration is 40 × 10-9
equivalents per liter (or 40
nEq/L),
• Mechanisms of maintaining normal pH
– buffering of H+
in both ECF &ICF
– respiratory compensation
– renal compensation
5. Acid Production in the Body
• Types of Acids
– Volatile acid
• CO₂
• Produced from aerobic metabolism
• CO₂ combines with H₂O to form the weak acid H₂CO₃⁻, which
dissociates into H⁺ and HCO₃⁻
• Carbonic anhydrase catalyzes the reversible reaction b/t CO₂ and
H₂O
– Non-volatile acid
• aka fixed acids
• Sulfuric acid, phosphoric acid
• ~ 40 -60 mmoles/day is produced
• Ketoacids, lactic acid, β-hydroxybutyric acid, glycolic acid, oxalic
acid & salicylic acid: over produced in disease or ingested
6. Buffers
• Prevents change in pH when H⁺ ions are
added to or removed from a sol.
• Most effective within 1.0 pH unit of the pK (–ve
logarithm of the [H⁺] at which ½ of the acid molecules are dissociated
and are undissociated)of the buffer
– i.e; in the linear portion of the titration curve
• Types
– Extracellular
– Intracellular
7. • Extracellular buffers
– Mostly HCO3-, which is produced from CO2 and H2O
• pK of CO2/HCO3- buffer pair is 6.1
– Phoshate
• Minor buffer
• pK of H2PO4-/HPO4- is 6.8
• Most important as a urinary buffer
– Excretion of H+ is called titratable acid
• Intracellular buffers
– Organic phosphates
• AMP, ADP, ATP, DPG
– Proteins
– Hemoglobin: major buffer
• Deoxyhemoglobin is better buffer than oxyhemoglobin
8. Henderson-Hasselbalch Equation
• Used to calculate pH
– where
• pH = -log10 [H+
] (pH units)
• pK = –ve logarithm of the [H⁺] at which ½ of the acid
molecules are dissociated and are undissociated
• [A-
] = Concentration of base form of buffer (mEq/L); is the
H acceptor
• [HA] = Concentration of acid form of buffer (mEq/L); is the
H donor
• When the conc. of A- and Ha are equal, the pH of
the sol. = the pH of the buffer
9. Ex of cal
• SAMPLE PROBLEM
– The pK of the HPO4
-2
/H2PO4
-
buffer pair is 6.8. Answer
two questions about this buffer: (1) At a blood pH of
7.4, what are the relative concentrations of the acid
form and the base form of this buffer pair? (2) At
what pH would the concentrations of the acid and
base forms be equal?
• SOLUTION
– The acid form of this buffer is H2PO4
-
, and the base
form is HPO4
-2
.The relative concentrations of the acid
and base forms are set by the pH of the solution and
the characteristic pK.
10. • Answering the first question: The relative
concentrations of acid and base forms at pH 7.4
are calculated with the Henderson-Hasselbalch
equation. (Hint: In the last step of the solution,
take the antilog of both sides of the equation!)
•
•
Therefore, at pH 7.4, the concentration of the
base form (HPO4
-2
) is approximately fourfold
that of the acid form (H2PO4
-
).
11. • Answering the second question: The pH at which there would be
equal concentrations of the acid and base forms can also be
calculated from the Henderson-Hasselbalch equation. When the
acid and base forms are in equal concentrations, HPO4
-2
/H2PO4
-
= 1.0.
•
•
The calculated pH equals the pK of the buffer. This important
calculation demonstrates that when the pH of a solution equals the
pK, the concentrations of the acid and base forms of the buffer are
equal. As discussed later in the chapter, a buffer functions best
when the pH of the solution is equal (or nearly equal) to the pK,
precisely because the concentrations of the acid and base forms are
equal, or nearly equal.
12. Titration Curves
• Describes how the pH of a buffered sol. changes as H ions are
added to it or removed from it
• As H⁺ ions are added to the sol., the HA form is produced
• As H⁺ ions are removed, the A⁻ form is produced
• A buffer is most effective in the linear proportion of the
titration curve, where the addition of removal of H causes
little change to pH
– Most effective physiologic buffer will have a pK with 1.0 pH unit of 7.4
(7.4 ±1.0)
– Outside of the effective range addition of removal of H⁺ changes sol.
drastically
• Based on Henderson-Hasselbalch Equation, when the pH of
the sol. = pK, the conc. of HA and A are equal
13. Figure 7-2 Titration curve of a weak acid (HA) and its conjugate base
(A-
). When pH equals pK, there are equal concentrations of HA and A-
.
15. Figure 7-3 Comparison of titration curves for H2PO4
-
/HPO4
-2
and
CO2/HCO3
-
. ECF, Extracellular fluid
16. Figure 7-4 Acid-base map. The relationships shown are between
arterial blood PCO2, [HCO3
-
], and pH.
17. Renal Acid-Base
• Reabsorption of HCO₃⁻
• Figure 7-5 Mechanism for reabsorption of filtered HCO3
-
in a
cell of the proximal tubule. ATP, Adenosine triphosphate.
18. Regulation of filtered HCO₃⁻
• Filtered load
– Increases in the filtered load leads to increased HCO₃⁻ reabsorption.
– If plasma [HCO3} becomes high (metabolic alkalosis), then filtration
will exceed reabsorption and excretion will occur
• PCO₂
– ↑ PCO₂ → ↑ HCO₃⁻ reabsorption; d/t ↑ICF [H⁺] for secretion
– ↓ PCO₂ → ↓ HCO₃⁻ reabsorption; d/t ↓ICF [H⁺] for secretion
• ECF volume
– Expansion → ↓ HCO₃⁻ reabsorption
– Contraction → ↑ HCO₃⁻ reabsorption (contraction alkalosis)
• Angiotensin II
– Stimulates Na⁺-H⁺ exchange and thus increases HCO₃⁻ reabsorption
– Contributes to contraction alkalosis secondary to ECF vol. expansion
19. Excretion of Fixed H⁺
• Excretion of H⁺ as titratable acid (H₂PO₄⁻)
• Figure 7-6 Mechanism for excretion of H+
as titratable acid.
ATP, Adenosine triphosphate
20. • Excretion of H⁺ as NH4 (ammonium)
• Figure 7-8 Mechanism of excretion of H+
as NH4
+
. In the proximal tubule, NH3 is produced
from glutamine in the renal cells, and NH4
+
is secreted by the Na+
-H+
exchanger. In the
collecting ducts, NH3 diffuses from the medullary interstitium into the lumen, combines with
secreted H+
in the lumen, and is excreted as NH4
+
. ATP, Adenosine triphosphate.
22. Figure 7-10 Values for simple acid-base disorders superimposed
on acid-base map.
23. • Metabolic Acidosis
– Over-production or injestion of fixed acid or loss
of base (↓HCO₃⁻) produces and increase in
arterial [H⁺] (acidemia)
– Primary disturbance = ↓[ HCO₃⁻]
– HCO₃⁻ is used to buffer the extra acid
– Compensation
• Respiratory compensation = hyperventilation
(kussmaul breathing – deep rapid respiration, common
in type 1 diabetics d/t keto acids)
• Renal compensation =
– ↑ excretion of H⁺ as titratable acid & NH₄
– ↑”new” HCO₃⁻ reabsorption
• Chronic metabolic acidosis = adaptive ↑ in NH₃
synthesis
24. Table 7-4. Causes of Metabolic Acidosis
Cause Examples Comments
Excessive production or ingestion of fixed H+
Diabetic ketoacidosis Accumulation of β-OH butyric acid and acetoacetic acid
↑ Anion gap
Lactic acidosis Accumulation of lactic acid during hypoxia
↑ Anion gap
Salicylate poisoning Also causes respiratory alkalosis
↑ Anion gap
Methanol/formaldehyde poisoning Converted to formic acid
↑ Anion gap
↑ Osmolar gap
Ethylene glycol poisoning Converted to glycolic and oxalic acids
↑ Anion gap
↑ Osmolar gap
Loss of HCO3
-
Diarrhea Gastrointestinal loss of HCO3
-
Normal anion gap
Hyperchloremia
Type 2 renal tubular acidosis (type 2 RTA) Renal loss of HCO3
-
(failure to reabsorb filtered HCO3
-
)
Normal anion gap
Hyperchloremia
Inability to excrete fixed H+
Chronic renal failure ↓ Excretion of H+
as NH4+
↑ Anion gap
Type 1 renal tubular acidosis (type 1 RTA) ↓ Excretion of H+
as titratable acid and NH4+
↓ Ability to acidify urine
Normal anion gap
Type 4 renal tubular acidosis (type 4 RTA) Hypoaldosteronism
↓ Excretion of NH4+
Hyperkalemia inhibits NH3 synthesis
Normal anion gap
25. • Serum anion gap
– Represents unmeasured anions in serum
• Phosphate, Citrate, Sulfate, protein
– Anion gap = [Na⁺] – ([Cl⁻] + [HCO₃⁻])
– Normal = 8 - 16 mEq/L
– In Metabolic acidosis, as HCO₃⁻ decreases, an
anion such as Cl⁻ must be increased for electro-
neutrality.
– ↑s by an ↑ in conc. of unmeasured anions (eg.
phoxphate, lactate) to replace HCO₃⁻
– Hyperchloremic metabolic acidosis: Cl⁻
(unmeasured anion) is increased to replace HCO₃⁻
with a normal anion gap
26. • Metabolic Alkalosis
– Loss of fixed H⁺ or gain of base → ↓ arterial [H⁺]
(alkalemia)
– Primary disturbance = ↑ [HCO₃⁻]
• Eg vomiting: lost of H⁺ in gastric acid, leaving behind HCO₃⁻ in
blood
– Compensation
• Respiratory Comp.: hypoventilation
• Renal comp: ↑ HCO₃⁻ excretion
– Filtered load exceeding reabsorption rate
– If ECF vol. contraction occurs, HCO₃⁻ reabsorption will ↑,
worsening the metabolic alkalosis
27. Table 7-5. Causes of Metabolic Alkalosis
Cause Examples Comments
Loss of H+ Vomiting Loss of gastric H+
HCO3
-
remains in the blood
Maintained by volume
contraction
Hypokalemia
Hyperaldosteronism Increased H+
secretion by
intercalated cells
Hypokalemia
Gain of HCO3
-
Ingestion of NaHCO3
-
Milk-alkali syndrome (chronic
disorder of the kidney; induced
by large amounts of alkali and
calcim in the Rx of peptic ucler,
can progress to renal failure)
Ingestion of large amounts of
HCO3
-
in conjunction with renal
failure
Volume contraction alkalosis Loop or thiazide diuretics ↑ HCO3
-
reabsorption due to ↑
angiotensin II and aldosterone
28. Figure 7-11 Generation and maintenance of metabolic alkalosis
with vomiting. ECF, Extracellular fluid
29. • Respiratory Acidosis
– D/t ↓ respiratory rate and retention of CO₂
– Pimary disturbance = ↑ arterial CO₂ → ↑[H⁺]
– Compensation
• No resp. compensation
• Renal comp
– ↑ excretion of H⁺ as titratable H⁺ and NH₄
– ↑reabsorption of “new” HCO₃⁻
– Acute resp. acidosis: renal comp. has not yet
occurred
– Chronic resp. acidosis: renal comp.
30. Table 7-6. Causes of Respiratory Acidosis
Cause Examples Comments
Inhibition of the medullary respiratory center Opiates, barbiturates, anesthetics
Lesions of the central nervous system
Central sleep apnea
Oxygen therapy Inhibition of peripheral chemoreceptors
Disorders of respiratory muscles Guillain-Barré syndrome, polio, amyotrophic
lateral sclerosis (ALS), multiple sclerosis
Airway obstruction Aspiration
Obstructive sleep apnea
Laryngospasm
Disorders of gas exchange Acute respiratory distress syndrome (ARDS) ↓ Exchange of CO2 between pulmonary
capillary blood and alveolar gas
Chronic obstructive pulmonary disease
(COPD)
Pneumonia
Pulmonary edema
31. • Respiratory Alkalosis
– d/t ↑ respiratory rate (hyperventilation)
– Primary disturbance = ↓ PCO₂
– Compensation
• No resp. comp.
• Renal comp.
– ↓H⁺ excretion
– ↑HCO₃⁻ excretion
– Acute resp. alkalosis: renal compensation has not
yet occurred
– Chronic resp. alkalosis: renal comp.
32. Table 7-7. Causes of Respiratory Alkalosis
Cause Examples Comments
Stimulation of the medullary
respiratory center
Hysterical hyperventilation
Gram-negative septicemia
Salicylate poisoning Also causes metabolic acidosis
Neurologic disorders (tumor; stroke)
Hypoxemia High altitude
Pneumonia; pulmonary embolism
Hypoxemia stimulates peripheral
chemoreceptors
Severe anemia
Mechanical ventilation
