1. Water, pH and dissociation
Lecture 3, Medical Biochemistry

2. Lecture 3 Outline Homeostasis The structure and function of water
Dissociation of weak acids and weak bases
pH and the Henderson-Hasselbalch equation
Buffers, biological/physiological examples

3. HOMEOSTASIS
The dynamic that defines the distribution of water and the maintenance of pH and electrolyte concentrations
Water distribution maintained by the kidneys, antidiuretic hormone, hypothalamic thirst response, respiration and perspiration
Clinically, need to be aware of water depletion caused by decreased intake (coma, wandering the desert) or increased loss (diarrhea, renal malfunction, over-exercise), and excess body water due to increased intake (too much I.V.) or decreased excretion (renal failure)

4. Structure of H20
From Lehninger, 2nd ed., Ch 4

5. WATER
Comprises approx 70% of human mass (45-60% intracellular, 25% extracellular/blood plasma)
dipolar: partial negative charge on oxygen, partial positive charge on hydrogens
dipolar nature leads to formation of many low energy hydrogen bonds

6. Water Solubility / Hydrophilic
From Lehninger, 2nd ed., Ch 4

7. Hydrophilic/Hydrophobic
From Lehninger, 2nd ed., Ch 4

8. Hydrophobicity
From Lehninger, 2nd ed., Ch 4

9. Hydrophobicity/Micelles
From Lehninger, 2nd ed., Ch 4

10. Summary of water and pH relationship
Very low dissociation of H2O to H+ or OH-
The ion product of H2O, Keq X 55.5 M, leads to this: [H+] = [OH-] = 1 X 10-7 M for pure H2O which is a constant in biological systems
Therefore, if [H+] > 10-7 M, then [OH-] must be less than 10-7 M, and vice versa.
Thus, if the negative logarithm of [H+] is derived ( pH = -log [H+] ), pure water would be pH = 7, acids pH < 7, and bases pH > 7

11. From Devlin, 3rd ed., Ch 1

12. Dissociation Constant and pH
From Marks, Marks, Smith, Ch 4

13. Henderson-Hasselbalch Equation

14. From Devlin, 3rd ed., Ch 1

15. From Devlin, 3rd ed., Ch 1

16. Sample pH problems
From Devlin, 3rd ed., Ch 1

17. Sample pH Problem (cont)
From Devlin, 3rd ed., Ch 1

18. Buffers
Definition: A weak acid plus its conjugate base that cause a solution to resist changes in pH when an acid or base are added
Effectiveness of a buffer is determined by: 1) the pH of the solution, buffers work best within 1 pH unit of their pKa ) the concentration of the buffer; the more present, the greater the buffering capacity

19. Physiological Buffers
Carbon Dioxide-Bicarbonate System; a major regulator of blood pH
Phosphate System; major regulator of cytosolic pH
[CO2] and [HCO3] are much higher than [PO4] in blood; the reverse is true in the cytosol, [PO4] >>> [HCO3]

20. Examples - Physiological Buffers
From Marks, Marks, Smith, Ch 4

21. From Marks, Marks, Smith, Ch 4

22. pH Titration Curves
From Lehninger, 2nd ed., Ch 4

23. Blood Bicarbonate and Metabolic Acidosis
The bicarbonate blood buffer in a normal adult
maintains the blood pH at about If the blood
pH drops below 7.35, the condition is referred to
as an ACIDOSIS. A prolonged blood pH below
7.0 can lead to death. Clinically for an acidosis,
the acid-base parameters (pH, [HCO3- ], [CO2] )
of the patients blood should be monitored. The
normal values for these are pH = 7.40;
[HCO3- ] = 24 mM; [CO2] = 1.2 mM.

24. Sample Problem – Metabolic Acidosis
The blood values of a patient were pH = and [CO2] = 1.1 mM. What is the patient’s blood [HCO3-] and how much of the normal [HCO3-] has been used in buffering the acid causing the condition?
The pK’ for [HCO3-]/[CO2] = 6.10

25. Solution Substitute into Henderson-Hasselbalch equation:
7.03 = log [HCO3-]/1.1 mM, or
0.93 = log [HCO3-]/1.1 mM
The anti-log of 0.93 = 8.5, thus:
8.5 = [HCO3-]/1.1 mM, or [HCO3-] = 9.4 mM
Since normal [HCO3-] equals 24 mM, there was a decrease of 14.6 mmol of [HCO3- per liter of blood in this patient. This would be approaching the point where, if left untreated, the HCO3- buffering capacity would be no longer effective in this patient.