Transcript Water, pH and dissociation Lecture 3, Medical Biochemistry Lecture 3 Outline • Homeostasis • The structure and function of water • Dissociation of weak.

Water, pH and dissociation
Lecture 3, Medical Biochemistry
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
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, overexercise), and excess body water due to
increased intake (too much I.V.) or decreased
excretion (renal failure)
Structure of H20
From Lehninger, 2nd ed., Ch 4
WATER
• Comprises approx 70% of human mass (4560% 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
Water Solubility / Hydrophilic
From Lehninger, 2nd ed., Ch 4
Hydrophilic/Hydrophobic
From Lehninger, 2nd ed., Ch 4
Hydrophobicity
From Lehninger, 2nd ed., Ch 4
Hydrophobicity/Micelles
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
Dissociation Constant and pH
From Marks, Marks, Smith, Ch 4
Henderson-Hasselbalch Equation
Sample pH problems
Sample pH Problem (cont)
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
2) the concentration of the buffer; the more
present, the greater the buffering capacity
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]
Examples - Physiological Buffers
From Marks, Marks, Smith, Ch 4
From Marks, Marks, Smith, Ch 4
From Lehninger, 2nd ed., Ch 4
Blood Bicarbonate and Metabolic
Acidosis
The bicarbonate blood buffer in a normal adult
maintains the blood pH at about 7.40. 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.
Sample Problem – Metabolic Acidosis
• The blood values of a patient were pH =
7.03 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
Solution
•
•
•
•
Substitute into Henderson-Hasselbalch equation:
7.03 = 6.10 + 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.