Water, pH and Buffers

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Water, pH and Buffers
College of Medicine Medical Sciences

• Lectures 1-3, MBC226-222-241-224
• Dr Ayyub Patel

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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

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HOMEOSTASIS

• The dynamic that defines the distribution of
  water and the maintenance of pH and electrolyte
  concentrations

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HOMEOSTASIS

• Water distribution maintained by the kidneys,
  antidiuretic hormone, hypothalamic thirst
  response, respiration and perspiration

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HOMEOSTASIS

• 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)

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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

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

• Very polar
• Oxygen is highly electronegative
• H-bond donor and acceptor
• High b.p., m.p., heat of vaporization, surface
  tension

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Water dissolves polar compounds
solvation shell or hydration shell
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Non-polar substances are insoluble in water
Many lipids are amphipathic
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How detergents work?
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Hydrogen Bonding of Water
One H2O molecule can associate with 4 other H20
molecules

• Ice 4 H-bonds per water molecule
• Water 2.3 H-bonds per water molecule

Crystal lattice of ice
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Biological Hydrogen Bonds
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(No Transcript)
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non-covalent interactions
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Relative Bond Strengths
Bond type kJ/mole H3C-CH3 88 H-H 104 I
onic 40 to 200 H-bond 2 - 20 Hydrophobic
interaction 3 -10 van der Waals 0.4 - 4
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Water Solubility / Hydrophilic
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From Lehninger, 2nd ed., Ch 4
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Hydrophilic/Hydrophobic
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From Lehninger, 2nd ed., Ch 4
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Hydrophobicity
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From Lehninger, 2nd ed., Ch 4
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Hydrophobicity/Micelles
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From Lehninger, 2nd ed., Ch 4
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Water and pH relationship

• In solution water shows a very low dissociation
  (probability of H in water is 1.8 x 10-9)
• H2O H OH-
• H is actually associated with a cluster of water
  molecules and exists in solution as H3O or H5O2
  or H7O3

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Ionization of Water
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Ionization of Water
H20 H OH-
Keq1.8 X 10-16M
Keq H OH- H2O
H2O 55.5 M
H2O Keq H OH-
(1.8 X 10-16M)(55.5 M ) H OH-
1.0 X 10-14 M2 H OH- Kw
If HOH- then H 1.0 X 10-7
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pH Scale

• Devised by Sorenson (1902)
• H can range from 1M and 1 X 10-14M
• using a log scale simplifies notation
• pH -log H
• Neutral pH 7.0

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Water and pH relationship

• For Dissociation of water,

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• Where denotes concentration and K is the
  dissociation constant

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Water and pH relationship

• 1 mole of water 18g
• 1 L of water contains
• 1000 18 55.56 mol (ie pure water 55.56M)
• Molar concentration of H (or OH-) ions can be
  calculated
• H 1.8x10-9 x 55.56 1.0x10-7
• In order to avoide using ve numbers, the H is
  expressed as pH which is ve log (base 10) of
  H
• pH - log10 H pH of pure water
  is 7
• Acidic solutions have pH lt 7 while basic
  solutions have pH gt 7

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Water and pH relationship

• H1x10-6 what is the pH?
• pH - log10 H
• H0.24x10-4 what is the pH?
• H3.4x10-3 what is the pH?

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4.6
2.5
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From Devlin, 3rd ed., Ch 1
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Henderson-Hasselbach Equation
Consider the dissociation of a general acid HA
HA H A-
We can define a dissociation constant (K) where
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Rearranging gives
Taking logarithms on both sides and multiplying
by -1 gives -logH -logK
log HA/A- or pH
pK log A-/HA
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Dissociation Constant and pH
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From Marks, Marks, Smith, Ch 4
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Henderson-Hasselbalch Equation
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Henderson-Hasselbalch Equation

• This equation can be used to determine the pH if
  the pK and ratio of the ionised and unionised
  forms is known.
• The pKa (a for acid) is the ve log of the
  dissociation constant of the acid. It is the pH
  at which the ratio of the ionised and unionised
  species is equal to 1. ie the molar
  concentration of the ionised and unionsed species
  is the same.
• Similarly pKb is ve log of the dissociation
  constant of the base

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From Devlin, 3rd ed., Ch 1
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From Devlin, 3rd ed., Ch 1
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Weak Acids and Bases Equilibria

• Strong acids / bases disassociate completely
• Weak acids / bases disassociate only partially
• Enzyme activity sensitive to pH
• weak acid/bases play important role in
• protein structure/function

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Acid/conjugate base pairs
HA H2O A- H3O HA A- H HA
acid ( donates H)(Bronstad Acid) A- Conjugate
base (accepts H)(Bronstad Base)
Ka pKa value describe tendency to loose
H large Ka stronger acid small Ka weaker
acid
Ka HA- HA
pKa - log Ka
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pKa values determined by titration
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Phosphate has three ionizable H and three pKas
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Buffers

• Buffers are aqueous systems that resist changes
  in pH when small amounts of a strong acid or base
  are added.
• A buffered system consist of a weak acid and its
  conjugate base.
• The most effective buffering occurs at the region
  of minimum slope on a titration curve
• (i.e. around the pKa).
• Buffers are effective at pHs that are within /-1
  pH unit of the pKa

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Henderson-Hasselbach Equation
HA weak acid A- Conjugate base
1) Ka HA- HA
2) H Ka HA A-
3) -logH -log Ka -log HA A-
H-H equation describes the relationship
between pH, pKa and buffer concentration
4) -logH -log Ka log A- HA
5) pH pKa log A- HA
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Case where 10 acetate ion 90 acetic acid

• pH pKa log10 0.1
• 
  
• 
  0.9
• pH 4.76 (-0.95)
• pH 3.81

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Case where 50 acetate ion 50 acetic acid

• pH pKa log10 0.5
• 
  
• 
  0.5
• pH 4.76 0
• pH 4.76 pKa

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Case where 90 acetate ion 10 acetic acid

• pH pKa log10 0.9
• 
  
• 
  0.1
• pH 4.76 0.95
• pH 5.71

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Cases when buffering fails

• pH pKa log10 0.99
• 
  
• 
  0.01
• pH 4.76 2.00
• pH 6.76
• pH pKa log10 0.01
• 
  
• 
  0.99
• pH 4.76 - 2.00
• pH 2.76

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

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Physiological Buffers

• Carbon Dioxide-Bicarbonate System a major
  regulator of blood pH
• Phosphate System major regulator of cytosolic pH
• Proteins as buffers
• CO2 and HCO3 are much higher than PO4 in
  blood the reverse is true in the cytosol, PO4
  gtgtgt HCO3

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Examples - Physiological Buffers
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From Marks, Marks, Smith, Ch 4
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From Marks, Marks, Smith, Ch 4
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pH Titration Curves
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From Lehninger, 2nd ed., Ch 4
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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.
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Sample Problem Metabolic Acidosis

• The blood values of a patient were pH 7.03 and
  CO2 1.1 mM. What is the patients 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

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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.

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