Key Concepts
Social Research (2000)
- ISSN: 19419260
- ISBN: 0470843799
- PubMed: 21904102
Available from www.resalliance.org
or
Abstract
The recent focus of federal funding on comparative effectiveness research underscores the importance of clinical trials in the practice of evidence-based medicine and health care reform. The impact of clinical trials not only extends to the individual patient by establishing a broader selection of effective therapies, but also to society as a whole by enhancing the value of health care provided. However, clinical trials also have the potential to pose unknown risks to their participants, and biased knowledge extracted from flawed clinical trials may lead to the inadvertent harm of patients. Although conducting a well-designed clinical trial may appear straightforward, it is founded on rigorous methodology and oversight governed by key ethical principles. In this review, we provide an overview of the ethical foundations of trial design, trial oversight, and the process of obtaining approval of a therapeutic, from its pre-clinical phase to post-marketing surveillance. This narrative review is based on a course in clinical trials developed by one of the authors (DJM), and is supplemented by a PubMed search predating January 2011 using the keywords "randomized controlled trial," "patient/clinical research," "ethics," "phase IV," "data and safety monitoring board," and "surrogate endpoint." With an understanding of the key principles in designing and implementing clinical trials, health care providers can partner with the pharmaceutical industry and regulatory bodies to effectively compare medical therapies and thereby meet one of the essential goals of health care reform.
Available from www.resalliance.org
Page 1
Key Concepts
Lecture 2: Water, Non-covalent Interactions, pH
BIOC 384 Fall 2011
BIOC 384 Fall 2011
Page 2
Key Concepts
• Properties of WATER are crucial to understanding the properties
(structural and functional) of biomolecules because the biological
milieu is primarily aqueous — water is the solvent for most
biomolecules.
• NON-COVALENT INTERACTIONS (hydrogen bonds, ionic interactions,
van der Waals interactions, and "hydrophobic interactions") are
individually much weaker than covalent bonds, but are absolutely
crucial to the structure and function of biomolecules.
• Most biomolecules have functional groups that are weak acids or
bases, and the ionization properties of those groups are crucial to
the structures and functions of the molecules; the PH determines
the state of ionization of biomolecular weak acids and bases.
• Biological systems, intracellular and extracellular, are BUFFERED.
• Properties of WATER are crucial to understanding the properties
(structural and functional) of biomolecules because the biological
milieu is primarily aqueous — water is the solvent for most
biomolecules.
• NON-COVALENT INTERACTIONS (hydrogen bonds, ionic interactions,
van der Waals interactions, and "hydrophobic interactions") are
individually much weaker than covalent bonds, but are absolutely
crucial to the structure and function of biomolecules.
• Most biomolecules have functional groups that are weak acids or
bases, and the ionization properties of those groups are crucial to
the structures and functions of the molecules; the PH determines
the state of ionization of biomolecular weak acids and bases.
• Biological systems, intracellular and extracellular, are BUFFERED.
Page 3
Learning Objectives
• Describe the properties of H2O: its polarity, hydrogen bonding capabilities, solvent
properties, and ionization properties.
• Explain the features of 4 types of noncovalent interactions (hydrogen bonds, ionic
interaction, van der Waals interactions, and hydrophobic interactions), including the
effect of solvent polarity and of distance on those interactions. Be able to identify non-
covalent interactions on a figure
• Explain the importance of weak interactions for biomolecules
• Describe the properties of amphipathic molecules
• Relate the state of protonation of a weak acid to its pKa and the pH of its environment.
Be able to interpret and sketch a pH titration curve (dominant form when pH < pKa,
when pH=pKa and when pH > pKa
• Know and use the Henderson-Hasselbalch equation to estimate the percentage of
charged molecules in a solution
• Describe the ionization properties of phosphate and what makes it a good buffer in the
physiological pH range
• Describe the properties of H2O: its polarity, hydrogen bonding capabilities, solvent
properties, and ionization properties.
• Explain the features of 4 types of noncovalent interactions (hydrogen bonds, ionic
interaction, van der Waals interactions, and hydrophobic interactions), including the
effect of solvent polarity and of distance on those interactions. Be able to identify non-
covalent interactions on a figure
• Explain the importance of weak interactions for biomolecules
• Describe the properties of amphipathic molecules
• Relate the state of protonation of a weak acid to its pKa and the pH of its environment.
Be able to interpret and sketch a pH titration curve (dominant form when pH < pKa,
when pH=pKa and when pH > pKa
• Know and use the Henderson-Hasselbalch equation to estimate the percentage of
charged molecules in a solution
• Describe the ionization properties of phosphate and what makes it a good buffer in the
physiological pH range
Page 4
Biochemistry: reactions
and interactions at a
number of different scales
and interactions at a
number of different scales
Page 5
Water is important for life
Page 6
Water is a unique molecule
What makes it so special?
1) Water is liquid over a wide range of temperatures.
2) Water is less dense as a solid than a liquid.
3) Water is an excellent solvent because of its polar nature.
What makes it so special?
1) Water is liquid over a wide range of temperatures.
2) Water is less dense as a solid than a liquid.
3) Water is an excellent solvent because of its polar nature.
Page 7
Interactions among Biomolecules in
Aqueous Solvent
• STRONG INTERACTIONS
– Covalent bond
• REVERSIBLE (WEAK) BONDS –
NON-COVALENT INTERACTIONS
– Ionic interactions
– Hydrogen bonds
– Van der Waals interactions
– Hydrophobic effect
Aqueous Solvent
• STRONG INTERACTIONS
– Covalent bond
• REVERSIBLE (WEAK) BONDS –
NON-COVALENT INTERACTIONS
– Ionic interactions
– Hydrogen bonds
– Van der Waals interactions
– Hydrophobic effect
Page 8
Covalent Bonds
• 2 atoms share a pair of electrons to fill orbitals
• Equal or nearly equal sharing a nonpolar group
– Examples: C–C and C–H bonds (not polar)
• Unequal sharing a polar group
– 1 atom has partial positive charge (δ+)
– other atom has partial negative charge (δ–)
– Example: an O–H bond (polar)
-0.66
0.33 0.33 d+
d-
d+
• 2 atoms share a pair of electrons to fill orbitals
• Equal or nearly equal sharing a nonpolar group
– Examples: C–C and C–H bonds (not polar)
• Unequal sharing a polar group
– 1 atom has partial positive charge (δ+)
– other atom has partial negative charge (δ–)
– Example: an O–H bond (polar)
-0.66
0.33 0.33 d+
d-
d+
Page 9
Hydrogen Bonds (H bonds)
• Special electrostatic interaction
• Occurs between a weakly acidic group (donor)
and a group with a lone pair of electrons
(acceptor)
H bond between two water molecules
Hydrogen bond
donor
Hydrogen bond
acceptor
• Special electrostatic interaction
• Occurs between a weakly acidic group (donor)
and a group with a lone pair of electrons
(acceptor)
H bond between two water molecules
Hydrogen bond
donor
Hydrogen bond
acceptor
Page 10
Hydrogen Bonds (H bonds)
H bond between two water molecules
In how many hydrogen
bonds can 1 H2O
molecule participate?
H bond between two water molecules
In how many hydrogen
bonds can 1 H2O
molecule participate?
Page 11
Structure of Water
• Water molecules form
hydrogen bonding network
• Special properties of water
– high boiling,
melting
temperature
– large heat of
vaporization
Ice Liquid water
Water has tetrahedral geometry
• Water molecules form
hydrogen bonding network
• Special properties of water
– high boiling,
melting
temperature
– large heat of
vaporization
Ice Liquid water
Water has tetrahedral geometry
Page 12
http://www.youtube.com/watch?v=t5ZFoU0S5iE
Page 13
Strength of a Hydrogen Bond
Hydrogen bond ~20 kJ/mol
(Covalent bond C-C ~350 kJ/mol)
Strength of a hydrogen bond depends on:
1) distance between atoms
2) direction (geometry)
Hydrogen bond ~20 kJ/mol
(Covalent bond C-C ~350 kJ/mol)
Strength of a hydrogen bond depends on:
1) distance between atoms
2) direction (geometry)
Page 14
Solvent Properties
of Water
• Water has a strong dipole
• Good solvent for Ions/
molecules with charged
groups and polar groups
of Water
• Water has a strong dipole
• Good solvent for Ions/
molecules with charged
groups and polar groups
Page 15
Hydrogen Bonds (H bonds) also occur between
other molecules besides water
C-H Hydrogen bonds?
other molecules besides water
C-H Hydrogen bonds?
Page 16
Hydrogen bonds
in biomolecules
• Not only for water
molecules
• Polar groups forms
hydrogen bond with water
• H-bonds determine
structure of biomolecules
Protein secondary structure
DNA pairing
in biomolecules
• Not only for water
molecules
• Polar groups forms
hydrogen bond with water
• H-bonds determine
structure of biomolecules
Protein secondary structure
DNA pairing
Page 17
Ionic Interactions
• Also called salt bridges
• Electrostatic interaction between ions of opposite
charge
• Coulomb’s Law:
E = Energy of interaction
q1 and q2 are charges on the atoms
D = dielectric constant
1 for vacuum
≈2 for hexane
4–20 inside protein
≈80 for H2O
r = distance between
charged atoms
r
+q1 –q2
Is the attractive force between two oppositely charged ions
greater in water or in hexane ?
E= kq1q2/Dr
k = proportionality constant
• Also called salt bridges
• Electrostatic interaction between ions of opposite
charge
• Coulomb’s Law:
E = Energy of interaction
q1 and q2 are charges on the atoms
D = dielectric constant
1 for vacuum
≈2 for hexane
4–20 inside protein
≈80 for H2O
r = distance between
charged atoms
r
+q1 –q2
Is the attractive force between two oppositely charged ions
greater in water or in hexane ?
E= kq1q2/Dr
k = proportionality constant
Page 18
Ionic Interactions
_
Examples of ionic interactions
between a protein and ligand (“CL”)
Common ionic interaction
seen in proteins
_
Examples of ionic interactions
between a protein and ligand (“CL”)
Common ionic interaction
seen in proteins
Page 19
van der Waals interactions
• Non-covalent interactions between electrically neutral
molecules
– interaction between dipoles
– very weak, but universal, can be strong in total
– liquid nitrogen (N2), methane (CH4), protein binding etc.
Attraction between dipoles of opposite sign
• Non-covalent interactions between electrically neutral
molecules
– interaction between dipoles
– very weak, but universal, can be strong in total
– liquid nitrogen (N2), methane (CH4), protein binding etc.
Attraction between dipoles of opposite sign
Page 20
van der Waals interactions
• Optimal attraction (lowest energy) between atoms
occurs at the “van der Waals contact distance”
– Strong repulsion if too close
– Weak attraction if too far
• Optimal attraction (lowest energy) between atoms
occurs at the “van der Waals contact distance”
– Strong repulsion if too close
– Weak attraction if too far
Page 21
Page 22
Hydrophobic "Interactions"
• hydrophobic effect, the "oil drop" effect
• association of non-polar groups with each
other in aqueous systems
• hydrophobic means
“water-hating”
• misnamed
“hydrophobic
‘interactions’“
• hydrophobic effect, the "oil drop" effect
• association of non-polar groups with each
other in aqueous systems
• hydrophobic means
“water-hating”
• misnamed
“hydrophobic
‘interactions’“
Page 23
Hydrophobic Effect
• Hydrophobic groups tend to
associate together to minimize
water/hydrophobic interface
• Water-nonpolar groups interaction is
unfavorable
• Hydrophobic groups tend to
associate together to minimize
water/hydrophobic interface
• Water-nonpolar groups interaction is
unfavorable
Page 24
Hydrophobic Effect
• Water-nonpolar groups interaction is
unfavorable
• Hydrophobic groups tend to
associate together to minimize
water/hydrophobic interface
• Water-nonpolar groups interaction is
unfavorable
• Hydrophobic groups tend to
associate together to minimize
water/hydrophobic interface
Page 25
http://www.youtube.com/watch?v=ETMmH2trTpM
Page 26
Amphipathic molecules
– contain polar and nonpolar (hydrophobic) regions
– polar region tends to be exposed to water
– nonpolar region tends to cluster together
minimize solvent
exposed surface
polar
nonpolar
– contain polar and nonpolar (hydrophobic) regions
– polar region tends to be exposed to water
– nonpolar region tends to cluster together
minimize solvent
exposed surface
polar
nonpolar
Page 27
Biological importance
• Many biomolecules are amphipathic
– proteins, certain vitamins, lipids etc
– “hydrophobic interactions” between lipids and
protein determines the structure of membranes
• Many biomolecules are amphipathic
– proteins, certain vitamins, lipids etc
– “hydrophobic interactions” between lipids and
protein determines the structure of membranes
Page 28
4 Types of Noncovalent (“Weak”) Interactions
Type Brief description and example Stabilization
Energy
(kJ/mol)
Length
(Å)
Hydrogen
bonds
Between a hydrogen atom covalently bonded to
an electronegative atom and a second
electronegative atom
10-30 1.8-3.0
Ionic
interactions
Interaction between ions of opposite charge 20-30 2.5
van der
Waals
interactions
Interaction between opposite dipoles
(permanent, temporary or induced dipoles)
1-5 1-2
Hydrophobic
effect
The tendency of nonpolar groups and molecules
to cluster together in aqueous solution
5-30 --
Type Brief description and example Stabilization
Energy
(kJ/mol)
Length
(Å)
Hydrogen
bonds
Between a hydrogen atom covalently bonded to
an electronegative atom and a second
electronegative atom
10-30 1.8-3.0
Ionic
interactions
Interaction between ions of opposite charge 20-30 2.5
van der
Waals
interactions
Interaction between opposite dipoles
(permanent, temporary or induced dipoles)
1-5 1-2
Hydrophobic
effect
The tendency of nonpolar groups and molecules
to cluster together in aqueous solution
5-30 --
Page 29
pH is controlled
pH log[H]
• Biological processes are
pH-dependent
• Arrangement of protons is essential
for enzyme activity
pH log[H]
• Biological processes are
pH-dependent
• Arrangement of protons is essential
for enzyme activity
Page 30
Why is it so important that pH is
controlled?
Example:
arrangement of protons is
essential for enzyme activity
controlled?
Example:
arrangement of protons is
essential for enzyme activity
Page 31
Acid and Bases
HA ⇌ H+ + A–
HA+ ⇌ H+ + A
Acid
(conjugate)
Base
Proton
Donor
Proton
Acceptor
Acetic Acid CH3COOH -> CH3COO
– + H+
Carboxyl Group: R-COOH -> R-COO– + H+
Ammonium Ion NH3
+ -> NH2 + H
+
Amino Group: R-NH3
+ -> R-NH2 + H
+
proton
pH log[H]
HA ⇌ H+ + A–
HA+ ⇌ H+ + A
Acid
(conjugate)
Base
Proton
Donor
Proton
Acceptor
Acetic Acid CH3COOH -> CH3COO
– + H+
Carboxyl Group: R-COOH -> R-COO– + H+
Ammonium Ion NH3
+ -> NH2 + H
+
Amino Group: R-NH3
+ -> R-NH2 + H
+
proton
pH log[H]
Page 32
Titration Curve
pH pKa log
[A]
[HA]
Henderson-Hasselbalch equation
How do we determine the pKa ?
pKa = pH at the midpoint
The pKa is the pH at which the
concentration of the protonated
form equals the concentration
of the deprotonated form
[HA] [A]
pH pKa log
[A]
[HA]
Henderson-Hasselbalch equation
How do we determine the pKa ?
pKa = pH at the midpoint
The pKa is the pH at which the
concentration of the protonated
form equals the concentration
of the deprotonated form
[HA] [A]
Page 33
Protonation
depends on pH
pH < pKa pKa < pH
dominant form protonated un-protonated
dominant form
HA ⇌ H+ + A–
HA
(charge = 0)
A–
(charge = -1)
dominant form
HA+ ⇌ H+ + A
HA+
(charge = +1)
A
(charge = 0)
• at lower pH, more protonated
• protonation state changes
when pH > pKa
• when pH ≈ pKa, two states
coexist
=HA
=A-
[HA]=[A-]
Percent protonated
0% 100%
depends on pH
pH < pKa pKa < pH
dominant form protonated un-protonated
dominant form
HA ⇌ H+ + A–
HA
(charge = 0)
A–
(charge = -1)
dominant form
HA+ ⇌ H+ + A
HA+
(charge = +1)
A
(charge = 0)
• at lower pH, more protonated
• protonation state changes
when pH > pKa
• when pH ≈ pKa, two states
coexist
=HA
=A-
[HA]=[A-]
Percent protonated
0% 100%
Page 34
Protonation
depends on pH
pH < pKa pKa < pH
dominant form protonated un-protonated
dominant form
HA ⇌ H+ + A–
HA
(charge = 0)
A–
(charge = -1)
dominant form
HA+ ⇌ H+ + A
HA+
(charge = +1)
A
(charge = 0)
• at lower pH, more protonated
• protonation state changes
when pH > pKa
• when pH ≈ pKa, two states
coexist
=HA
=A-
[HA]=[A-]
Percent protonated
0% 100%
depends on pH
pH < pKa pKa < pH
dominant form protonated un-protonated
dominant form
HA ⇌ H+ + A–
HA
(charge = 0)
A–
(charge = -1)
dominant form
HA+ ⇌ H+ + A
HA+
(charge = +1)
A
(charge = 0)
• at lower pH, more protonated
• protonation state changes
when pH > pKa
• when pH ≈ pKa, two states
coexist
=HA
=A-
[HA]=[A-]
Percent protonated
0% 100%
Page 35
Example
• The pKa of a molecule is about 8. At pH 7, the
percent that would be charged would be
approximately:
– 10 %?
– 50 %?
– 90 %?
• The pKa of a molecule is about 8. At pH 7, the
percent that would be charged would be
approximately:
– 10 %?
– 50 %?
– 90 %?
Page 36
Page 37
If the pKa of aspirin is 3.5, what percentage of aspirin is protonated in the stomach
at pH 1.5?
What percentage of aspirin is protonated in the small intestine at pH 6.0?
at pH 1.5?
What percentage of aspirin is protonated in the small intestine at pH 6.0?
Page 38
Titration Curve and Buffer
• Buffering region:
small change in pH
against titration
• pKa ± 1
• Buffer – Mixture of
weak acids and
conjugate bases
• Buffering region:
small change in pH
against titration
• pKa ± 1
• Buffer – Mixture of
weak acids and
conjugate bases
Page 39
Physiologically Important Buffer
Systems – Phosphate Buffer
• In the cytoplasm of all cells
• Effective range 5.9 – 7.9
• pH of extracellular fluid and most cytoplasmic compartments
= 6.9 – 7.4
• other biologically important buffers
– bicarbonate buffer (controls blood pH)
– proteins (Hemoglobin)
H2PO4
– ⇌ H+ + HPO4
2– (pKa=6.86)
Systems – Phosphate Buffer
• In the cytoplasm of all cells
• Effective range 5.9 – 7.9
• pH of extracellular fluid and most cytoplasmic compartments
= 6.9 – 7.4
• other biologically important buffers
– bicarbonate buffer (controls blood pH)
– proteins (Hemoglobin)
H2PO4
– ⇌ H+ + HPO4
2– (pKa=6.86)
Page 40
• A brief review from general chemistry is posted on D2L (Gen
Chem Review.pdf) which includes some practice problems.
• NOTE: you will need a simple scientific calculator capable of
doing log and ln calculations. If you don’t already have one,
please purchase one and practice using it before the exam.
• Programmable calculators will not be allowed for the exams
Chem Review.pdf) which includes some practice problems.
• NOTE: you will need a simple scientific calculator capable of
doing log and ln calculations. If you don’t already have one,
please purchase one and practice using it before the exam.
• Programmable calculators will not be allowed for the exams
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