Chapter - Clayton State University

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Transcript Chapter - Clayton State University

Section 1
Introduction to Biochemical
Principles
Chapter 3
Water: The Matrix of Life
Section 3.1: Molecular Structure of Water
Water is essential for life
Water’s important properties include:
Chemical stability
Remarkable solvent properties
Role as a biochemical reactant
Hydration of organism
Section 3.1: Molecular Structure of Water
Water has a tetrahedral
geometry
Oxygen is more electronegative
than hydrogen (3.5 vs. 2.1)
Figure 3.2 Tetrahedral
Structure of Water
Section 3.1: Molecular Structure of Water
Larger oxygen atom has partial negative charge
(d-) and hydrogen atoms have partial positive
charges (d+)
Figure 3.3 Charges on a Water Molecule
Figure 3.4 Water Molecule
Section 3.1: Molecular Structure of Water
Bond between oxygen and hydrogen is polar
Water is a dipole because the positive and
negative charges are separate
Figure 3.5 Molecular
Dipoles in an Electric
Field
Section 3.1: Molecular Structure of Water
An electron-deficient hydrogen of one
water is attracted to the unshared
electrons of water forming a hydrogen
bond
Can occur with oxygen, nitrogen,
and sulfur (i.e., when there is
significant difference in EN
between H and higher EN atom)
Has electrostatic (i.e., opposite
charges) and covalent (i.e., electron
sharing) characteristics
Large numbers of hydrogen bonds
lead to extended network
Section 3.2: Noncovalent Bonding
Noncovalent interactions are electrostatic
Weak individually, but play vital role in biomolecules
because of cumulative effects
Individual H-bonds not strong (~20 kJ/mol), but
collectively, they aggregate
Why water has high MP, BP
Section 3.2: Noncovalent Bonding
Three most important noncoavalent bonds:
Ionic interactions
Van der Waals forces
Hydrogen bonds
Section 3.2: Noncovalent Bonding
Ionic Interactions
Oppositely charged ions attract one another
Biochemistry primarily investigates the interaction of charged
groups on molecules, which differs from ionic interactions like those of
ionic compounds (e.g., NaCl)
Ionized amino acid
side chains can
form salt bridges
with one another
Section 3.2: Noncovalent Bonding
Van der Waals Forces:
Occur between molecules
with permanent, and/or
induced dipoles
Figure 3.8 Dipolar Interactions
Three types of interactions:
Dipole-dipole
Two permanent dipoles align
themselves by charge
(e.g., H-bonds)
Dipole-induced dipole
Permanent dipole induces transient
dipole in another molecule by
distorting its electron distribution
(e.g., carbonyl induces dipole in
aromatic ring pi electrons)
Induced dipole-induced dipole
Transient dipole in 1 molecule induces
transient dipole in another molecule
(e.g., Stacking of base rings in
DNA: individually weak but
collectively strong)
Section 3.3: Thermal Properties of Water
Water’s melting and boiling points are exceptionally high
due to hydrogen bonding
Each water molecule can form four hydrogen bonds
with other water molecules
Forms extended network of hydrogen bonds
This explains thermal properties of water: It requires a
lot of E input into system to melt ice and boil water, in
order to overcome E of H-bonds
Section 3.3: Thermal Properties of Water
Water has an exceptionally high heats of fusion, and
vaporization.
Takes a lot of energy to convert water from solid to
liquid, and liquid to gas.
Related to H-bonding.
Section 3.3: Thermal Properties of Water
Water has a high heat capacity.
Heat capacity is energy needed to change T of 1 g of
substance by 1 °C.
Water can absorb a lot of E before T increases.
q = mCΔT
q = Energy in Joules
m = mass (g)
C = heat capacity
ΔT = Change in temperature
Q: How much E is required to raise T
of 1 Kg of water from 65 °C to 75 °C?
C of water = 4.184 J/g·°C
Most living organisms are comprised of ~50-95% water.
The high water content and heat capacity of water help
maintain internal T of organisms.
Evaporation of water from body is cooling mechanism.
Human adults eliminate ~1200g of water daily in expired
air, sweat, and urine.
Section 3.4: Solvent Properties of Water
Figure 3.10 Solvation
Spheres
Water is an ideal biological solvent
Dissolves ions, sugars, many amino acids (ionic and polar substances)
Water forms solvation spheres around ions, molecules, thus separating
them
Does not dissolve lipids and some amino acids
Allows biological processes (e.g., protein folding and membrane
formation)
Section 3.4: Solvent Properties of Water
Structured Water
Water is rarely free flowing
Most of the time, water
molecules in cell are noncovalently associated with
macromolecules and other
cellular components
(membranes, proteins, etc.)
A single layer of water
molecules on the surface of a
biomolecule attracts an
second layer, and so on.
Forms complex threedimensional bridges between
cellular components
Figure 3.11 Diagrammatic
View of Structured Water
Section 3.4: Solvent Properties of Water
Figure 3.12 Amoeboid
Movement
Sol-Gel Transitions
Cytoplasm has properties of a gel (semi-solid colloidal
mixture)
Transition from gel to sol (more liquid state) important in
cell movement
Amoeboid motion provides an example of regulated, cellular,
sol-gel transitions
Caused by reversible polymerization of G-actin to form Factin (gel becomes more solid, expands, then gel becomes
more liquid, creating a contractile force that pulls the gel
forward)
Section 3.4: Solvent Properties of Water
Figure 3.13 The
Hydrophobic Effect
Hydrophobic Molecules and the Hydrophobic Effect
Small amounts of nonpolar substances are excluded from the
solvation network forming droplets
This hydrophobic effect results from the solvent properties of
the water and is stabilized by van der Waals interactions
Water molecule form a cage around the hydrophobic droplets,
isolating them.
This effect generates stable lipid membranes and contributes
to protein folding.
Section 3.4: Solvent Properties of Water
Amphipathic Molecules
Figure 3.14 Formation of Micelles
Contain both polar and
nonpolar groups
Amphipathic molecules
form micelles when mixed
with water
Important feature for
the formation of
cellular compartments,
membranes
Section 3.4: Solvent Properties of Water
Figure 3.15 Osmotic
Pressure
Osmotic Pressure
Osmosis is the spontaneous passage of solvent molecules
through a semipermeable membrane
Osmotic pressure is the pressure required to stop the net
flow of water across the membrane
Over time, water diffuses from side A (more dilute) to
side B (more concentrated)
Osmotic pressure depends on solute concentration
Section 3.4: Solvent Properties of Water
Can be measured with an osmometer or calculated
( =iMRT, p. 89)
Cells may gain or lose water because of the
environmental solute concentration
Solute concentration differences between the cell and
the environment can have important consequences.
(a) When cells are placed in isotonic solution (i.e.,
concentrations of solute and water are the same on
both sides of selectively permeable membrane), no
movement of water is observed in either direction.
(b) A cell in a hypotonic solution (i.e., solution has
lower solute concentration), water moves into cells.
For example, red blood cells in pure water will
absorb water, swell and rupture.
(c) In a hypertonic solution (i.e., higher solute
concentration), water moves out of cell and cell
shrivels.
Figure 3.17 Effect of Solute Concentration on Animal Cells
Section 3.4: Solvent Properties of Water
Most macromolecules have little effect on cellular osmolality, but
proteins can, because their ionizable amino acid side chains attract
ions of opposite charge.
Size of an ions solvation sphere is inversely related to its charge
density. For example, Na+ has a smaller radius than K+, but larger
hydration sphere due to charge density. It requires less energy to
remove solvation sphere from K+, so there is a tendency for more K+
to accumulate in the cell.
Membrane Potential
Intracellular proteins tend
to produce a significant net
negative charge in cell,
which is partially offset by
K+ ions.
The extracellular
environment has a net
positive charge due to Na+
ions.
This asymmetric charge
distribution results in
formation of an electrical
gradient (membrane
potential) which provides
the means for electrical
conduction, active
transport, and passive
transport
Self-Ionization of Water



Pure water contains H2O molecules.
In addition, small but equal amounts of H3O+ and OHions are also present.
The reason for this is that in one liter of pure water 1.0 x
10-7 moles of water molecules behave as Brønsted
acids and donate protons to another 1.0 x 10-7 moles of
water molecules, which act as Brønsted bases. The
reaction is:
K
H2O (l) + H2O (l) ⇆ H3O+ (aq) + OH− (aq)



As a result, absolutely pure water contains 1.0 x 10-7
mol/L of both H3O+ and OH-.
The term neutral is used to describe any water solution
in which the concentrations of H3O+ and OH- are equal.
Thus, pure water is neutral because each of the ions is
present at a concentration of 1.0 x 10-7 M.
Self-Ionization of Water

Equilibrium in reactions
K
• Equilibrium of water
K
 Kw is ion product of water
at 25°C and 1 atm pressure,
Kw is a constant, but temperature dependent. The concentration of either acid or base may
change, but Kw does not.
If [H3O+] > [OH-], solution acidic, e.g., [H3O+] = 1.0 x 10-4, so [OH-] = 1.0 x 10-10
Section 3.5: Ionization of Water
Acids, Bases, and pH (BL definitions)
An acid is a proton donor
A base is a proton acceptor
Most organic molecules that donate or accept
protons are weak acids or weak bases
A deprotonated product of a dissociation reaction
is a conjugate base
Section 3.5: Ionization of Water
The pH scale can be used to
measure hydrogen ion
concentration
pH=-log[H+]
Figure 3.18 The pH Scale and the pH Values of
Common Fluids
Section 3.5: Ionization of Water
pKa is used to express the
strength of a weak acid
Lower pKa equals a stronger
acid
pKa=-logKa
Ka is the acid dissociation
constant
Figure 3.18 The pH Scale and the pH Values of
Common Fluids
Section 3.5: Ionization of Water
Section 3.5: Ionization of Water
Buffers
Regulation of pH is universal and essential for all
living things
Certain diseases can cause changes in pH that can
be disastrous
Acidosis (blood pH falls below 7.35; blood pH below
7.00 leads to death) and Alkalosis (blood pH above
7.45, leads to convulsions and respiratory arrest)
Buffers help maintain a relatively constant
hydrogen ion concentration
Commonly composed of a weak acid and its
conjugate base
Section 3.5: Ionization of Water
Buffers Continued
Establishes an
equilibrium between
buffer’s components
Follows Le Chatelier’s
principle
Equilibrium shifts in
the direction that
relieves the stress
Figure 3.19 Titration of Acetic Acid
with NaOH
Section 3.5: Ionization of Water
Henderson-Hasselbalch Equation
Establishes the relationship between pH and pKa for
selecting a buffer
Buffers are most effective when they are composed of
equal parts weak acid and conjugate base
Biological buffers usually contain more conj. base,
since biological systems generate acid during
metabolism
Best buffering occurs 1 pH unit above and below the
pKa
Henderson-Hasselbalch Equation
pH = pKa + log
[A-]
[HA]
Section 3.5: Ionization of Water
Buffering Capacity
Capacity to maintain a certain pH depends on:
Molar concentration of acid-conj. Base pair
Ratio of their concentrations
Increasing concentration increases buffering capacity
Buffer concentration is sum of concentration of weak
acid and its conjugate base
Section 3.5: Ionization of Water
Worked Problem 3.5 (Page 91)
Calculate the pH of a mixture of 0.25 M acetic acid (CH3COOH)
and 0.1 M sodium acetate (NaC2H3O2)
The pKa of acetic acid is 4.76
Solution:
pH = pKa + log
pH = 4.76 + log
[acetate]
[acetic acid]
[0.1]
[0.25]
= 4.76 + 0.398 = 4.36
What happens if we add HCl to a concentration of 0.75 M HCl?
Section 3.5: Ionization of Water
In-Class Problem
Calculate the pH of a mixture prepared by mixing 150 mL of 0.10 M HCl
with 300 mL of 0.10 M sodium acetate (NaC2H3O2), and diluting the
mixture to 1.0 L. The pKa of acetic acid is 4.76. (CH3COOH) (NaC2H3O2)
Section 3.5: Ionization of Water
In-Class Problem
Calculate the pH of a mixture prepared by mixing 150 mL of 0.10 M HCl
with 300 mL of 0.10 M sodium acetate (NaC2H3O2), and diluting the
mixture to 1.0 L. The pKa of acetic acid is 4.76. (CH3COOH) (NaC2H3O2)
Solution:
250  × 0.10  = 25  
300  × 0.10  = 30  
pH = pKa + log
pH = 4.76 + log
[acetate]
[acetic acid]
5
25
= 4.76 – 0.70 = 4.06
Section 3.5: Ionization of Water
Figure 3.20 Titration of
Phosphoric Acid with
NaOH
Weak Acids with Multiple Ionizable Groups
Each ionizable group can have its own pKa
Protons are released in a stepwise fashion
Section 3.5: Ionization of Water
Physiological Buffers
Buffers adapted to solve specific physiological
problems within the body
Bicarbonate Buffer
One of the most important buffers in the blood
CO2 + H2O  H+ + HCO3- (HCO3- is bicarbonate):
This is a reversible reaction
Carbonic anhydrase is the enzyme responsible
Section 3.5: Ionization of Water
Phosphate Buffer
Consists of H2PO4-/HPO42(weak acid/conjugate base)
H2PO4-  H+ + HPO42Important buffer for
intracellular fluids
Protein Buffer
Proteins are a significant
source of buffering capacity
(e.g., hemoglobin)
Figure 3.21 Titration of
H2PO4- by Strong Base