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Title: Information for BTY221 students


1
Information for BTY221 students
TEXTBOOK BIOCHEMISTRY by Garret Grisham 3rd
Edition
PRACTICALS Thursday at 10H50 , Venue
N105 Co-ordinator Ms E Anthony
TESTS Test 1 Wed 22 Feb 2006 at 17H00 Test 2
Wed 22 March 2006 at 17H00 Venue - OA, OB and OD
(To be confirmed)
2
EXAMINATIONS Final examinations held in June
2006. A pass 40 in final examination, final
mark of 50 To qualify for a supplementary
examination (i) must have a year mark of 50 or
(ii) must be have obtained a final mark of
between 45 and 50.
EVALUATION MARK Tests (T1 T2)
20 Tuts (tut1 tut2 tut3)
10 Practicals 30 EXAM MARK
40
  • Note
  • Absence from any test /examination requires
  • an authentic doctors certificate.
  • Read the notice-boards!

3
Chapter 1
  • Chemistry Is the Logicof Biological Phenomena
  • Biochemistry
  • by
  • Reginald Garrett and Charles Grisham

4
Essential Question
  • Despite the spectacular diversity of life, the
    elaborate structure of biological molecules, and
    the complexity of vital mechanisms, are life
    functions ultimately interpretable in chemical
    terms?

5
Outline
  • What Are the Distinctive Properties of Living
    Systems?
  • What Kinds of Molecules Are Biomolecules?
  • What Is the Structural Organization of Complex
    Biomolecules?
  • How Do the Properties of Biomolecules Reflect
    Their Fitness to the Living Condition?
  • What Is the Organization and Structure of Cells?
  • What Are Viruses?

6
On Life and Chemistry...
  • Living things are composed of lifeless
    molecules (Albert Lehninger)
  • Chemistry is the logic of biological phenomena
    (Garrett and Grisham)

7
1.1 - Distinctive Properties of Living Systems
  • Organisms are complicated and highly organized
    (molecules, organelles, cells)
  • Biological structures serve functional purposes
  • Living systems are actively engaged in energy
    transformations (photosynthesis, metabolism, ATP,
    NADPH)
  • Living systems have a remarkable capacity for
    self-replication (DNA)

8
Figure 1.2The food pyramid. Photosynthetic
organisms at the base capture light energy.
Herbivores and carnivores derive their energy
ultimately from these primary producers.  
9
Figure 1.5The DNA double helix. Two
complementary polynucleotide chains running in
opposite directions can pair through hydrogen
bonding between their nitrogenous bases. Their
complementary nucleotide sequences give rise to
structural complementary.
10
1.2 - Biomolecules The Molecules of Life
  • H, O, C and N make up 99 of atoms in the human
    body
  • ELEMENT PERCENTAGE
  • Oxygen 63
  • Hydrogen 25.2
  • Carbon 9.5
  • Nitrogen 1.4

11
1.2 - Biomolecules The Molecules of Life
  • What property unites H, O, C and N and renders
    these atoms so appropriate to the chemistry of
    life?
  • Answer Their ability to form covalent bonds by
    electron-pair sharing.

12
1.2 - Biomolecules The Molecules of Life
  • What are the bond energies of covalent bonds?
  • Bond Energy kJ/mol
  • H-H 436
  • C-H 414
  • C-C 343
  • C-O 351

13
Figure 1.6Covalent bond formation by e- pair
sharing.
  • C is versatile
  • Can form up to 4 covalent bonds
  • Can bond with itself
  • Tetrahedral nature of 4 C bonds -gt
  • structural variety

14
Figure 1.7Examples of the versatility of CC
bonds in building complex structures linear
aliphatic, cyclic, branched, and planar.  
15
1.3 - A Biomolecular Hierarchy
  • Simple Molecules are the Units for Building
    Complex Structures
  • Metabolites and Macromolecules
  • Organelles
  • Membranes
  • The Unit of Life is the Cell

16
Inorganic precursors
metabolites
Figure 1.8Molecular organization in the cell is
a hierarchy.
Building blocks
macromolecules
Supramolecular complexes
organelles
The cell
17
1.4 - Properties of Biomolecules Reflect Their
Fitness to the Living Condition
  • Macromolecules and Their Building Blocks Have a
    Sense or Directionality
  • Macromolecules are Informational
  • Biomolecules Have Characteristic
    Three-Dimensional Architecture
  • Weak Forces Maintain Biological Structure and
    Determine Biomolecular Interactions

18
Figure 1.9(a) Amino acids build proteins by
connecting the a-carboxyl C atom of one amino
acid to the a -amino N atom of the next amino
acid in line. (b) Polysaccharides are built by
combining the C-1 of one sugar to the C-4 O of
the next sugar in the polymer. (c) Nucleic acids
are polymers of nucleotides linked by bonds
between the 3 -OH of the ribose ring of one
nucleotide to the 5 -PO4 of its neighboring
nucleotide. All three of these polymerization
processes involve bond formations accompanied by
the elimination of water (dehydration synthesis
reactions).
b
a
c
19
IN008
Figure 1.10The sequence of monomeric units in a
biological polymer has the potential to contain
information if the diversity and order of the
units are not overly simple or repetitive.
Nucleic acids and proteins are information-rich
molecules polysaccharides are not.
20
Non-covalent forces determine Biomolecular
structure and function
  • Eg H bonds, Van der Waals forces, ionic and
    hydrophobic interactions
  • Van der Waals forces-gt induced interactions
    between molecules or atoms as their e- clouds
    fluctuate in time, ve nuclei attract e- , weak
    but additive repulsive
  • H bonds -gt between H atoms bonded to e- atoms and
    another e- atoms specific and directional

21
Figure 1.12Van der Waals packing is enhanced in
molecules that are structurally complementary.
Gln121 represents a surface protuberance on the
protein lysozyme. This protuberance fits nicely
within a pocket (formed by Tyr101, Tyr32, Phe91,
and Trp92) in the antigen-binding domain of an
antibody raised against lysozyme. (See also
Figure 1.16.) (a) A space-filling representation.
(b) A ball-and-stick model. (From Science 233751
(1986), figure 5.)
22
Figure 1.13The van der Waals interaction energy
profile as a function of the distance, r, between
the centers of two atoms. The energy was
calculated using the empirical equation U  
B/r12 - A/r6. (Values for the parameters B 11.5
x 10-6 kJnm12/mol and A 5.96 x 10-3 kJnm6/mol
for the interaction between two carbon atoms are
from Levitt, M., 1974, Journal of Molecular
Biology 82393-420.)
23
Figure 1.14Some of the biologically important H
bonds and functional groups that serve as H bond
donors and acceptors.
24
Figure 1.15Ionic bonds in biological molecules.
25
1.4 - Properties of Biomolecules Reflect Their
Fitness to the Living Condition
  • Important numbers!
  • van der Waals 0.4-4.0 kJ/mole
  • Hydrogen bonds 12-30 kJ/mole
  • Ionic bonds 20 kJ/mole
  • Hydrophobic interactions lt40 kJ/mole

26
Two Important Points About Weak Forces
  • Biomolecular Recognition is Mediated by Weak
    Chemical Forces
  • Weak Forces Restrict Organisms to a Narrow Range
    of Environmental Conditions

27
Figure 1.16 Structural complementary the
pieces of a puzzle, the lock and its key, a
biological macromolecule and its ligandan
antigenantibody complex. (a) The antigen on the
right (green) is a small protein, lysozyme, from
hen egg white. The part of the antibody molecule
(IgG) shown on the left in blue and yellow
includes the antigen-binding domain. (b) This
domain has a pocket that is structurally
complementary to a surface protuberance (Gln121,
shown in red between antigen and antigen-binding
domain) on the antigen. (See also Figure
1.12.)(photos, courtesy of Professor Simon E. V.
Philips)
28
Figure 1.17Denaturation and renaturation of the
intricate structure of a protein.
29
Figure 1.18 Metabolism is the organized release
or capture of small amounts of energy in
processes whose overall change in energy is
large. (a) For example, the combustion of glucose
by cells is a major pathway of energy production,
with the energy captured appearing as 30 to 38
equivalents of ATP, the principal energy-rich
chemical of cells. The ten reactions of
glycolysis, the nine reactions of the citric acid
cycle, and the successive linked reactions of
oxidative phosphorylation release the energy of
glucose in a stepwise fashion and the small
packets of energy appear in ATP. (b) Combustion
of glucose in a bomb calorimeter results in an
uncontrolled, explosive release of energy in its
least useful form, heat.
30
Enzymes catalyse metabolic reactions
31
Organization and Structure of Cells
  • Prokaryotic cells
  • A single (plasma) membrane
  • no nucleus or organelles
  • Eukaryotic cells
  • much larger in size than prokaryotes
  • 103-104 times larger!
  • Nucleus plus many organelles
  • ER, Golgi, mitochondria, etc.

32
Figure 1.21 This bacterium is Escherichia coli,
a member of the coliform group of bacteria that
colonize the intestinal tract of humans. E. coli
organisms have rather simple nutritional
requirements. They grow and multiply quite well
if provided with a simple carbohydrate source of
energy (such as glucose), ammonium ions as a
source of nitrogen, and a few mineral salts. The
simple nutrition of this lower organism means
that its biosynthetic capacities must be quite
advanced. When growing at 37C on a rich organic
medium, E. coli cells divide every 20 minutes.
Subcellular features include the cell wall,
plasma membrane, nuclear region, ribosomes,
storage granules, and cytosol (Table 1.5).
(photo, Martin Rotker/Phototake, Inc. inset
photo, David M. Phillips/The Population
Council/Science Source/Photo Researchers, Inc.)
33
Figure 1.22This figure diagrams a rat liver
cell, a typical higher animal cell in which the
characteristic features of animal cells are
evident, such as a nucleus, nucleolus,
mitochondria, Golgi bodies, lysosomes, and
endoplasmic reticulum (ER). Microtubules and the
network of filaments constituting the
cytoskeleton are also depicted. (photos, top,
Dwight R. Kuhn/Visuals Unlimited middle, D.W.
Fawcett/Visuals Unlimited bottom, Keith
Porter/Photo Researchers, Inc.)
34
Figure 1.23This figure diagrams a cell in the
leaf of a higher plant. The cell wall, membrane,
nucleus, chloroplasts, mitochondria, vacuole, ER,
and other characteristic features are shown.
(photos, top, middle, Dr. Dennis
Kunkel/Phototake, NYC bottom, Biophoto
Associates)  
35
Figure 1.25The virus life cycle. Viruses are
mobile bits of genetic information encapsulated
in a protein coat. The genetic material can be
either DNA or RNA. Once this genetic material
gains entry to its host cell, it takes over the
host machinery for macromolecular synthesis and
subverts it to the synthesis of viral-specific
nucleic acids and proteins. These virus
components are then assembled into mature virus
particles that are released from the cell. Often,
this parasitic cycle of virus infection leads to
cell death and disease.
36
Chapter 2
  • Water The Medium of Life
  • Biochemistry
  • by
  • Reginald Garrett and Charles Grisham

37
Essential Question
  • What are the properties of water that render it
    so suited to its role as the medium of life?

38
Outline
  • What Are the Properties of Water?
  • What is pH?
  • What Are Buffers, and What Do They Do?
  • Does Water Have a Unique Role in the Fitness of
    the Environment?

39
2.1 - What Are the Properties of Water?
  • High b.p., m.p., heat of vaporization, surface
    tension
  • Bent structure makes it polar (dipole)
  • Non-tetrahedral bond angles -gt low density of ice
  • H-bond donor and acceptor, cooperative
  • Potential to form four H-bonds per water

40
Figure 2.1 The structure of water. Two lobes of
negative charge formed by the lone-pair electrons
of the oxygen atom lie above and below the plane
of the diagram. This electron density contributes
substantially to the large dipole moment and
polarizability of the water molecule. The dipole
moment of water corresponds to the OH bonds
having 33 ionic character. Note that the HOH
angle is 104.3, not 109, the angular value
found in molecules with tetrahedral symmetry,
such as CH4. Many of the important properties of
water derive from this angular value, such as the
decreased density of its crystalline state, ice.
(The dipole moment in this figure points in the
direction from negative to positive, the
convention used by physicists and physical
chemists organic chemists draw it pointing in
the opposite direction.)
41
Comparison of Ice and Water
  • Issues H-bonds and Motion
  • Ice 4 H-bonds per water molecule (open lattice
    structure, holds molecules apart, lt density)
  • Water 2.3 H-bonds per water molecule (random
    network)
  • Ice H-bond lifetime - about 10 microsec
  • Water H-bond lifetime - about 10 psec
  • (10 psec 0.00000000001 sec)
  • Thats "one times ten to the minus eleven second"!

42
Figure 2.2 The structure of normal ice. The
hydrogen bonds in ice form a three-dimensional
network. The smallest number of H2O molecules in
any closed circuit of H-bonded molecules is six,
so this structure bears the name hexagonal ice.
Covalent bonds are represented as solid lines,
whereas hydrogen bonds are shown as dashed lines.
The directional preference of H bonds leads to a
rather open lattice structure for crystalline
water and, consequently, a low density for the
solid state. The distance between neighboring
oxygen atoms linked by a hydrogen bond is
0.274nm. Since the covalent H-O bond is 0.095nm,
the H-O hydrogen bond length in ice is 0.18 nm.
43
Figure 2.3 The fluid network of H bonds linking
water molecules in the liquid state. It is
revealing to note that, in 10 psec, a photon of
light (which travels at 3 x 108 m/sec) would move
a distance of only 0.003m.
44
The Solvent Properties of Water Derive from Its
Polar Nature
  • Ions are always hydrated in water and carry
    around a "hydration shell"
  • Water forms H-bonds with polar, nonionic solutes
    (eg sugars)
  • Hydrophobic interactions with non-polar solutes-gt
    H bond network forms a clathrate structure around
    the solute molecule

45
Figure 2.4 Hydration shells surrounding ions in
solution. Water molecules orient so that the
electrical charge on the ion is sequestered by
the water dipole. For positive ions (cations),
the partially negative oxygen atom of H2O is
toward the ion in solution. Negatively charged
ions (anions) attract the partially positive
hydrogen atoms of water in creating their
hydration shells.
46
Hydrophobic Interactions
  • A nonpolar solute "organizes" water
  • The H-bond network of water reorganizes to
    accommodate the nonpolar solute
  • This is an increase in "order" of water
  • This is a decrease in ENTROPY

47
Figure 2.5 Formation of a clathrate structure by
water molecules surrounding a hydrophobic solute.
48
Amphiphilic Molecules
  • Also called "amphipathic
  • Refers to molecules that contain both polar and
    nonpolar groups
  • Equivalently - to molecules that are attracted to
    both polar and nonpolar environments
  • Good examples - fatty acids

49
Figure 2.6 An amphiphilic molecule sodium
palmitate. Amphiphilic molecules are frequently
symbolized by a ball and zig-zag line structure
,              ,where the ballrepresents the
hydrophilic polar head andthe zig-zag represents
the nonpolar hydrophobic hydrocarbon tail.
50
Figure 2.7 Micelle formation by amphiphilic
molecules in aqueous solution. Negatively charged
carboxylate head groups orient to the micelle
surface and interact with the polar H2O molecules
via H bonding. The nonpolar hydrocarbon tails
cluster in the interior of the spherical micelle,
driven by hydrophobic exclusion from the solvent
and the formation of favorable van der Waals
interactions. Because of their negatively charged
surfaces, neighboring micelles repel one another
and thereby maintain a relative stability in
solution.
51
Colligative Properties
  • Set of changes in water introduced by solutes
  • Solutes give order and restrict dynamic interplay
    between H2O molecules
  • Makes it more difficult for water to freeze
  • Eg 1M solution at 1 atm pressure lt freezing by
    1.86?C, gt boiling by 0.543?C, lowers vapour
    pressure, osmotic pressure relative to water22.4
    atm (at 25?C)

52
Figure 2.8 The osmotic pressure of a 1 molal (m)
solution is equal to 22.4 atmospheres of
pressure. (a) If a nonpermeant solute is
separated from pure water by a semipermeable
membrane through which H2O passes freely, (b)
water molecules enter the solution (osmosis) and
the height of the solution column in the tube
rises. The pressure necessary to push water back
through the membrane at a rate exactly equaled by
the water influx is the osmotic pressure of the
solution. (c) For a 1 m solution, this force is
equal to 22.4 atm of pressure. Osmotic pressure
is directly proportional to the concentration of
the nonpermeant solute.
53
Figure 2.9 The ionization of water.
54
Figure 2.10 The hydration of H3O (hydronium
ions). Solid lines denote covalent bonds dashed
lines represent the H bonds formed between the
hydronium ion and its waters of hydration.
55
2.2 What is pH?
  • The pH Scale
  • A convenient means of writing small
    concentrations
  • pH -log10 H
  • Sørensen (Denmark)
  • If H 1 x 10 -7 M
  • Then pH 7
  • Neutral H OH-

56
Dissociation of Weak Electrolytes
  • Electrolyte-gt generates ions in solution
  • Consider a weak acid, HA
  • The acid dissociation constant is given by
  • HA ? H A-
  • Ka H A -
  • ____________________ HA

57
The Henderson-Hasselbalch Equation
  • Know this! You'll use it constantly.
  • For any acid HA, the relationship between the
    pKa, the concentrations existing at equilibrium
    and the solution pH is given by
  • pH pKa log10 A
  • HA

58
Figure 2.11 The titration curve for acetic acid.
Note that the titration curve is relatively flat
at pH values near the pKa.  In other words, the
pH changes relatively little as OH- is added in
this region of the titration curve.
59
Consider the Dissociation of Acetic Acid
60
Consider the Dissociation of Acetic Acid
  • Another case....
  • What happens if exactly 0.5 eq of base is added
    to a solution of the fully protonated acetic
    acid?
  • With 0.5 eq OH added
  • pH pKa log10 0.5
  • 0.5
  • pH 4.76 0
  • pH 4.76 pKa

61
Consider the Dissociation of Acetic Acid
A final case to consider.... What is the pH if
0.9 eq of base is added to a solution of the
fully protonated acid? With 0.9 eq OH
added pH pKa log10 0.9

0.1 pH 4.76 0.95 pH 5.71
62
Figure 2.12The titration curves of several weak
electrolytes acetic acid, imidazole, and
ammonium. Note that the shape of these different
curves is identical. Only their position along
the pH scale is displaced, in accordance with
their respective affinities for H ions, as
reflected in their differing pKa values.
63
Figure 2.13 The titration curve for phosphoric
acid. The chemical formulas show the prevailing
ionic species present at various pH values.
Phosphoric acid (H3PO4) has three titratable
hydrogens and therefore three midpoints are seen
at pH 2.15 (pK1), pH 7.20 (pK2), and pH 12.4
(pK3).
64
2.3 What Are Buffers, and What Do They Do?
  • Buffers are solutions that resist changes in pH
    as acid and base are added
  • Most buffers consist of a weak acid and its
    conjugate base
  • Note in Figure 2.14 how the plot of pH versus
    base added is flat near the pKa
  • Buffers can only be used reliably within a pH
    unit of their pKa

65
Figure 2.14 A buffer system consists of a weak
acid, HA, and its conjugate base, A-. The pH
varies only slightly in the region of the
titration curve where HA A-. The unshaded
box denotes this area of greatest buffering
capacity. Buffer action when HA and A- are both
available in sufficient concentration, the
solution can absorb input of either H or OH-,
and pH is maintained essentially constant.
66
Figure 2.15 pH versus enzymatic activity. The
activity of enzymes is very sensitive to pH. The
pH optimum of an enzyme is one of its most
important characteristics. Pepsin is a
protein-digesting enzyme active in the gastric
fluid. Trypsin is also a proteolytic enzyme, but
it acts in the more alkaline milieu of the small
intestine. Lysozyme digests the cell walls of
bacteria it is found in tears.
67
Figure 2.17 The pKa values and pH range of some
Good buffers.
68
2.4 Does Water Have a Unique Role in the Fitness
of the Environment?
  • Vital for life
  • Powerful solvent hydrophobic interactions
    important in biology
  • Medium for ionisation
  • Climatic buffer
  • Osmosis helps form shape and structure

69
Chapter 3
  • Thermodynamics of Biological Systems
  • Biochemistry
  • by
  • Reginald Garrett and Charles Grisham

70
Essential Question
  • Living things require energy
  • What are the laws and principles of
    thermodynamics that allow us to describe the
    flows and interchanges of heat, energy, and
    matter in biochemical systems?

71
Outline
  • What Are the Basic Concepts of Thermodynamics?
  • What Can Thermodynamic Parameters Tell Us About
    Biochemical Events?
  • What Is the Effect of pH on Standard-State Free
    Energies?
  • What Is the Effect of Concentration on Net Free
    Energy Changes?
  • Why Are Coupled Processes Important to Living
    Things?
  •  What Are the Characteristics of High-Energy
    Biomolecules ?
  • What Are the Complex Equilibria Involved in ATP
    Hydrolysis?

72
3.1 - What Are the Basic Concepts of
Thermodynamics?
  • The system the portion of the universe with
    which we are concerned
  • The surroundings everything else
  • Isolated system cannot exchange matter or energy
  • Closed system can exchange energy
  • Open system can exchange either or both

73
Figure 3.1 The characteristics of isolated,
closed, and open systems. Isolated systems
exchange neither matter nor energy with their
surroundings. Closed systems may exchange energy,
but not matter, with their surroundings. Open
systems may exchange either matter or energy with
the surroundings.
74
The First LawThe Total Energy of an Isolated
System is Conserved
  • E (or U) is the internal energy - a function that
    keeps track of heat transfer and work expenditure
    in the system
  • E is heat exchanged at constant volume
  • E is independent of path
  • E2 - E1 ?E q w
  • q is heat absorbed BY the system
  • w is work done ON the system

75
Work
  • Mechanical work -gt movement through some distance
    caused by an application of force
  • Many biological examples
  • w-P ?V (Ppressure, V volume)
  • ?E, w and q have same unit cal (kcal) or Joule
  • If during mechanical work, no change in volume
    occurs, then
  • ?Eq

76
Enthalpy
  • A function for constant pressure
  • H E PV
  • If P is constant, ?H q
  • ?H is the heat absorbed at constant P
  • Volume is approx. constant for biochemical
    reactions (in solution)
  • So ?H is approx. same as ?E
  • ?H?, ?E? -gt standard state

77
Calorimeter can measure Heat (?H)
Figure 3.2 Diagram of a calorimeter. The reaction
vessel is completely submerged in a water bath.
The heat evolved by a reaction is determined by
measuring the rise in temperature of the water
bath. 
78
The Second Law Systems Tend Toward Disorder and
Randomness
  • Systems tend to proceed from ordered to
    disordered states
  • The entropy change for (system surroundings) is
    unchanged in reversible processes and positive
    for irreversible processes
  • All processes proceed toward equilibrium - i.e.,
    minimum potential energy

79
Entropy
  • A measure of disorder
  • An ordered state is low entropy
  • A disordered state is high entropy
  • dSreversible dq/T
  • Where dSreversible is ?S (entropy change) in a
    reversible system, q is heat transferred, T is
    temp

80
The Third LawWhy Is Absolute Zero So Important?
  • The entropy of any crystalline, perfectly ordered
    substance must approach zero as the temperature
    approaches 0 K
  • At T 0 K (abs 0, -273?C), entropy is exactly
    zero
  • Heat capacity (Cp) -gt at constrant pressure, the
    amt of heat 1 mole can store as temp is raised by
    1 degree
  • For a constant pressure process
  • Cp dH/dT

81
Free Energy
  • Hypothetical quantity - allows chemists to asses
    whether reactions will occur
  • Gibbs Free Energy -gt G H - TS
  • For any process at constant P and T
  • ?G ?H - T?S
  • If ?G 0, reaction is at equilibrium
  • If ?G lt 0, reaction proceeds as written
    (exergonic)
  • If ?G gt 0, reaction proceeds in reverse direction
    (endergonic)

82
?G versus ?Go
  • How can we calculate the free energy change for
    rxns not at standard state?
  • Consider a reaction A B ? C D
  • Then
  • ?G ?Go RT ln (CD/AB)
  • At equilibrium, ?G 0 and (CD/AB) Keq
    so
  • ?G? -RT ln Keq

83
Figure 3.4 The dependence of G on temperature
for the denaturation of chymotrypsinogen.at pH3
(Adapted from Brandts, J. F., 1964. The
thermodynamics of protein denaturation. I. The
denaturation of chymotrypsinogen. Journal of the
American Chemical Society 8642914301.)
84
Figure 3.5 The dependence of S on temperature
at pH 3 for the denaturation of chymotrypsinogen.
(Adapted from Brandts, J. F., 1964. The
thermodynamics of protein denaturation. I. The
denaturation of chymotrypsinogen. Journal of the
American Chemical Society 8642914301.)
85
3.2 What Can Thermodynamic Parameters Tell Us
About Biochemical Events?
  • Comparison of multiple thermodynamic parameters
    is useful (ie ?H, ?S, ?G)
  • Can be used to predict whether or not a reaction
    will proceed as written

86
3.3 What Is the Effect of pH on Standard-State
Free Energies?
  • Need a modified standard state () for reactions
    in which H are produced
  • ?G ?G? RT ln H

87
3.4 What Is the Effect of Concentration on
Net Free Energy Changes?
  • ?G can differ from the standard state value if
    the concentrations of the reactants and products
    differ widely from unit activity (1M)
  • For AB ? CD
  • ?G ?G ? RT ln (C D / A B)

88
3.5 Why Are Coupled Processes Important to Living
Things?
  • Many important cellular reactions must run
    against their thermodynamic potential ie in
    direction of positive ?G
  • Do this by coupling to a favorable reaction

89
3.6 What Are the Characteristics of High-Energy
Biomolecules ?
  • Energy Transfer -gt A Crucial Biological Need
  • Energy acquired from sunlight or food must be
    used to drive endergonic (energy-requiring)
    processes in the organism
  • Two classes of biomolecules do this
  • Reduced coenzymes (NADH, FADH2)
  • High-energy phosphate compounds - free energy of
    hydrolysis more negative (-25 kJ/mol).

90
High-Energy Biomolecules
  • Study Table 3.3!
  • Note what's high - PEP (?G?-62.2 kJ/mol)
  • Note what's low - sugar phosphates
  • (?G? -13.9 kJ/mol)
  • Note what's in between - ATP
  • (?G?-30.5 kJ/mol)
  • Note difference (Figure 3.8) between overall free
    energy change and the energy of activation for
    phosphoryl-group transfer

91
Figure 3.8 The activation energies for phosphoryl
group-transfer reactions (200 to 400 kJ/mol) are
substantially larger than the free energy of
hydrolysis of ATP (-30.5 kJ/mol).
92
ATP
  • An Intermediate Energy Shuttle Device
  • PEP and 1,3-BPG are created in the course of
    glucose breakdown
  • Their energy (and phosphates) are transferred to
    ADP to form ATP
  • But ATP is only a transient energy carrier - it
    quickly passes its energy to a host of
    energy-requiring processes

93
Figure 3.9The triphosphate chain of ATP contains
two pyrophosphate linkages, both of which release
large amounts of energy upon hydrolysis.
94
Phosphoric Acid Anhydrides
  • Why ATP does what it does!
  • ADP and ATP are examples of phosphoric acid
    anhydrides
  • Large negative free energy change on hydrolysis
    is due to
  • electrostatic repulsion
  • stabilization of products by ionization and
    resonance
  • entropy factors due to hydrolysis and subsequent
    ionisation

95
Figure 3.10 (a) Electrostatic repulsion between
adjacent partial positive charges (on carbon and
phosphorus, respectively) is relieved upon
hydrolysis of the anhydride bonds of acetic
anhydride and phosphoric anhydrides. The
predominant form of pyrophosphate at pH values
between 6.7 and 9.4 is shown. (b) The competing
resonances of acetic anhydride and the
simultaneous resonance forms of the hydrolysis
product, acetate.
(a)
(b)
96
Figure 3.11Hydrolysis of ATP to ADP (and/or of
ADP to AMP) leads to relief of electrostatic
repulsion.
97
Phosphoric-Carboxylic Anhydrides
  • These mixed anhydrides - also called acyl
    phosphates - are very energy-rich
  • Acetyl-phosphate ?G -43.3 kJ/mol
  • 1,3-BPG ?G -49.6 kJ/mol
  • Bond strain, electrostatics, and resonance are
    responsible

98
Figure 3.12 The hydrolysis reactions of acetyl
phosphate and 1,3-bisphosphoglycerate.
99
Enol Phosphates
  • Phosphoenolpyruvate (PEP) has the largest free
    energy of hydrolysis of any biomolecule
  • Formed by dehydration of 2-phospho-glycerate
  • Hydrolysis of PEP yields the enol form of
    pyruvate - and tautomerization (bond
    rearrangement) to the keto form is very favorable

100
enolase
Pyruvate kinase
Figure 3.13 Phosphoenolpyruvate (PEP) is
produced by the enolase reaction (in glycolysis)
and in turn drives the phosphorylation of ADP to
form ATP in the pyruvate kinase reaction. 
101
Figure 3.14 Hydrolysis and the subsequent
tautomerization account for the very large G
of PEP.
102
3.7 What Are the Complex Equilibria Involved in
ATP Hydrolysis?
  • ATP has five dissociable protons
  • pKa values range from 0-1 to 6.95
  • Free energy of hydrolysis of ATP is relatively
    constant from pH 1 to 6, but rises steeply at
    high pH
  • Since most biological reactions occur near pH 7,
    this variation is usually of little consequence
  • Metal ions affect the ?G? of hydrolysis of ATP

103
Figure 3.16 The pH dependence of the free energy
of hydrolysis of ATP. Because pH varies only
slightly in biological environments, the effect
on ?G is usually small.
104
Figure 3.17The free energy of hydrolysis of ATP
as a function of total Mg2 ion concentration at
38C and pH 7.0. (Adapted from Gwynn, R. W., and
Veech, R. L., 1973. The equilibrium constants of
the adenosine triphosphate hydrolysis and the
adenosine triphosphate-citrate lyase reactions.
Journal of Biological Chemistry 24869666972.) 
105
The Effect of Concentration
  • Free energy changes are concentration dependent
  • We will use the value of -30.5 kJ/mol for the
    standard free energy of hydrolysis of ATP
  • But at non-standard-state conditions (in a cell,
    for example), the ?G is different because ATP,
    ADP and Pi differ!
  • ?G ?G ? RT ln (?ADP ?Pi) / ?ATP
  • In typical cells, the free energy change for ATP
    hydrolysis is typically -50 kJ/mol

106
Figure 3.18The free energy of hydrolysis of ATP
as a function of concentration at 38C, pH 7.0.
The concentrations C of ATP, ADP, and Pi
assumed to be equal.
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