Transcript Document

BIOCHEMISTRY A GMS BI 555/755
Course manager: Konstantin Kandror
See Courseinfo (www.bumc.bu.edu)
Every Monday and Wednesday, 10:00 to 12:00 Room
R103
Every other Friday, 11:30 to 12:30 Room R103
(Computer room L1110 on September 11)
09/09/09
GMS BI 555/755 Lecture 1.
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GMS BI 555/755: Biochemistry
Resources
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Instructor for lectures 1-3: Joseph Zaia, X8-6762, [email protected]
Course notes, print them out BEFORE the lecture.
Biochemistry, 6th ed. Berg, Tymoczko, Stryer (hard copy)
Engines, Energy and Entropy John Fenn.
National Center for Biotechnology Information (on-line text books)
– www.ncbi.nih.gov
– Stryer 5th ed
– Molecular Cell Biology Lodish et al.
– Molecular Biology of the Cell, 4th ed, Alberts et al.
– Many other texts
Wikipedia for some basic concepts
Discussion groups
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GMS BI 555/755 Lecture 1: . Bioenergetics: Molecular Interactions and
Thermodynamics within Biological systems
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Readings:
Biochemistry, Berg et al, 6th ed, pp 6-17.
Supplementary: Molecular Cell Biology, Lodish, el al, Chap. 2, “Chemical
Foundations” on-line at www.ncbi.nih.gov
1. Bioenergetics
a. First Law of Thermodynamics
b. Second law of Thermodynamics
c. Free energy
d. Flow of energy in living systems
2. Non-covalent interactions
a. Covalent bonds
b. Hydrogen bonding
c. Ionic interactions
d. Van der Waal’s forces
e. Hydrophobic effect
f. Protein specificity
3. Weak acids and bases in biological systems.
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The cell as a system: what principles underlie it’s organization and energetics
and, by extension, life itself?
Sir Arthur Eddington in The Nature of the Physical World:
“The second law of thermodynamics holds, I think, the supreme position among the laws of Nature. If someone
points out to you that your pet theory of the universe is in disagreement with Maxwell’s equations - then so much
the worse for Maxwell’s equations. If it is found to be contradicted by observation, well, these experimentalists do
bungle things sometimes. But if your theory is found to be against the second law of thermodynamics I can give
you no hope; there is nothing for it but to collapse in deepest humiliation.”
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The Universe as a thermodynamic
system: energy/matter is neither created
nor destroyed but converts between
different forms.
• System: matter within a defined region
of space
A system is open if
it can exchange energy and
matter with its surroundings,
closed if it can exchange
energy but not matter, and
isolated if it can exchange
neither energy nor matter.
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Types of energy: Potential and Kinetic
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Kinetic energy:
– Thermal energy - energy of motion, heat energy. Thermal regulation in
warm blooded animals used for maintenance of ideal temperature for
enzymatic reactions; cells do not use heat to do work (exception of
thermophilic bacteria)
– Radiant energy – energy of light (photons) may be converted to heat
energy. May be used to generate chemical bond energy in
photosynthesis.
Potential energy: stored energy
– Energy stored in chemical bonds: (i.e. C6H12O6,) used to drive
metabolism
– Energy of concentration gradients sequestered by biological
membranes. Ions, nutrients, waste, protons, etc. maintained in
compartment specific concentrations.
– Electric potential: all cells maintain a gradient of electrical charge across
their plasma membrane.
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Energy as heat: transferred
to surroundings to cause
random molecular motions
due to increase in
temperature. Dissipation.
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Energy as work: directed
energy that has a focused
effect.
• Lifting a weight
• Transporting a nutrient across
a membrane
• Synthesizing a protein
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Biochemical thermodynamics
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1st Law: The amount of energy in a system plus surroundings is constant.
– Energy cannot be created or destroyed but can be converted between
different forms. Mass is a type of energy.
– Any energy released in the formation of chemical bonds must be used
to break other chemical bonds, converted to heat or light, or stored in
another form
2nd Law: The amount of entropy (randomness) in a system plus
surroundings always increases.
– The universe is expanding. The amount of disorder in the universe
resulting from a (bio) chemical process always increases.
– In order to create biological order, an organism must expend energy,
but the net entropy in the universe increases.
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Metabolism: An organism must continuously maintain order with respect to
its environment. It must create order within itself at the expense of that of
its environment (waste). It must free itself, through metabolism, from the
disorder (decay) it must create in order to exist (1930s→1950s).
Gibbs free energy (Gibbs’ Law, 1870’s):
ΔG = ΔH-TΔS
H = enthalpy (chemical bond energy)
S = entropy (randomness)
ΔG <0 for spontaneous processes
ΔG >0 reverse reaction spontaneous
ΔG = 0 both forward and reverse reactions occur at equal rates
J. Willard Gibbs
Chemical Physicist
Gibbs free energy explains the driving
force behind (biological and nonbiological) chemical change.
• ΔS proportional to heat transferred from
system to surroundings
• ΔS inversely proportional to temperature
Thermodynamically, the amount of energy capable of doing work during a chemical reaction is measured
quantitatively by the change in the Gibbs free energy.
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ΔG = ΔH - TΔS
The total entropy change is given by the expression
Substituting equation 1 into equation 2 yields
Multiplying by -T gives
The function -TΔS has units of energy and is referred to as free energy or
Gibbs free energy, after Josiah Willard Gibbs, who developed this function
in 1878:
The free-energy change, ΔG, will be used throughout to describe the
energetics of biochemical reactions.
Recall that the Second Law of Thermodynamics states that, for a reaction
to be spontaneous, the entropy of the universe must increase. Examination
of equation 3 shows that the total entropy will increase if and only if
Rearranging gives TΔSsystem > ΔH, or entropy will increase if and only if
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ΔG = ΔH-TΔS
Mixing of two solutes by
removing a barrier
ΔH
ΔS
ΔG
2H2 + O2 → 2H2O
ΔH
ΔS
ΔG
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ΔG = ΔH-TΔS
Entropy (S): randomness or disorder in a system. Entropy increases as a system
becomes more disordered and decreases as it becomes more ordered.
Chambers separated by
permeable membrane
H2O
1M
glucose
H2O
ΔH = 0
ΔS = >0
H2O
0.5 M
glucose
nutrients
A cell
H2O
0.5 M
glucose
ΔH = 0
ΔS = <0
Entropically favorable because
glucose is allowed to diffuse
over the maximum available
volume achieving an increase
in randomness
Entropically unfavorable:
increase in order required
to maintain homeostasis
of nutrients inside cell
and waste outside cell.
nutrients
waste
waste
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Formation of the DNA double helix
• An exothermic reaction (ΔH < 0) that increases
entropy (ΔS > 0) occurs spontaneously (ΔG < 0).
• An endothermic reaction (ΔH > 0) will occur
spontaneously if ΔS increases enough so that the T
ΔS term can overcome the positive ΔH.
• If the conversion of reactants into products results in
no change in free energy (ΔG = 0), then the system is
at equilibrium; that is, any conversion of reactants to
products is balanced by an equal conversion of
products to reactants.
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ΔS <0 (more ordered)
ΔH <0 (heat released)
ΔG <0 (spontaneous process, heat
released through formation
of multiple H-bonds)
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ΔG = ΔH-TΔS
Enthalpy ΔH = negative when the reactants have more chemical bond
energy than the products. Excess energy is given off, usually as heat.
• Combustion is an exothermic process:
2C8H18 + 25O2 → 16CO2 + 18H2O + energy
• Metabolism is an exothermic process:
C6H12O6 + 6O2 → 6CO2 + 6H2O + energy
ΔH = positive when the products have more chemical bond energy than the
reactants. When enthalpy is positive, energy must be added to chemical
bonds for the reaction to occur.
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What does ΔG tell us about a biochemical reaction?
• A reaction occurs spontaneously only if ΔG <0 (exothermic or exergonic)
• A system is at equilibrium when ΔG = 0 (steady state)
• A reaction in which ΔG>0 requires input of energy (endothermic or endergonic)
• ΔG = Gproducts – Greacants, it does not depend on the reaction path
(i.e. combustion vs metabolism)
• ΔG provides no information on the reaction rate
Activation energy
Greactants
ΔG
Gproducts
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Parameters affecting ΔG of a reaction: temperature, pressure, concentrations of reactants and products
Most biological reactions differ from standard conditions, particularly in the concentrations of
reactants. However, we can estimate free-energy changes for different temperatures and initial
concentrations, using the equation
where R is the gas constant of 1.987 cal/(degree · mol), T is the temperature (in degrees Kelvin), and Q is the initial ratio of
products to reactants. For the interconversion of glyceraldehyde 3-phosphate (G3P) and dihydroxyacetone phosphate
(DHAP)
Q = [DHAP]/[G3P] and ΔG°′ = −1840 cal/mol. Equation 2-8 for ΔG then becomes
ΔG may be calculated for any set of concentrations of DHAP and G3P.
•If the initial concentrations of both DHAP and G3P are 1 M, then ΔG = ΔG°′ = −1840 cal/mol, because RT ln 1 = 0.
•If [DHAP] = 0.1 M and [G3P] = 0.001 M, then Q = 0.1/0.001 = 100, and
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Molecular Cell Biology, Lodish et al.
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An Unfavorable Chemical Reaction Can Proceed If It Is Coupled with an Energetically
Favorable Reaction
Many chemical reactions in cells are energetically unfavorable (ΔG > 0) and will not proceed
spontaneously. One example is the synthesis of small peptides (e.g., glycylalanine) or proteins
from amino acids. Cells are able to carry out a reaction that has a positive ΔG by coupling it to a
reaction that has a negative ΔG of larger magnitude, so that the sum of the two reactions has a
negative ΔG. Suppose that the reaction
has a ΔG°′ of +5 kcal/mol and that the reaction
has a ΔG°′ of −10 kcal/mol. In the absence of the second reaction, there would be much more A
than B at equilibrium. The occurrence of the second process, by which X becomes Y + Z, changes
that outcome: because it is such a favorable reaction, it will pull the first process toward the
formation of B and the consumption of A.
The ΔG°′ of the overall reaction will be the sum of the ΔG°′ values of each of the two partial
reactions:
The overall reaction releases energy. In cells, energetically unfavorable reactions of the type
A ⇌ B + X are often coupled to the hydrolysis of the compound adenosine triphosphate (ATP), a
reaction with a negative change in free energy (ΔG°′ = −7.3 kcal/mol), so that the overall reaction
has a negative ΔG°′.
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Molecular Cell Biology, Lodish et al.
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Ion transport using Na+ gradients
Lodish et al Molecular Cell Biology chap 15.
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Flow of energy in living systems
(controlled oxidation and reduction)
Diagram demonstrating the flow of energy in living organisms. Arrows point in the direction in which energy flows. We
focus only on the most common processes and do not include less ubiquitous ones, such as bioluminescence.
(Adapted from D.A. Harris, Bioenergetics at a glance, Blackwell Science, Oxford [1995].)
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The energy Required to Break Some Important Covalent Bonds Found in Biological Molecules
Type of Bond
Energy (kcal/mol)
SINGLE BOND
Type of Bond
Energy (kcal/mol)
DOUBLE BOND
O—H
110
C=O
170
H—H
104
C=N
147
P—O
100
C=C
146
C—H
99
P=O
120
C—O
84
C—C
83
TRIPLE BOND
S—H
81
C≡O
C—N
70
C—S
62
N—O
53
S—S
51
*
195
Note that double and triple bonds are stronger than single bonds.
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Chemical Bonds: Covalent
Types of covalent bonds in biochemistry: proteins
Amino acids
R1
O
R3
H2N
OH
OH
OH
H2N
O
H2N
O
R2
R1
O
R3
H
N
OH
H2N
N
H
O
R2
Covalent bonding involves sharing of
a pair of electrons in the form of
orbital overlap
O
Peptide: amide linked
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Chemical Bonds: Covalent
Types of covalent bonds in biochemistry: carbohydrates
OH OH
OH
O
HO
OH
HO
HO
OH
Gal
OH OH
HO
OH
O
HO
OH
NHAc
GlcNac
OH OH
OH
O
O
NHAc
OH
OH OH
OH
O
O
O
HO
O
O
NHAc
Polylactosamine
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Monosaccharides
(polyhydroxy aldehydes/ketones)
O
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OH
O
OH
O
HO
O
OH
NHAc
Glycosidic bond: Ether linkage
Non-covalent bonding 22
Non-covalent interactions
Hydrogen bonding:
O
H
H
O
O
H
H
H
H
O
H
+ 5 kcal/mol
H
• H-bond formation releases a small amount of energy
relative to the typical energies of a covalent bond.
• Average kinetic energies of molecules at 25C is 0.6
kcal/mol
• H-bonds are weak and transient in nature
• In combination, H-bond interactions become significant
for holding molecules together non-covalently
• H-bonds underly the unique chemical properties of
water (solvent of life).
• Much higher boiling point than similarly sized nonhydrogen bonding molecules
• Water is very dense; solid state has lower density
than liquid state (ice floats)
• Polar protic nature enabled evolution of
biomolecules
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Structure of ice, showing H-bonds
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Non-covalent interactions
Hydrogen bonding: general form
0.26-0.31 nm
(twice covalent bond length)
Donors, N-H and O-H bonds in biomolecules
Acceptors: O: and N: lone pairs. Numerous interactions possible in
biomolecules.
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Hydrogen bond formation in proteins and nucleic acids releases 1-2
kcal/mol, much lower than for bond formation among water
molecules. The reason for this is that in order for a new proteinwater H-bond to form, an H-bond with water must break.
Water-free microenvironments enable formation of H-bonds between polar groups
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Ionic interactions
Electronegativity values of main-group
elements in the periodic table. Atoms located
to the upper right tend to have high
electronegativity, fluorine being the most
electronegative. Elements with low
electronegativity values, such as the metals
lithium, sodium, and potassium, are often
called electropositive. The electronegativities
of several atoms abundant in biological
molecules differ enough that they form polar
covalent bonds (e.g., O—H, N—H) or ionic
bonds (e.g., Na+Cl−). Because the inert gases
(He, Ne, etc.) have complete outer shells of
electrons, they neither attract nor donate
electrons, rarely form covalent bonds, and
have no electronegativity values.
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E = (kq1q2)/Dr2
(Coulomb’s law)
k = proportionality constant
D = dielectric constant = 80 for water
r = distance
Ionic interactions
q1
•
•
r
q2
Atoms with large electronegativity differences do not share electrons by
forming covalent bonds. Rather, electrons are completely transferred,
forming a positive and a negative ion.
X▪ + Y▪ → X+ + YNa▪ + Cl▪ → Na+ + ClMany ionic compounds are readily soluble in water because of release of
energy when they become hydrated (energy of hydration); they become
surrounded by a shell of water molecules (hydration sphere)
Hδ+
Hδ+
H2O
solvation
Hδ+
(KCL lattice)
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Cl-
Oδ-
K+▪▪▪▪▪▪▪▪▪▪Oδ-
Hδ+
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Ionic interactions
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Energetics of solvation of ionic compounds is a balance between the energy
of the ionic lattice versus those of the hydrated ions
Na▪ + Cl▪ → Na+ + ClReadily soluble, energy of solvation greater than lattice energy
Ca3(PO4)2
→
Insoluble due to high lattice energy
•
Ions must lose their hydration shell as they pass through ion transport
channels in biological membranes. The energetics of these processes
influences ion selective transport.
• (Mg2+), six water molecules held tightly in
place by electrostatic interactions between
the two positive charges on the ion and the
partial negative charge on the oxygen of
each water molecule.
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Ionic interactions: H-bonds revisited
δ- δ+
δ+
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δ-
• Polarity and H-bonding capability (donating
and accepting) of water makes it an excellent
solvent for polar molecules
• It has a high (80) dielectric constant (D) and
weakens electrostatic interactions and Hbonds between other molecules by
competing for interactions
• Water free microenvironments exist in
biological systems to enable strong
interactions between polar molecules
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Van der Waal’s forces (aka London dispersion forces)
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Two atoms approaching one another induce transient dipoles resulting in a
weak, non-specific attractive force.
δ-
•
•
•
δ+ δ- δ+
Occur in all molecules, polar and non-polar
Responsible for cohesion of non-polar liquids (alkanes). Stronger
interactions (ionic, H-bonding) override van der Waal’s attractions.
Van der Waal’s contact: when the repulsive forces of a pair of atom electron
clouds is balanced by van der Waal’s attractive forces.
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Van der Waal’s forces
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Van der Waal’s energy is ~1kcal/mol, weaker than H-bonds, only
slightly higher than thermal energy
Van der Waal’s forces become significant for large molecules in
which domains with complementary shapes allow for many
contacts
Thus they are important driving forces behind the energetics of
intra- and inter-molecular protein interactions:
– Protein folding
– Multi-protein complexes (molecular machines)
– Enzyme-substrate binding
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Hydrophobic interactions
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Non polar molecules lack polar and/or ionic groups and do not form ionic or
H-bonding interactions with water molecules.
Water must form an ordered cage surrounding a hydrophobic molecule.
Thus there is an energetic cost to the dissolution of a hydrophobic molecule
in water causing hydrophobic molecules to remain associated with one
another (a droplet) or with a surface (vessel wall) rather than go into
solution.
Hydrophobic interactions are thus driven by entropic effects. There is a
strong tendency of hydrophobic molecules to interact with one another
rather than become solvated in water.
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Non-covalent interactions give rise to protein specificity
The binding of a
hypothetical pair of proteins
by two ionic bonds, one
hydrogen bond, and a large
combination of hydrophobic
and van der Waals
interactions. The structural
complementarity of the
surfaces of the two molecules
gives rise to this particular
combination of weak bonds
and hence to the specificity of
binding between the
molecules.
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A spectrum of macromolecular non-covalent interactions (binding partners) gives rise to biological activity
Multiple weak bonds stabilize specific associations between large molecules. (Left) In this hypothetical complex,
seven noncovalent bonds bind the two protein molecules A and B together, forming a stable complex. (Right) Because
only four noncovalent bonds can form between proteins A and C, this interaction may be too weak for the A-C complex
to exist in cells.
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Ionization of water
Water dissociates into hydronium (H3O+) and hydroxyl (OH-) ions. For simplicity, we refer to the hydronium ion
as a hydrogen ion (H+) and write the equilibrium as
The equilibrium constant Keq of this dissociation is given by
in which the terms in brackets denote molar concentrations. Because the concentration of water (55.5 M) is
changed little by ionization, expression 1 can be simplified to give
in which Kw is the ion product of water. At 25°C, Kw is 1.0 × 10-14.
Note that the concentrations of H+ and OH- are reciprocally related. If the concentration of H+ is high, then the
concentration of OH- must be low, and vice versa. For example, if [H+] = 10-2 M, then [OH-] = 10-12 M
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Definition of pH and pK
The pH of a solution is a measure of its concentration of H+. The pH is defined as
The ionization equilibrium of a weak acid is given by
The apparent equilibrium constant Ka for this ionization is
The pKa of an acid is defined as
Inspection of equation 4 shows that the pKa of an acid is the pH at which it is half dissociated, when [A-]=[HA].
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The Henderson-Hasselbalch Equation Relates pH and Keq of an Acid-Base System
Many molecules used by cells have multiple acidic or basic groups, each of which can
release or take up a proton. In the laboratory, it is often essential to know the precise state
of dissociation of each of these groups at various pH values. The dissociation of an acid
group HA, such as acetic acid (CH3COOH), is described by
The equilibrium constant Ka for this reaction is
By taking the logarithm of both sides and rearranging the result, we can derive a very
useful relation between the equilibrium constant and pH as follows:
or
Substituting pH for −log [H+] and pKa for −log Ka, we have
Henderson-Hasselbalch equation: pKa of any acid is equal to the pH at which half the
molecules are dissociated and half are neutral (undissociated).
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Buffers: weak acids that resist pH changes near their pKa
The titration curve of acetic acid
(CH3COOH). The pKa for the
dissociation of acetic acid to hydrogen
and acetate ions is 4.75. At this pH, half
the acid molecules are dissociated.
Because pH is measured on a
logarithmic scale, the solution changes
from 91 percent CH3COOH at pH 3.75 to
9 percent CH3COOH at pH 5.75. The
acid has maximum buffering capacity in
this pH range.
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The titration curve of
phosphoric acid (H3PO4).
This biologically ubiquitous
molecule has three hydrogen
atoms that dissociate at
different pH values; thus,
phosphoric acid has three pKa
values, as noted on the graph.
The shaded areas denote the
pH ranges — within one pH
unit of the three pKa values —
where the buffering capacity of
phosphoric acid is maximum.
In these regions the addition
of acid (or base) will cause the
least change in the pH.
Phosphodiester bond (pKa ~ 3)
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pH mediated DNA denaturation
• As pH is increased, the extent of deprotonatoin of the N-1 nitrogen atom of Guanine
increases. At pH > 9.7, the N-1 proton is no longer there to participate in base pairing through
H-bonding
• At pH <5, some H-bond acceptors become protonated and unable to participate in H-bonding,
disrupting the structure of DNA.
O
O
N
N
pKa = 9.7
NH
N
N
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N
N
NH2
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N
NH2
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Proteins in solution
pKa ~2
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pKa ~9
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