Transcript 4. Water (2)

1. The molecular logic of life (1)
2. Cells (2)
3. Biomolecules (2)
4. Water (2)
5. Amino acids…etc (2)
6. Protein structure (4); Protein action (2)
7. Protein function (4); Experimental
techniques (2)
8. Enzymes: catalysis and regulation (6)
midterm: review and Q&A (2); exam (2)
Chapter 2 Molecular Logic of
Life
Some Important Chemical Concepts
and Principles for Studying
Biochemistry
1. Living matter is composed mostly of the
lighter elements
1.1 The composition of living matter is strikingly
different from that of its physical environment (1810s)
1.2 The elements found in living organisms also exist in
nature (especially in seawater and atmosphere).
1.2.1 99% of the mass of living organisms are
made of H, O, N, and C.
1.2.2 H, O, N, and C are the lightest elements
capable of forming one, two, three, and four bonds (in
general, lightest elements form the strongest bonds).
1.2.3 The trace elements, although represent a
miniscule fraction in living organisms, all are
absolutely essential to life (Fe, Cu, Mn, Zn, I, Mg).
Jellyfish and the sea water
2. Carbon was selected as the key element for life
due to its versatile bonding capacity
2.1 Carbon accounts for more than one-half the cell dry
weight.
2.2 Each carbon atom can form very stable single bonds
with one, two, three, or four other carbon atoms, and
double or triple bonds can also be formed between two
carbon atoms.
2.3 Covalently linked carbon atoms can form linear
chains, branched chains, and cyclic and cagelike（笼形
的）structures.
2.4 To these carbon skeletons are added functional
groups conferring specific activities to the molecules.
2.4 Molecules containing covalently bonding carbon
backbones are called organic compounds (including
mainly alcohols, amines, aldehydes and ketones,
carboxylic acids, sulfhydryls, … etc. Most biomolecules
are organic compounds.
2.5 Carbon atoms have a characteristic tetrahedral
arrangement of their four single bonds. Carbon-carbon
single bonds have freedom of rotation, but not double nor
triple bonds.
2.6 No other chemical element has the capacity to form
molecules of such widely different sizes and shapes or
with such a variety of functional groups.
Filled outer electron shells are
more stable: covalent bonds by
sharing unpaired electrons between
two atoms.
Versatility of carbon in forming
covalent bonds
The end group of Arginine’s side chain
Histidine’s side chain group
Cysteine’s functional group
3. Organic biomolecules have three
dimensional structures
3.1 The central special feature of organic
compounds is not their compositions but the way
their atoms are combined, i.e., their structures
(realized between 1820s-1860s). Corollary: two
substances may show the same chemical formula
but be physically and chemically different
materials (different structures and functions).
Covalent bond
length = sum of
covalent radii
Fisher
Ball-and-stick
Space filling
Convention used in organic chemistry for configuration
Light absorbing pigment in rhodopsin
an integral membrane protein
Convention from organic chemistry
(R,R) isomer
(S,S) isomer
High
energy
Low
energy
3.2 Compounds of carbon can often exist in two or more
chemically indistinguishable stereoisomers (having the
same formula, the same joining/bonds between atoms,
but different arrangements in 3D space).
3.2.1 Four different functional groups can be
bonded to a carbon in two different spatial arrangements,
making two stereoisomers. Such a carbon is called
asymmetric carbon or chiral carbon, or a carbon with
chirality.
3.2.2 The two stereoisomers of a chiral carbon are
nonsuperimposable(不能重叠的） mirror images of each
other. They are called enantiomers to each other (like a
pair of right and left hands).
3.2.3 The two enantiomers have identical chemical
properties but are different in a physical property called
optical activity. One rotates the plane of the planepolarized light to the left, the other to the right.
3.3 Configuration and conformation define the
different aspects of the three dimensional structure of
biomolecules.
3.3.1 Configuration defines the spatial
arrangement of the groups attached to an asymmetric
carbon or two double-bonded carbon atoms.
Configurational isomers can not be interconverted
without breaking one or more covalent bonds.
3.3.2 Conformation refers to the numerous
possible spatial arrangement of atoms due to free
rotations around single bonds. Conformational isomers
can be interconverted without breaking covalent bonds.
3.4 The three dimensional structure of organic
molecules can be illustrated by different ways.
Perspective model specifies the 3-D structure,
only applying to small molecules. Ball-and-stick
model shows the relative bond angles and lengths.
Skeleton model shows only the framework of a
molecule. Space-filling model shows the atoms in
proportional to their van der Waals radius, with
more realistic volumes and surfaces.
4. Interactions between biomolecules are
stereospecific
4.1 Between pairs of enantiomers, usually only one form
is biologically active. For example, only L-amino acids
(S) are found in proteins and only D-glucose (R) is
biologically active.
4.2 Usually only one chiral form of a biomolecule is
generated in living cells due to enzyme specificity (1975
Nobel Prize was for this discovery).
4.2.1 In (organic) chemical synthesis the two
enantiomers are usually synthesized in equal amounts.
4.2.2 Two stereoisomers may have totally opposite
biological effects (e.g., aspartame（天冬酰苯丙氨酸甲
酯）, a sugar substitute, and its stereoisomer, bitter).
It has no symmetry but complementary.
Conformational flexibility of a protein
香菜
绿薄荷
Chewing gums
Neutral sweet
(commercial name)
5. Chemical reactions between biomolecules are
the broken and formation of covalent bonds
5.1 Covalent bonds are formed by sharing the outer
shell electrons between two atoms.
5.1.1 Atoms tend to attain “filled-shell” conditions
by gaining, losing, or sharing electrons.
5.1.2 When two atoms have the same
electronegativity (electron affinity，电负性), the
covalent bond is nonpolar; when different, the bond is
polarized.
5.1.3 The strength of a bond is expressed as bond
energy (in joules or calories): that is, the amount of
energy required to break the bond; or gained by the
surroundings when the two atoms form the bond.
5.2 Many biochemical reactions occur when
nucleophiles (nuclear-seeking groups or atoms,
rich in electrons，亲核试剂) attack electrophiles.
5.2.1 Functional groups containing O, N,
and S are important biological nucleophiles.
5.2.2 Positively charged cations (H+, metals)
often act as biological electrophiles.
5.2.3 A carbon atom can act as both
(carbonium ion or carbon anion) depending on
the bonds and functional groups surrounding it.
Oxidation states of carbons in biomolecules
Most
reduced
Glucose,
source of
electrons
for
metabolism
Most
oxidized
Oxidation, losing electrons
Reduction, gaining electrons
An oxidation-reduction reaction
Carbonium ion
All have “protruding”
lone electron pairs
B1吸引一个质子
这样形成了一个C=C双键
来自羰基的电子促使O与H
离子形成了O－H键
An isomerization reaction
Four resonance structures of double bonds
More accurate representation using hybrid orbitals
SN2 reaction, the leaving group is ADP.
Transition state which exists transiently
Condensation
（缩和）
reactions
Elimination of a water molecule
tRNA is a better leaving group for removal during condensation.
In this reaction, an amino acid is first activated by tRNA.
A hydrolysis reaction of a peptide bond, a nucleophilic attack
on the carbonyl carbon by a water molecule
6. Free energy change (G) determines
whether a biochemical reaction can occur
spontaneously
6.1 The total energy of the universe (closed)
remains constant in any process (First law of
thermodynamics). Energy is generally defined as
the capacity to do work.
6.2 A process (e.g., a chemical reaction) can occur
spontaneously only if the sum of the entropies
(randomness) of the system (open) and its
surrounding (the universe) increases (Second law
of thermodynamics).
6.3 Free energy change (G) of the system is a
composite function that is a direct measure of the
entropy change of the universe:
G=H-TS
Each of the parameters are of the system
(with properties of the surrounding not included).
 H is the change in enthalpy (heat transferred to
the surrounding) of the system. T is the absolute
temperature of the system.  S is the entropy
change of the system.
6.4 The G of a reaction depends on the
difference of the free energies of the product and
the reactant, which are determined by their
structures and concentrations only. Therefore,
G is independent of the path (molecular
mechanism) of the transformation from the
reactant to the product.
6.5 The G provides no information about the
reaction rate. That is, a spontaneous reaction
may occur at a nonperceptible（不可感知的）
rate (never occur)!
6.6 the standard free energy Go’ and mass action
ln([product]/[reactant]):
G = Go’ + RT ln([product]/[reactant])
Go’ is defined as the free energy change of the
reaction under the standard condition, that is, when
the initial concentration of each component is 1.0 M,
the pH is 7.0, and the temperature is 25 C.
6.7 Go’ is related to the equilibrium constant
(standard condition) Keq’(by definition):
Go’ = -RT lnKeq’ = -2.303RT lgKeq’
Keq’ = 10 - G
o’ /1.36
?
Combination and
coupling
of endergonic and
exergonic
reactions; one is
driven by
the other.
Adenine
g
b
a
Ribose
7. The rate of a chemical reaction is determined
by its activation energy (G‡ ), which has no
relation with G
7.1 Each reaction has an activation energy barrier
to get the reactants to the transition state (high
energy state). The energy required to overcome
this energy barrier is called activation energy
(G‡ ).
7.2 The relationship between the rate constant (k),
and the activation energy(G‡ ) is inverse and
exponential:
k=
‡ /RT
G
e
7.3 A small decrease in G‡ results in a large
increase in reaction rate. Enzymes catalyze
biochemical reactions by lowering the activation
energy (thus increasing the rate). Enzymes
exercise no effect on G, therefore do not affect
reaction equilibrium (an enzyme can never
make a reaction of positive G to occur!)
Equillibrium and non-equillibrium
Non-equillibrium but steady-state for living systems
Energy cycle in living systems
8. Reversible interactions within and between
biomolecules are mediated by four
noncovalent weak interactions: hydrogen
bonds, ionic bonds, van der Waals interactions,
hydrophobic interactions
8.1 Hydrogen bonds are formed by two
electronegative atoms (usually N or O, not C in
biomolecules) sharing one hydrogen atom
8.1 The atom to which the H is more
tightly linked is called hydrogen donor, the
other hydrogen acceptor
8.1.2 The bond energy of a H bond is about 8-21
kcal/mol (that of a N-H or O-H covalent bond is about
100 kcal/mol), which is small enough to be broken by
thermal motion of molecules at room temperature
(25C).
8.1.3 The distance between the two hydrogen
bonded electronegative atoms is about 2.9 Å. (short
range)
8.1.4 An important feature of hydrogen bonds is
that they are highly directional, strongest when the
three atoms are colinear.
8.1.5 Hydrogen bonds are abundant and
important in biomolecules. (often responsible for
specificity)
8.2 Ionic bonds occur between two oppositely
charged groups (electrostatic interactions):
F = Q1Q2/ r2
where  is the dielectric constant.
8.2.1 The bonds are much stronger in less
polar environment (in protein,  is smaller, 2.0-4.0;
in water 80.0).
8.2.2 The optimal distance for an ionic bond is
about 2.8 Å.
8.2.3 The bond energy for attraction is about
42 kcal/mol, for repulsion -21 kcal/mol.
8.3 van der Waals interaction is a weak nonspecific
attraction existing between any two atoms at certain
distance (short range).
8.3.1 It is resulted from the interaction of the
induced transient opposite electric dipoles of two
nearby atoms. (Quantum mechanics origin)
8.3.2 It is strongest when the two atoms are 3-4
Å apart.
8.3.3 Any two atoms can not be brought
together closer than the sum of their van der Waals
radii (repulsion due to electron clouds) ( which are
bigger than the covalent radii of the corresponding
atoms).
8.3.4 The bond energy for each van der
Waals bond is about 4 kcal/mol (even weaker
than H bonds).
8.3.5 The contribution of this bond is
significant only when large numbers of them exist
simultaneously.
8.3.6 van der Waals generates very strong
repulsive forces between atoms when they are too
close to each other (hard sphere model). This is
the origin of steric hindrance.
8.4 Hydrophobic interaction is a result of nonpolar
molecules (groups) clustering together in water.
8.4.1 This interaction is not a result of intrinsic
attraction of nonpolar groups.
8.4.2 It is the result of the increased entropy of
freed water molecules when nonpolar molecules are
clustered together. This is because the solvation of the
nonpolar groups need to bind water molecules and the
bound water molecules are more ordered/structured
and have less entropy (randomness). The strength
depends on surface areas.
8.4.3 This interaction is a major driving force in
the folding of macromolecules and formation of
membranes.
8.5 Weak interactions are crucial for the structure
and function of biomolecules
8.5.1 The weak interactions, although
individually weak, accumulate to a significant level
when occur in large numbers. This is the so called
accumulative or collective effect.
8.5.2 Weak interactions maximize when a
macromolecule forms its three-dimensional structure
and interacts with other complementary (cognate)
molecules (producing stability).
8.5.3 The transient nature of weak interactions
confers flexibility for macromolecules and their
interactions.
A brief review of a few other principles
9. All aspects of cell structure and function
are adapted to the physical and chemical
properties of water
9.1 The water molecule is dipolar and highly
cohesive.
9.1.1 The electron outer shell of the
oxygen atom is roughly tetrahedral, with two
hydrogen atoms at two corners, and two pairs
of electrons (lone pairs) at the other two
corners. (sp3)
9.1.2 The oxygen atom is partially negative
and the two hydrogen atoms are partially positive.
9.1.3 Each water molecule can serve as both
H donor and acceptor, forming hydrogen bonds
with as many as four neighboring water molecules,
providing strong cohesive forces between them.
9.1.4 Water has a higher melting point,
boiling point, and heat of vaporization than most
other common solvents due to their high
cohesiveness (hydrogen bonding capacity). (table).
Ordered hydrogen bonding network
9.2 Water is an excellent solvent for polar molecules.
9.2.1 It competes for hydrogen bonds or forms
oriented/structured shells around ions.
9.2.2 Biomolecules are mostly water soluble
(hydrophilic, “water-loving”) molecules.
9.2.3 Water-free microenvironments are also
formed in biological systems to maximize polar
interactions (dielectric constant of water is 80, acting
as a electric screen/shield, while in the protein interior
2-4).
9.3 Hydrophobic (“water hating”, nonpolar)
groups tend to be squeezed/driven together by
water, forming specific biological structures
(interior of globular proteins, biomembranes).
9.4 Water molecules have a slight tendency to undergo
reversible ionization to yield H+ and OH-.
9.4.1 The product of [H+][OH-] in aqueous
solutions at 25C is always 1X10-14 M2 (the measured
Keq for pure water is 1.8X10-16M at 25C and the
concentration of water is 55.5 M) (calculation).
9.4.2 pH is defined as the negative logarithm of
the molar concentration [H+] (pH scale is logarithm, not
arithmetic).
9.4.3 pH scale designates the actual concentration
of [H+] (thus of [OH-], remember their product is a
constant), in any solution in the range between 1.0 M
H+ and 1.0 M OH- (pH value from 1.0 to 14.0).
9.4.4 pH can be accurately measured
using glass electrodes (which is selectively
sensitive to [H+]). The pH meter.
9.4.5 The structure and function of
biomolecules are widely affected by the pH of
the solutions (pH is frequently monitored and
controlled in biochemical reactions).
10. Weak acids and weak bases are common in
biomolecules and their function as pH buffer.
10.1 Weak acids (proton donors) and weak bases
(proton acceptors) do not ionize completely when
dissolved in water.
10.1.1 strong acids (e.g., hydrochloric acids,
sulfuric and nitric acids) and strong bases (e.g.,
NaOH and KOH) ionizes completely when
dissolved in water.
10.1.2 A proton donor and its corresponding
proton acceptor make up a conjugate acid-base
pair
10.1.3 The characteristic tendency of each
weak acid for losing its proton is reflected by its
dissociation constant Ka (stronger acids have
larger Ka value).
10.1.4 For convenience, Ka is converted to
pKa (the negative logarithm).
10.2 The titration curves of weak acids can be fitted by
the Henderson-Hasselbach equation
10.2.1 The amount of weak acids in a solution can
be determined by titrating with a strong base of known
concentration (the solution volume changes little during
titration).
10.2.2 Titration curves are made from plotting the
pH of the solution against the amount of strong base
added, or the fraction (the OH- equivalents) of the total
amount (1.0 OH- equivalents) of strong base required to
neutralize the weak acid (until [proton donor]~0, [proton
acceptor]=initial weak acid concentration, that is, all
weak acid molecules are ionized, no more buffering
effect, pH increases rapidly).
10.2.3 Titration curves of weak acids have nearly
identical shapes (reflecting the same law behind the
phenomenon)
10.2.4 The Henderson-Hasselbalch equation fits
the titration curves of all weak acids!
pH = pKa + log[proton acceptor]/[proton donor]
10.2.5 The pH value at which the conjugate acidbase pair is at equimolar concentration equals to the
pKa value of the weak acid. The plateau of the curve
gives the pKa. The 0.5 OH- equivalents gives the pKa.
10.3 The titration curves of weak acids indicate that a
conjugate acid-base pair can act as a buffer (resisting to
pH changes of the system).
10.3.1 There is a relatively flat zone on all titration
curves of weak acids extending about 0.5 pH units on
either side of the pKa values.
10.3.2 pH changes little in the flat zone when H+
or OH- are added to the system (comparing when added
to pure water). (One should observe this in experiment).
10.3.3 Buffering action is a result of the ionization
of water and weak acid reaching equilibrium
simultaneously (as governed by Kw and Ka).
10.3.4 Many biochemical structures and
processes are affected by pH due to the
involvement of groups behaving as weak acids
and weak bases.
10.3.5 pH values in biological systems are
usually strictly kept constant (near pH 7.0) by
buffering pairs of conjugate acid-bases (e.g., the
phosphate buffer H2PO4-/HPO42- in cytoplasm
and the bicarbonate buffer H2CO3/HCO3- in
blood, better with CO2 gas to dissolved H2CO3
conversion).