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 aremore stable: covalent bonds bysharing unpaired electrons betweentwo atoms.

Versatility of carbon in formingcovalent 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 bondlength = sum ofcovalent radii

Fisher Ball-and-stick Space filling

Convention used in organic chemistry for configuration

Light absorbing pigment in rhodopsinan integral membrane protein

Convention from organic chemistry

(R,R) isomer (S,S) isomer

Highenergy

Lowenergy

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 plane-polarized 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

Chewing gums

绿薄荷

香菜

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 oxidized

Mostreduced

Glucose,source ofelectronsfor metabolism

An oxidation-reduction reaction

Oxidation, losing electrons

Reduction, gaining electrons

Carbonium ion

All have “protruding”lone electron pairs

An isomerization reaction

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

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 attackon 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 - Go’ /1.36

?

Combination and couplingof endergonic and exergonic reactions; one is driven bythe other.

Adenine

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 = e- G‡ /RT

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 (25C).

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 25C is always 1X10-14 M2 (the measured Keq for pure water is 1.8X10-16M at 25C 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 acid-base 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).