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1:IV: Biochemical bonds/weak interactions - Biology

1:IV: Biochemical bonds/weak interactions - Biology


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A. Hydrogen bonds (H-bonds)

The "H-bond donor" - a hydrogen bound to an electronegative atom (usually "O" or "N") in an organic compound forms this bond with an "H-bond acceptor" - an electronegative atom in the same or another compound (also usually and "O" or an "N").
H-bond donors: the "H" in -OH, -NH3+
H-bond acceptors: an "O" in -OH, -C=O, -CO2- or the "N" in -NH2 or the equivalent.

B. Salt bridges

The weak interaction between fully charged ions (e. g. Na+, K+, Cl-, -CO2-, -NH3+) when they are dissolved in an aqueous solution. The water shields the ions making their normally strong interactions much weaker.

C. Van der Waals interactions

(D. Hydrophobic interactions)

Some take home information

Bond/interactionenergylength
Van der Waals0.08 kJ/mole for single atoms (~7 kJ for benzene rings)0.1 - 0.17 nm
H-bond12 - 30 kJ/mol0.2 - 0.27 nm
Ionic Bond~500 kJ/mol
Salt Bridge~30 kJ/mol
Covalent Bond350 - 450 kJ/mol0.1 - 0.15 nm
Hydrophobic interaction0.01 kJ/mol per square angstrom (13 kJ/mol for benzene rings)

Two quarters of organic chemistry are one of the prerequisites for BIS102. What organic chemistry should I remember? Largely the properties of a few kinds of molecules that are important in biochemistry. At a minimum, you should recall the structure and properties of the following compounds/functional groups:

  1. Carboxylic acids
  2. Alcohols
  3. Amines
  4. Aromatic rings
  5. Alkanes
  6. C-C single and double bonds (e. geometry - tetrahedral, free rotation vs. planar, restricted rotation)
  7. Aldehyde/ketone (i. e. carbonyl oxygen)

Also, you should be familiar with the basics of the processes of nucleophilic attack, acid catalysis and base catalysis.


Biochemistry Exam 1

Any molecule that can release a proton has an acid and conjugate base form.
Can describe ionization using equilibrium constant - changed name to acid dissociation constant.
Acid dissocitation constant is Ka, molecules that are strong acids release proton really easily, they have a high Ka. Something that has low Ka is a base that tends to stay in acid form. Reaction tends to go to the right. If reaction goes in other direction then it's a base.
pKa= -logKa (its opposite day)
Low pKa means strong acid, high pKa means strong base
Acetic acid(vinegar) is a weak acid. -log of Ka to get pKA. pKAbelow 7 so means acid. When pH = pKA, 50% of molecule will be in HA form and 50% will be in A- form.
CH3COOHóCH3COO- + H+
HA A-

Important to recognize pKAis property of molecule (acetic acid). pH changes, if lower pH,
Lowering pH means increasing H+ concentratin, which means reaction will go to the left
Much of metabolism governendby principle that changing concentration on one side - Le Chattiers Principle
If change concentration on either side, if you increase the concentration it flows the other way.
Reactions are responding to changing concentrations
Le Chattiersprinciple affects which way reaction goes, at any other pH one form is dominant

HA «H++ A-
•Acid dissociation constantdescribes equilibrium.

8.5 = 7 + log([A-]/[HA])
8.5 - 7 = log([A-]/[HA])
1.5 = log([A-]/[HA])
10(1.5)= [A-]/[HA] = 31.6 (31.6x more A-than HA)


Abstract

Experimental charge density distributions in a series of ionic complexes of 1,8-bis(dimethylamino)naphthalene (DMAN) with four different acids: 1,2,4,5-benzenetetracarboxylic acid (pyromellitic acid), 4,5-dichlorophthalic acid, dicyanoimidazole, and o-benzoic sulfimide dihydrate (saccharin) have been analyzed. Variation of charge density properties and derived local energy densities are investigated, over all inter- and intramolecular interactions present in altogether five complexes of DMAN. All the interactions studied <[O···H···O] - , C−H···O, [N−H···N] + , O−H···O, C−H···N, Cπ···Nπ, Cπ···Cπ, C−H···Cl, N−H + >follow exponential dependences of the electron density, local kinetic and potential energies at the bond critical points on the length of the interaction line. The local potential energy density at the bond critical points has a near-linear relationship to the electron density. There is also a Morse-like dependence of the laplacian of rho on the length of interaction line, which allows a differentiation of ionic and covalent bond characters. The strength of the interactions studied varies systematically with the relative penetration of the critical points into the van der Waals spheres of the donor and acceptor atoms, as well as on the interpenetration of the van der Waals spheres themselves. The strong, charge supported hydrogen bond in the DMANH + cation in each complex has a multicenter character involving a <[Me2N−H····NMe2] + ····X δ- > assembly, where X is the nearest electronegative atom in the crystal lattice.

In papers with more than one author, the asterisk indicates the name of the author to whom inquiries about the paper should be addressed.


Hydrogen Bonding

A hydrogen bond is a strong intermolecular force created by the relative positivity of hydrogen atoms.

Learning Objectives

Describe the properties of hydrogen bonding.

Key Takeaways

Key Points

  • Hydrogen bonds are strong intermolecular forces created when a hydrogen atom bonded to an electronegative atom approaches a nearby electronegative atom.
  • Greater electronegativity of the hydrogen bond acceptor will lead to an increase in hydrogen-bond strength.
  • The hydrogen bond is one of the strongest intermolecular attractions, but weaker than a covalent or an ionic bond.
  • Hydrogen bonds are responsible for holding together DNA, proteins, and other macromolecules.

Key Terms

  • electronegativity: The tendency of an atom or molecule to draw electrons towards itself, form dipoles, and thus form bonds.
  • hydrogen bond: The attraction between a partially positively charged hydrogen atom attached to a highly electronegative atom (such as nitrogen, oxygen, or fluorine) and another nearby electronegative atom.
  • intermolecular: A type of interaction between two different molecules.

Forming a Hydrogen Bond

A hydrogen bond is the electromagnetic attraction created between a partially positively charged hydrogen atom attached to a highly electronegative atom and another nearby electronegative atom. A hydrogen bond is a type of dipole-dipole interaction it is not a true chemical bond. These attractions can occur between molecules (intermolecularly) or within different parts of a single molecule (intramolecularly).

Hydrogen bonding in water: This is a space-filling ball diagram of the interactions between separate water molecules.

Hydrogen Bond Donor

A hydrogen atom attached to a relatively electronegative atom is a hydrogen bond donor. This electronegative atom is usually fluorine, oxygen, or nitrogen. The electronegative atom attracts the electron cloud from around the hydrogen nucleus and, by decentralizing the cloud, leaves the hydrogen atom with a positive partial charge. Because of the small size of hydrogen relative to other atoms and molecules, the resulting charge, though only partial, is stronger. In the molecule ethanol, there is one hydrogen atom bonded to an oxygen atom, which is very electronegative. This hydrogen atom is a hydrogen bond donor.

Hydrogen Bond Acceptor

A hydrogen bond results when this strong partial positive charge attracts a lone pair of electrons on another atom, which becomes the hydrogen bond acceptor. An electronegative atom such as fluorine, oxygen, or nitrogen is a hydrogen bond acceptor, regardless of whether it is bonded to a hydrogen atom or not. Greater electronegativity of the hydrogen bond acceptor will create a stronger hydrogen bond. The diethyl ether molecule contains an oxygen atom that is not bonded to a hydrogen atom, making it a hydrogen bond acceptor.

Hydrogen bond donor and hydrogen bond acceptor: Ethanol contains a hydrogen atom that is a hydrogen bond donor because it is bonded to an electronegative oxygen atom, which is very electronegative, so the hydrogen atom is slightly positive. Diethyl ether contains an oxygen atom that is a hydrogen bond acceptor because it is not bonded to a hydrogen atom and so is slightly negative.

A hydrogen attached to carbon can also participate in hydrogen bonding when the carbon atom is bound to electronegative atoms, as is the case in chloroform (CHCl3). As in a molecule where a hydrogen is attached to nitrogen, oxygen, or fluorine, the electronegative atom attracts the electron cloud from around the hydrogen nucleus and, by decentralizing the cloud, leaves the hydrogen atom with a positive partial charge.

Interactive: Hydrogen Bonding: Explore hydrogen bonds forming between polar molecules, such as water. Hydrogen bonds are shown with dotted lines. Show partial charges and run the model. Where do hydrogen bonds form? Try changing the temperature of the model. How does the pattern of hydrogen bonding explain the lattice that makes up ice crystals?

Applications for Hydrogen Bonds

Hydrogen bonds occur in inorganic molecules, such as water, and organic molecules, such as DNA and proteins. The two complementary strands of DNA are held together by hydrogen bonds between complementary nucleotides (A&T, C&G). Hydrogen bonding in water contributes to its unique properties, including its high boiling point (100 °C) and surface tension.

Water droplets on a leaf: The hydrogen bonds formed between water molecules in water droplets are stronger than the other intermolecular forces between the water molecules and the leaf, contributing to high surface tension and distinct water droplets.

In biology, intramolecular hydrogen bonding is partly responsible for the secondary, tertiary, and quaternary structures of proteins and nucleic acids. The hydrogen bonds help the proteins and nucleic acids form and maintain specific shapes.


Formation of Hydrophobic Interactions

The mixing hydrophobes and water molecules is not spontaneous however, hydrophobic interactions between hydrophobes are spontaneous. When hydropobes come together and interact with each other, enthalpy increases ( ( Delta H ) is positive) because some of hydrogen bonds that form the clathrate cage will be broken. Tearing down a portion of the clathrate cage will cause the entropy to increase ( ( Delta S ) is positive), since forming it decreases the entropy.

According to the Equation ( ef)

Result: ( Delta ) is negative and hence hydrophobic interactions are spontaneous.


Molecular Forces in Anesthesia

Several observations and ideas on possible anesthesia mechanism have appeared, as theories, and have created considerable controversies, indicating that these proposed ideas are more in the nature of interesting, thought-provoking hypotheses, than they are actual scientific theories. All these hypotheses are based on the correlations between anesthetic potency and some physical properties of the relevant agents, such as solubility, polarizability, refractivity. These properties are all intimately interrelated and are consequences of the atomic structure. This chapter discusses the basic molecular interactions in which anesthetic agents can be involved in their biological surroundings. This study is generally restricted to the so-called “inert gaseous anesthetic agents” that exert their biological effects without undergoing any change in their own chemical structures. Attention has been focused on the behavior and distribution of these discrete molecules in their biological surrounding. The various binding forces, through which the so-called “inert gaseous anesthetic agents” associate with macromolecules, are discussed to demonstrate what conditions are prerequisite for binding. If the molecular natures of particular binding sites are determined, it will be feasible to elucidate how such an interaction affects the function of the systems. Recent advances in protein crystallography have made now it possible to study drug-protein interaction on a truly molecular level.

Present address of both authors: Department of Pharmacology, University of California, San Francisco, California.


Strong and Weak Hydrogen Bonds in Protein–Ligand Recognition

The hydrogen bond has justifiably been termed the ‘master key of molecular recognition’. It is an interaction that is weaker than the covalent bond and stronger than the van der Waals interaction. The ubiquity and flexibility of hydrogen bonds make them the most important physical interaction in systems of biomolecules in aqueous solution. Hydrogen bonding plays a significant role in many chemical and biological processes, including ligand binding and enzyme catalysis. In biological processes, both specificity and reversibility are important. Weaker interactions can be made and broken more easily than stronger interactions. In this context, it is of interest to assess the relative significance of strong and weak interactions in the macromolecular recognition processes. Is protein–ligand binding governed by conventional, that is, electrostatic N–H…O and O–H…O hydrogen bonds, or do weaker interactions with a greater dispersive component such as C–H…O hydrogen bonds also play a role? If so, to what extent are they significant? Most proteins, involving as they do, main chains, side chains, and differently bound forms of water, do not really have a static fixed structure, but rather have a dynamic, breathing nature. This tendency may to some extent be lessened by the ligands which are small molecules, but in the end, it is reasonable to expect that the strong and weak hydrogen bonds inside the protein and also at the protein–ligand interface will also have dynamic character arguably, the weaker the hydrogen bond, the greater its dynamic character. These are often central to the much debated mechanisms of binding such as conformational selection and induced fit. All protein–ligand interactions must compete with interactions with water both the protein and the ligand are solvated before complexation and lose their solvation shell on complex formation. Conversely, the entropic cost of trapping highly mobile water molecules in the binding site is large. However, in favorable cases, these losses are suitably compensated by the enthalpic gain resulting from water-mediated hydrogen bonds. In effect, the enthalpy–entropy balance is a fine one, and for a water molecule to be able to contribute to binding affinity, it has to be in a binding site that provides the maximum number of hydrogen-bond partners at the optimum distance and orientation. In summary, hydrogen bonds are crucial to the recognition of ligands by proteins. Integration of knowledge gained from more high-quality protein–ligand structures into theoretical and computational molecular models will be an exciting challenge in the coming years.

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H-Bonds and Water

H-bonding occurs in water. In the liquid state they are rapidly being formed and broken as the mobile particles move over each other. Note in figure (PageIndex<2>) that there are two type of O-H bonds, the intramolecular O-H bond within a molecule (bond length = 1.01Å) and the intermolecular bond between atoms (bond length = 1.75Å). The closeness of the bond length indicates that the intramolecular bond is very strong, and of comparable magnitude to the intramolecular one. On a macroscopic scale this is obvious to anyone who falls on ice (frozen water) and quickly realizes how hard it is, and how strong the bonds are that hold the water molecules to each other.

Figure (PageIndex<2>): Hydrogen bonding interactions within water. Note the similarity in length between the intermolecular O---H bond, and the intramolecular O-H bond.

Water also has two lone pairs and two H atoms attached to the highly electronegative oxygen. This means each water molecule can participate in up to 4 bonds (two where it is the h-bond acceptor, and two where it is the h-bond donor). One interesting consequence of this is that water forms a 3D crystalline structure that is sort of based on a distorted tetrahedron. That is, the oxygen is sp 3 hybridized with a tetrahedral electronic geometry, having two bonding orbitals and two lone pairs. All of these are involved with hydrogen bonds. The lone pairs are functioning as H-bond acceptors, and the hydrogen on the bonding orbitals are functioning as h-bond donors. So each oxygen is attached to 4 hydrogens, two are 1.01Å covalent bonds and two are 1.75 Å hydrogen bonds, and this results in a structure like figure 11.5.3, which has lots of void space, and the consequence that ice is less dense than liquid water and floats.

Figure (PageIndex<3>): Hydrogen bonding and the crystal structure of ice (left), which floats on liquid water (rights, wikicommons image credit)

The fact that ice floats has great ramifications for life. The ice, with it's void space, acts as an insulator. If the ice sunk to the bottom, lakes would completely freeze and aquatic life like fish would not be able to survive the winters.

H-bonding and Boiling Points

Boiling points are an indicator of intermolecular forces, and we will look at the phenomena of boiling in more detail in a later section of this chapter. From the kinetic molecular theory of gases we learned that the kinetic energy of a molecule is proportional to its absolute temperature, where KE Translation = 1/2mv 2 , (m is the mass, and v is the velocity). We also learned that there is a velocity profile with different molecules moving at different speeds, but that heavier molecules tend to move slower than lighter ones (remember macroscopic observables like liquids and boiling are the result of the interactions of a huge number of molecules which possess a distribution of energies). So all things equal, we would anticipate that it is easier to boil a lighter molecule than a heavier one, and we would predict the heavier one to have a higher boiling point.

Figure (PageIndex<4>). shows the trend for various hydrides of different families of the periodic table. Our predictions based on molar mass work for group IV (CH4 through SnH4), but does not work for groups VII, VI and V, where the hydrides of the second period (N,O,F) have higher boiling points than would be predicted from their mass trends.

Figure (PageIndex<4>): Periodic trends in boiling point for various hydrides.

There are a multitude of questions that can come out of figure (PageIndex<4>), and you should familiarize yourself with this figure. You should also revisit this figure after we have covered the section on boiling. Right now let is suffice that to boil a substance you need to overcome the intermolecular forces that hold the liquid together (the cohesive forces), and that as you increase the temperature you increase the kinetic energy of the molecules (so they can escape the liquid and become a vapor).

First, lets look at the group 4 hydrides (CH4, SiH4, GeH4& SnH4),. As we go down the table from carbon to silicon to germanium to Tin we see an increase in the boiling points. These are symmetric nonpolar molecules and there are two reasons why going down this family shows an increase in the boiling point.

  1. Going down group 4 the mass increases, requiring a higher temperature for molecules to gain enough kinetic energy to vaporize.
  2. Going down the table the valence electrons occupy more diffuse orbitals, resulting in a higher polarizability, which in turn results in larger London dispersion forces, which requires more

So both of these trends would indicate an increase in the boiling point as we go down the family.

Second, ammonia, hydrofluoric acid and water show a deviation in this trend. That is, they are lighter, but there is an increase in their boiling points. This is due to hydrogen bonding. That is, although these molecules are lighter, they have very strong intermolecular forces, which must be overcome for them to boil.

Third, water has a higher boiling point than HF, yet fluorine is more electronegative than oxygen, it is also smaller, and so you would expect the HF hydrogen bond to be stronger than the OH hydrogen bond. In fact, this is all try, and what this argument does not take into account is the number of hydrogen bonds. That is, the cohesive forces that hold the liquid together are not just the strength of the bonds, but also the number of bonds. Each HF molecule has one H, and 3 lone pairs on the fluorine. So in a macroscopic system like a mole of HF (remember that to be a liquid, you need a lot of molecules), each HF would on average be involved with 2 bonds, one involving the hydrogen (hydrogen donor), and one involving a fluorine lone pair (hydrogen acceptor), and simple speaking, there are not enough hydrogens to use up all the lone pairs, (two of the fluorine's lone pairs are not involved in H bonds.) In the case of water, the number of lone pairs equals the number of hydrogens, and so each water molecule can on the average, be involved with 4 hydrogen bonds. So even though the individual hydrogen bonds in water may be weaker than in HF, there are more of them, making the boiling point of water higher than HF.

Exercise 11.5.2: Look at Real Molecules (Exam Material)

Before reading on, look at the following three molecules and their boiling points, and try and answer the following 2 questions.

1. Why is pentane's boiling point so much lower than the other two?

Pentane is non polar and the other two have hydrogen bonds through the OH group.

2. Why does butan-1-ol have a higher boiling point than 2-methylpropan-1-ol?

Both of these molecules are isomers with the same chemical constituents (C4H10O) and both have an OH group than can be involved with hydrogen bonds. But butan-1-ol is more diffuse, thus more polarizable and has stronger van der Waals (London dispersion) interactions


Molecular Biology of the Gene(6th Edition) epub 下载 mobi 下载 pdf 下载 txt 下载

Molecular Biology of the Gene(6th Edition) epub 下载 mobi 下载 pdf 下载 txt 下载

Molecular Biology of the Gene(6th Edition) pdf epub mobi txt 下载

Though completely up-to-date with the latest research advances, the Sixth Edition of James D. Watson’s classic book, Molecular Biology of the Gene retains the distinctive character of earlier editions that has made it the most widely used book in molecular biology. Twenty-two concise chapters, co-authored by six highly respected biologists, provide current, authoritative coverage of an exciting, fast-changing discipline. Mendelian View of the World, Nucleic Acids Convey Genetic Information,The Importance of Weak Chemical Interactions, The Importance of High Energy Bonds, Weak and Strong Bonds Determine Macromolecular Interactions, The Structures of DNA and RNA, Genome Structure, Chromatin and the Nucleosome, The Replication of DNA, The Mutability and Repair of DNA,Homologous Recombination at the Molecular Level, Site-Specific Recombination and Transposition of DNA, Mechanisms of Transcription 13 RNA Splicing, Translation, The Genetic Code, Transcriptional Regulation in Prokaryotes, Transcriptional Regulation in Eukaryotes, Regulatory RNAs, Gene Regulation in Development and Evolution, Genomics and Systems Biology, Techniques of Molecular Biology, Model Organisms. Intended for those interested in learning more about the basics of Molecular Biology.

James D. Watson was Director of Cold Spring Harbor Laboratory from 1968 to 1993, President from 1994 to 2003, and is now its Chancellor. He spent his undergraduate years at the University of Chicago and received his Ph.D. in 1950 from Indiana University. Between 1950 and 1953, he did postdoctoral research in Copenhagen and Cambridge, England. While at Cambridge, he began the collaboration that resulted in the elucidation of the double-helical structure of DNA in 1953. (For this discovery, Watson, Francis Crick, and Maurice Wilkins were awarded the Nobel Prize in 1962.) Later in 1953, he went to the California Institute of Technology. He moved to Harvard in 1955, where he taught and did research on RNA synthesis and protein synthesis until 1976. He was the first Director of the National Center for Genome Research of the National Institutes of Health from 1989 to 1992. Dr. Watson was sole author of the first, second, and third editions of Molecular Biology of the Gene, and a co-author of the fourth and fifth editions. These were published in 1965, 1970, 1976, 1987, and 2003, respectively. He is also a co-author of two other textbooks: Molecular Biology of the Cell and Recombinant DNA.

Tania A. Baker is the Whitehead Professor of Biology at the Massachusetts Institute of Technology and an Investigator of the Howard Hughes Medical Institute. She received a B.S. in biochemistry from the University of Wisconsin, Madison, and a Ph.D. in biochemistry from Stanford University in 1988. Her graduate research was carried out in the laboratory of Professor Arthur Kornberg and focused on mechanisms of initiation of DNA replication. She did postdoctoral research in the laboratory of Dr. Kiyoshi Mizuuchi at the National Institutes of Health, studying the mechanism and regulation of DNA transposition. Her current research explores mechanisms and regulation of genetic recombination, enzyme-catalyzed protein unfolding, and ATP-dependent protein degradation. Professor Baker received the 2001 Eli Lilly Research Award from the American Society of Microbiology and the 2000 MIT School of Science Teaching Prize for Undergraduate Education and was elected as a fellow of the American Academy of Arts and Sciences in 2004. She is co-author (with Arthur Kornberg) of the book DNA Replication, Second Edition.

Stephen P. Bell is a Professor of Biology at the Massachusetts Institute of Technology and an Investigator of the Howard Hughes Medical Institute. He received B.A. degrees from the Department of Biochemistry, Molecular Biology, and Cell Biology and the Integrated Sciences Program at Northwestern University and a Ph.D. in biochemistry at the University of California, Berkeley in 1991. His graduate research was carried out in the laboratory of Dr. Robert Tjian and focused on eukaryotic transcription. He did postdoctoral research in the laboratory of Dr. Bruce Stillman at Cold Spring Harbor Laboratory, working on the initiation of eukaryotic DNA replication. His current research focuses on the mechanisms controlling the duplication of eukaryotic chromosomes. Professor Bell received the 2001 ASBMB–Schering Plough Scientific Achievement Award, the 1998 Everett Moore Baker Memorial Award for Excellence in Undergraduate Teaching at MIT and the 2006 MIT School of Science Teaching Award.

Alexander A. F. Gann is Editorial Director of Cold Spring Harbor Laboratory Press, and a faculty member of the Watson School of Biological Sciences at Cold Spring Harbor Laboratory. He received his B.Sc in microbiology from University College London and a Ph.D. in molecular biology from The University of Edinburgh in 1989. His graduate research was carried out in the laboratory of Noreen Murray and focused on DNA recognition by restriction enzymes. He did postdoctoral research in the laboratory of Mark Ptashne at Harvard, working on transcriptional regulation, and that of Jeremy Brockes at the Ludwig Institute of Cancer Research at University College London, where he worked on newt limb regeneration. He was a Lecturer at Lancaster University, U.K., from 1996 to 1999, before moving to Cold Spring Harbor Laboratory. He is co-author (with Mark Ptashne) of the book Genes & Signals (2002).

Michael Levine is a Professor of Molecular and Cell Biology at the University of California, Berkeley, and is also Co-Director of the Center for Integrative Genomics. He received his B.A. from the Department of Genetics at University of California, Berkeley, and his Ph.D. with Alan Garen in the Department of Molecular Biophysics and Biochemistry from Yale University in 1981. As a postdoctoral fellow with Walter Gehring and Gerry Rubin from 1982-1984, he studied the molecular genetics of Drosophila development. Professor Levine's research group currently studies the gene networks responsible for the gastrulation of the Drosophila and Ciona (sea squirt) embryos. He holds the F. Williams Chair in Genetics and Development at University of California, Berkeley. He was awarded the Monsanto Prize in Molecular Biology from the National Academy of Sciences in 1996, and was elected to the American Academy of Arts and Sciences in 1996 and the National Academy of Sciences in 1998.

Richard M. Losick is the Maria Moors Cabot Professor of Biology, a Harvard College Professor, and a Howard Hughes Medical Institute Professor in the Faculty of Arts & Sciences at Harvard University. He received his A.B. in chemistry at Princeton University and his Ph.D. in biochemistry at the Massachusetts Institute of Technology. Upon completion of his graduate work, Professor Losick was named a Junior Fellow of the Harvard Society of Fellows when he began his studies on RNA polymerase and the regulation of gene transcription in bacteria. Professor Losick is a past Chairman of the Departments of Cellular and Developmental Biology and Molecular and Cellular Biology at Harvard University. He received the Camille and Henry Dreyfuss Teacher-Scholar Award, is a member of the National Academy of Sciences, a Fellow of the American Academy of Arts and Sciences, a Fellow of the American Association for the Advancement of Science, a Fellow of the American Academy of Microbiology, a member of the American Philosophical Society, and a former Visiting Scholar of the Phi Beta Kappa Society. Professor Losick is the 2007 winner of the Selman A. Waksman Award of the National Academy of Sciences.

PART 1 Chemistry and Genetics
1 Mendelian View of the World
2 Nucleic Acids Convey Genetic Information
3 The Importance of Weak Chemical Interactions
4 The Importance of High Energy Bonds
5 Weak and Strong Bonds Determine Macromolecular Interactions

PART 2 Maintenance of the Genome
6 The Structures of DNA and RNA
7 Genome Structure, Chromatin and the Nucleosome
8 DNA replication
9 The Mutability and Repair of DNA
10 Homologous Recombination at the Molecular Level
11 Site-Specific Recombination and Transposition of DNA

PART 3 Expression of the Genome
12 Mechanisms of Transcription
13 Splicing
14 Translation

PART 4 Regulation
15 The Genetic Code
16 Transcriptional Regulation in Prokaryotes
17 Transcriptional Regulation in Eukaryotes
18 Regulatory RNAs
19 Gene Regulation in Development and Evolution
20 Genomics and Systems Biology

PART 5 Methods
21 Techniques of Molecular Biology
22 Model Organisms
· · · · · · (收起)


1:IV: Biochemical bonds/weak interactions - Biology

The nature of the bonding and a definite preference for an eclipsed geometry in several substituted but-2-ynes, including certain novel derivatives are uncovered and examined. In particular, we consider the molecular species R<sub>3</sub>CCCCR<sub>3</sub> (where R= H, F, Cl, Br, I, and CN), their R<sub>3</sub>CBNCR<sub>3</sub> analogues, and a few novel exo-bridge systems with intramolecular hydrogen bonds running parallel to the CCCC chain. In some cases, the potential energy surfaces are remarkably flatso flat, in fact, that free rotation is predicted for those molecules at very low temperatures. A systematic investigation of the bonding in the halogenated butynes demonstrates that the eclipsed conformation actually becomes more stable relative to the staggered form as R becomes larger and less electron-withdrawing. The rotational barriers (the differences in energy between the eclipsed and staggered geometries) are magnified significantly, however, in a special case where selected R groups at the ends of the R<sub>3</sub>CCCCR′<sub>3</sub> molecule form hydrogen bonds parallel to the CCCC core. In those systems, the hydrogen bonds serve as a weak locking mechanism that favors the eclipsed conformation. A comparison of HF and uncorrected DFT methods versus the MP2­(full), CCSD­(T), and other dispersion-corrected methods confirms that correlation accounts to a significant extent for barriers in substituted butyne compounds. In the hydrogen-bonded systems, the barriers are comparable to and larger in some cases than the barriers observed for the more extensively studied ethane molecule


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