Life Sciences 1a - Harvard Universitysites.fas.harvard.edu/~lsci1a/9-21notes.pdf · Unlike the macromolecules of life, the small molecules used by living systems are extremely diverse

Professor David Liu and Brian Tse, Life Sciences 1a page 1

Life Sciences 1aAn Integrated Introduction

to the Life Sciences

Lecture Slides Set 2Fall 2006-2007Prof. David R. Liu


“In my hunt for the secret of life, I started research inhistology. Unsatisfied by the information that cellularmorphology could give me about life, I turned tophysiology. Finding physiology too complex I took uppharmacology. Still finding the situation too complicated Iturned to bacteriology. But bacteria were even toocomplex, so I descended to the molecular level, studyingchemistry and physical chemistry. After twenty years’work, I was led to conclude that to understand life wehave to descend to the electronic level, and to the worldof wave mechanics. But electrons are just electrons, andhave no life at all. Evidently, on the way I lost life; it hadrun out between my fingers.”

– Albert Szent-Gyorgyi, Personal Reminisces (1937 Nobel Laureate in Physiology)


Lectures 2 & 3: The molecules of life and their chemical foundations1. The molecules of life comprise macromolecules and small molecules2. Understanding the molecules of life: chemical structures and bonding

a. Connectivity versus conformationb. The nature of atomsc. Covalent bonding and formal chargesd. Ionic bondinge. Electronegativity and hydrogen bondsf. The bonding continuum and bond polarity

3. Organic molecules and how to draw thema. The molecules of life are organic moleculesb. The geometries of organic moleculesc. Drawing organic moleculesd. Understanding “arrow-pushing” notation

4. Stereochemistry and the molecules of lifea. Stereoisomers and enantiomersb. Drawing stereoisomersc. The tragedy of thalidomided. Geometric isomerse. The role of geometric isomers in vision

5. The molecular components of HIV

Required: IIB lecture notes;McMurry Ch. 1 (review), Ch.3, Ch. 5, p. 235-237, p. 354-370; Alberts pp. 39-59, 66,78-79

Lecture Readings

1. The molecules of life comprise macromolecules and small molecules

In the last lecture, we learned how observations of ill patients eventually lead to thediscovery that HIV is the causal agent of AIDS. Scientists have long realized that fighting adisease in the most effective way possible requires an understanding of both the causativeagent as well as the life processes affected by the disease. As Albert Szent-Gyorgyi’s quotereflects, this understanding can come at many different levels, ranging from populationsurveys of affected humans to a sub-atomic analysis of an atom involved in the disease. Formany problems in the life sciences, a useful compromise between an understanding that istoo simplistic and one that contains an unhelpful level of detail is an examination at themolecular level. Over the next several lectures we will introduce the chemical principlesunderlying the molecules of life.

Life is complex— so complex that scientists and philosophers continue to debate whatconstitutes a living system. Despite this complexity, all living systems share certainfundamental features such as the ability to propagate, to respond to changes in theenvironment, and to protect components essential to their proper function. Each of thesefundamental features of life is made possible by a collection of chemical structures that wecall “the molecules of life”. Because living systems as different as bacteria and humans allrely on this same set of molecules, they have been the focus of intense scientific research.


H2N

O

HO

OH

OOPHO

OH

O

N

N

N

N

NH2

HO

NH

O

HN

NH

O

O

HO

OH

BiologicalPolymers

nucleotidemonomer

nucleic acid polymer

protein polymer

amino acidmonomer

O

O

N

N

N

NN

O

PO

OO

O

PO

OO

O

N

N

O

H

O

CH3H

H

OP

O

O

O

PO

OO

O

N

N

O

H

O

CH3

monomers polymer

The molecules of life vary greatly in size, composition, and shape. Some of the mostimportant molecules for living systems are biological polymers, such as DNA, RNA, andproteins. Polymers consist of many repeating units, or monomers. Because these polymersare typically made of hundreds or thousands of atoms, they are commonly referred to asmacromolecules. Biological macromolecules play some of the most important roles in livingsystems. They are responsible for storing genetic information to be passed down to futuregenerations as well as accelerating chemical reactions required by the cell. These moleculesand their roles will be discussed in much greater detail throughout this course. Don’t worry ifyou aren’t yet be able to understand the structures of these molecules shown in the followingfigures; their details are not important at this point. These structures are shown simply to giveyou a sense of the diverse nature of the macromolecules of life.


Biological Small Molecules

Glucose(Sugar)

Pyridoxine(Vitamin B6)

N

H

OH

OH

HO

CH3

Cholesterol(Steroid)

NH

HO

NH2

Serotonin(Neurotransmitter)

OHO

HO

OHOH

HO

O

OHH2N

S

Methionine(Amino Acid)

OOPO

PO

PHO

HO OH OH

O O O

N

N

N

N

NH2

HO OH

Adenosine Triphosphate(Nucleotide)

HH3C

HO

H H

H3C

CH3

H3C

H3C

In addition to macromolecules, small molecules also play a central role in living systems.Although there is no discrete size cut-off that distinguishes “small” molecules frommacromolecules, most small molecules relevant to life contain fewer than about 100 atoms.Unlike the macromolecules of life, the small molecules used by living systems are extremelydiverse in their basic chemical structures. Their more diverse structures also imply that smallmolecules are synthesized, both in the laboratory and in the cell, by a much more diversecollection of chemical reactions than those used to make macromolecules. Indeed, as we willlearn later in this course, macromolecules are typically generated by repeating one type ofchemical reaction over and over, while small molecules are synthesized through the use ofthousands of different reactions.


Small Molecules on a Macromolecular Scale

OOPHO

OH

O

N

N

N

N

NH2

HO OH

DNA

Small molecule(Adenosine monophosphate)

Protein(DNA Ligase)

A major recurring theme in this course is that the structure of a molecule determines itschemical and biological properties. Because small molecules are structurally more varied thanmacromolecules, their chemical and biological roles are in some ways more diverse than thoseof macromolecules. Nutrients such as simple sugars and vitamins, signaling molecules such assteroids and neurotransmitters, and the monomers that are linked together to form DNA, RNA,and proteins are all examples of small molecules essential to a wide variety of living processes.The smaller sizes of these molecules compared with macromolecules also enable them toperform roles in living systems that are not possible for macromolecules, such as to quicklyenter cells or travel between different cells. As we will see in a later part of this course, thevast majority of drugs to treat human diseases are also small molecules for some of thesereasons.


Small Connectivity Differences Can DramaticallyAlter the Properties of Molecules

OOPHO

OH

O

N

N

N

N

NH2

HO OH

OOPHO

OH

O

N

N

N

N

NH2

HO

RNA monomer DNA monomer

no change!

water1 day

water1 day

RNA degraded RNA DNA

• Structure = connectivity + conformation• Even modest changes in connectivity can have profound effects on function

2. Understanding the molecules of life: chemical structures and bonding

Connectivity versus conformation

The word “structure” applied to the molecules of life can have several distinct meanings. Inthis course, we will focus on two of them: (i) the connectivity of atoms within a molecule (thatis, the map of chemical bonds within a molecule), and (ii) the relative three-dimensionalorientation of chemically bonded atoms within a molecule (also called a molecule’sconformation). Altering either type of structure within a molecule can cause profound changesin that molecule’s properties. For example, DNA and RNA differ in the map of their chemicalbonds by only a single change in each monomer, yet as a result DNA is much more resistant todegradation in water than RNA (the basis of this difference will be explained in the anupcoming lecture).


One Connectivity, Two Conformations

Change in saltconcentration

(only for certainDNA sequences)

B-form DNA Z-form DNA

• Similarly, changes in conformation can also dramatically affect function

Many of the molecules of life can adopt multiple conformations and can change betweenthese conformations without breaking any connections between atoms. As a dramaticexample, consider this preview of the famous double-helical structure of DNA. The mostcommon conformation of double-stranded DNA is known as B-form DNA. Now compare thisstructure with a second conformation of DNA, called Z-form DNA, that certain double-strandedsequences can adopt in the presence of very high concentrations of salt. These two forms ofDNA are clearly very different even though the connectivity of the atoms within these twoconformations is identical. As you probably already guessed, the biological properties of thesetwo conformations of the same DNA molecule are also very different; B-form DNA, but not Z-form DNA, is recognized and processed by the machinery in the cell.


The Periodic Table of the Elements

H, C, N, O, P, S (aka SPONCH) = 99% of living matter

The nature of atoms

To understand the structure of the molecules that make up living systems requiresunderstanding the nature of atoms and chemical bonds. The molecules of life are made ofatoms, the smallest units of the approximately 100 elements of the periodic table. The vastmajority of biological matter— about 99% in fact— are made of just six kinds of these atoms:carbon, hydrogen, nitrogen, oxygen, sulfur, and phosphorus (conveniently remembered asSPONCH).


Basic Structure of Atoms

usually 7 neutrons

7 protons

NLewis dotstructure

A nitrogen atom(nucleus greatly enlarged for clarity)

Electron clouds(“orbitals”)

7 electrons total5 valence electrons(in the “outer shell”)

Atoms are made of positively charged protons, neutral neutrons, and negatively chargedelectrons. The protons and neutrons form the tiny atomic nucleus at the center of an atom.Despite drawings you may have seen that liken electrons to planets orbiting a star, electronsactually exist as clouds of electron density surrounding the nucleus. These clouds are calledorbitals and have specific three-dimensional shapes. Differences in the number of protons,neutrons, and electrons within each atom give rise to the different properties of each element.

Although the nuclei of atoms as well as the electrons closest to the nuclei rarely changeduring the course of the chemical reactions described in this course, the outermost electronseach atom frequently undergo changes, both in their number and in the shapes of the spacethat they occupy. The outermost electrons of an atom are also called the valence electrons.Students learning chemistry often draw dots around an atom’s symbol to represent its valenceelectrons. Representations of atoms and molecules that use dots to keep track of valenceelectrons are called Lewis dot structures.


The Excerpted Periodic Table

period

group

Number ofvalenceelectrons: 1 2 3 4 5 6 7 0

(8)

Electron shell being filledshell 1

shell 2

shell 3

6C

12.011carbon

Atomic number(number of protons)

Atomic mass(weighted averageof 98.9% carbon-12 +1.1% carbon-13)

Let’s take a closer look at the part of the periodic table that includes the SPONCHelements— the first three rows. The rows of the periodic table are called periods, while thecolumns are called groups. Each of the first two periods represent the first or second electron“shell” of an atom. You can therefore see that there are up to two electrons in the first shelland up to eight electrons in the second shell. The third electron shell can hold up to 18electrons, the first eight of which are represented by the elements in the third period. Thegroups of the first three periods are designated (from left to right) 1A, 2A, 3A, 4A, 5A, 6A, 7A,and 8A. These group designations tell you how many valence electrons are present in eachelement. For example, carbon is in group 4A, and therefore has four valence electrons. Thegroup 8A elements include elements with filled sets of electron shells, however, and thereforeare considered to have zero (rather than eight) valence electrons.

The atomic number (the whole numbers on the periodic table started with 1 for hydrogen, 2for helium, etc.) is the number of protons in an atom of a given element. The atomic mass(the number with several decimal places, such as 12.011 for carbon) is the weighted averageof the total mass of the protons, neutrons, and electrons of one atom considering the naturallyoccurring abundance of the various isotopes (atoms that differ only in the number of neutrons)of that element, in atomic mass units (amu). A protons or a neutron represents about 1 amu,while electrons possess almost negligible mass (~1/2000th of one amu) . For example,carbon-13 is a form of carbon containing an extra neutron compared with the more commoncarbon-12 and therefore has an atomic mass of 13 rather than 12. Because carbon on earth is~1.1% C-13 + ~98.9% C-12, carbon’s atomic mass is listed as 12.011.


= =

Covalent Bonding Basics

= =

HOH

OH

H

C

O

O

C

O

O

C

C

H

H

C

C

H

H

• Single covalent bond = two shared valence electrons• Double bond = four shared valence electrons• Triple bond = six shared valence electrons

Water =

Carbon dioxide = Acetylene =

Typical covalent bondlength: ~1 Å (10-10 m)

Covalent bonding

There are a variety of ways to connect the atoms represented on the periodic table to formmolecules; these connections are called chemical bonds. Most of the chemical bonds betweenatoms in the molecules of life are covalent bonds. In a covalent bond, valence electrons areshared by two atoms. In a single covalent bond, represented by a single short line connectingtwo atoms, one pair of electrons is shared. For example, a water molecule is made of oneoxygen atom connected to each of two hydrogen atoms through single covalent bonds. In adouble covalent bond (represented by two stacked lines between atoms), the two bondedatoms share two pairs of electrons, and in a triple covalent bond (represented by three stackedlines), three pairs of electrons are shared. Since non-covalent chemical bonds cannot formdouble or triple bonds, the terms “double bond” and “triple bond” are used to refer to doublecovalent and triple covalent bonds. Lewis dot structures of molecules represent covalentbonds as pairs of dots located between two adjacent atoms. While Lewis structures ofcovalent bonds are useful for reminding you that single, double, and triple covalent bondscontain two, four, and six valence electrons, respectively, they are not as easy to draw or asefficient at conveying the structure of bonds in a molecule as the much more common use ofsimple single, double, or triple lines.


The Octet Rule

C

Carbon

N

Nitrogen

O

Oxygen

Methane

C

H

H

H HN

H

H H

Ammonia Water

O

H

H

H

H H

OH

H

CH

H

H

H NH

H

H

H H

Net result:Neutral C makes 4 bondsNeutral N makes 3 bondsNeutral O makes 2 bondsNeutral H makes 1 bond

H

4 valenceelectrons

5 valenceelectrons

6 valenceelectrons

8 valenceelectrons

8 valenceelectrons

8 valenceelectrons

• Many period 2 & 3 atoms tend to form bonds until they have 8 total valenceelectrons Hydrogen

(period 1)

Different types of atoms have different number of valence electrons available for covalent bonding.Forming a covalent bond is energetically favorable, and therefore each type of atom tends to form themaximum number of covalent bonds such that all of its valence electrons participate in covalent bonds.Many “mid-sized” atoms in the second and third rows of the periodic table (including carbon, nitrogen,oxygen, and sometimes sulfur and phosphorus) also have a strong tendency to arrange their electrons insuch a way that a total of eight valence electrons are associated with every atom; this is the so-called“octet rule”. By keeping in mind these two simply tendencies of atoms, you will be able to predict a greatdeal about their chemical and biological behavior.

As a result of these principles, every type of atom tends to form a specific number of covalent bonds.A neutral hydrogen atom contains only one electron, and therefore it is no surprise that it can form onlyone covalent bond, such as the bond in hydrogen gas (H–H). Neutral carbon has four valence electrons,and therefore tends to form four covalent bonds. When a carbon atom forms four single covalent bondsto four other atoms (such as in CH4, or methane), it has both engaged all of its valence electrons incovalent bonds and has also satisfied the octet rule. Any combination of single, double, or triple bonds toa carbon atom (such as in O=C=O, or carbon dioxide) that results in four total bonds will also meet thesecriteria. Nitrogen has five valence electrons but cannot make five covalent bonds without violating theoctet rule (since five bonds would cause ten electrons to be associated with one atom). Therefore,neutral nitrogen strongly prefers to make three covalent bonds (representing a total of six sharedelectrons) and to keep the remaining pair of valence electrons to itself. This pair of non-bonded electronsis commonly known as a lone pair. As a result, a nitrogen that participates in a total of three covalentbonds will fulfill the octet rule (six shared electrons + two non-bonded electrons = eight electrons total).If you repeat this analysis for neutral oxygen, which has six valence electrons, you will come to theconclusion that oxygen tends to participate in two covalent bonds and to keep two pairs of non-bondedelectrons to itself.


Formal Charges on Atoms• To satisfy these bonding principles, atoms of the same element that make

different numbers of covalent bonds have different charged states• Charge = group number – # of covalent bonds – # of non-bonded electrons

OC

+

H Cl••••

••

H Cl••••

•••• –

•• ••NH

••

OC

•• ••

NH

••

O•••• H

••–

••O••••H–

OC

•• ••

NH

•• –

+

(hydrochloric acid,In your stomach) (an important form of the

peptide bond)

(the first step in the reactioncatalyzed by HIV protease)

Note how each non-hydrogen atomabove still satisfies the octet rule!

Group 6 – 2 covalent bonds– 4 non-bonded electrons

= 0 net chargeGroup 6 – 1 covalent bond– 6 non-bonded electrons

= –1 net charge

Group 5 – 4 covalent bonds– 0 non-bonded electrons

= +1 net charge

When an atom makes an unusual number of covalent bonds while still satisfying the octetrule, the outcome is usually that the atom is no longer neutral but instead adopts a formalcharge (typically +1 or -1 because the number of electrons associated with the atom no longerequals its number of protons). As you will see throughout this course, these formal chargesoften play crucial roles in defining the biological properties of a molecule. A formal chargeusually arises when a covalent bond is broken, and both of the electrons making up that covalentbond end up on one of the two atoms. As a result, one of the two atoms is left with a positivecharge (has lost a previously shared electron), and the other atom is left with a negative charge(has gained a previously shared electron). For example, when the covalent bond between H andCl in H-Cl (hydrochloric acid, found in your stomach) is broken, both electrons end up on the Clatom, resulting in Cl– (which satisfies the octet rule).

A formula to calculate the formal charge on any atom is as follows: formal charge = # ofvalence electrons in that atom type when neutral – # of covalent bonds to the atom – # of non-bonded electrons associated with the atom. Since the # of valence electrons when neutral issimply the group number of the atom for the atoms commonly found in the molecules of life, thisformula reduces to: formal charge = group # – # of covalent bonds – # of non-bondedelectrons. Convince yourself that the formal charges of all the atoms in the figures above arecorrect.

As you become familiar with atoms and molecules, you will learn to intuitively and instantlyrecognize the correct formal charge on an atom simply by examining the type of atom, thenumber of covalent bonds made to the atom, how many lone pairs are present. For example, anoxygen atom that makes one covalent bond and has three lone pairs has a formal charge of –1;likewise a nitrogen atom that makes four bonds and has no lone pairs has a formal charge of +1.


1st bond

2nd bond

C

H

H

H

CH

H

H

Only Single Bonds Can Rotate Freely

• Single bonds (but not double or triple bonds) can rotate withoutdisrupting electron overlap, leading to different conformations

C

H

C

H

H

H

Single bond Double bond

The vast majority of the molecules of life can adopt multiple conformations. We’ve alreadyseen the example of DNA existing in either a B-form (right-handed helix) conformation, or in aZ-form (left-handed helix) conformation. What property of chemical bonds makes it possiblefor a single molecule to adopt different conformations?

The answer is that single covalent bonds can rotate, provided that the bonded atoms arenot geometrically constrained from bond rotation (for example, by being in a ring). Covalentbonds are stable so long as the electron orbitals containing the shared pairs of valenceelectrons can overlap in space. In a single bond, this overlap lies along the line connecting thetwo nuclei of the bonded atoms (the orange orbitals above). Even if these two atoms (and theother groups connected to them) were to rotate with respect to each other, the valenceelectrons participating in a single covalent bond can still overlap and the bond remains intact.Rotating a single bond causes a change in bond angle.

The situation is quite different, however, when we consider double or triple bonds. In thesecases, the electron orbitals that must overlap for the double or triple bond to exist lie bothalong as well as outside of the line connecting the two bonded atoms’ nuclei (the purpleorbitals above). As a result, rotating a double or triple bond destroys the overlap of valenceelectrons that make up the bond, breaking the double or triple bond.

You can now understand which molecules— and even which parts of a molecule— canundergo bond rotation, leading to different conformations.


OOPO

OH

O

N

N

N

N

NH2

O

=

=

Bonding Rules Apply to All Structures(Large and Small)

Hydrogens are oftenomitted for clarity...

C C

C

O

C

COP

O

O

O

NC

CN

C

NC

NC

N

O

H H

H

H

H

H

H H

H

H

HH

DNA

Octet Expansion:Phosphorous = 3 or 5 bonds

Sulfur = 2, 4, or 6 bonds

We are now ready to appreciate the covalent structure of the molecules of life. If youclosely examine the structure of DNA, you will see that each of the carbon, nitrogen, oxygen,and hydrogen atoms in DNA obey the principles of covalent bonding described above. Youmay also notice that the phosphorus atoms in DNA each make a total of five bondsrepresenting ten electrons, a violation of the octet rule. In the case of larger atoms (period 3and higher) such as phosphorus and sulfur, more than eight valence electrons can beassociated with each atom. This is referred to as an “octet expansion”.

Octet expansion is possible because the outermost shell of electrons for elements starting inperiod 3 can hold more than eight electrons. While the period 1 atoms H and He have only the“1s” electron orbital available (which holds a maximum of two electrons), the outermost shellsof period 2 atoms are made of 2s and 2p orbitals (which hold up to two and up to sixelectrons, respectively, for a total of eight valence electrons). The outer shells of period threeatoms, however, consist of 3s (up to two electrons), 3p (up to six electrons), and 3d (up to tenelectrons) orbitals, which can collectively hold up to 18 electrons.


Lectures 2 & 3: An introduction to the molecules of life1. The molecules of life comprise macromolecules and small molecules2. Understanding the molecules of life: chemical structures and bonding

a. Connectivity versus conformationb. The nature of atomsc. Covalent bonding and formal chargesd. Ionic bondinge. Electronegativity and hydrogen bondsf. The bonding continuum and bond polarity

3. Organic molecules and how to draw thema. The molecules of life are organic moleculesb. The geometries of organic moleculesc. Drawing organic moleculesd. Understanding “arrow-pushing” notation

4. Stereochemistry and the molecules of lifea. Stereoisomers and enantiomersb. Drawing stereoisomersc. The tragedy of thalidomided. Geometric isomerse. The role of geometric isomers in vision

5. The molecular components of HIV


Ionic Bonding

Na Cl Na Cl -

attraction

+

+ -

repulsion

--

A molecule that haslost an electron bonds to a

molecule that has gained one.These charged molecules are called ions.

(both ions satisfythe octet rule)

Ionic bonding

In addition to the covalent bond, a second type of chemical bond is the ionic bond. An ionicbond exists between atoms that are positively and negatively charged. These charged atoms,or ions, each have an unequal number of protons and electrons, accounting for their overallpositive (more protons than electrons) or negative (more electrons than protons) charge.Since positively charged matter is attracted to negatively charged matter, positively chargedions (also called cations) tend to associate with negatively charged ions (also called anions).This association is an ionic bond. Perhaps the simplest example of an ionic bond is commontable salt, or sodium chloride. In table salt, each sodium atom (which ordinarily has onevalence electron) donates that electron to a neighboring chlorine atom (which ordinarily hasseven valence electrons). The resulting Na+ cation has lost its original valence electron,leaving it with a new outer shell of eight electrons that were previously buried inside the shelloccupied by the original valence electron. The Na+ cation therefore satisfies the octet rule, anenergetically favorable situation. Because the Na+ cation now contains 11 protons but only 10electrons, it acquires a single positive overall charge. Conversely, the Cl– anion ends up witheight valence electrons (also satisfying the octet rule) and a single net negative charge.


Ionic Bonds in Protein-DNA Complexes

OBase

O

P

OBase

O

O

O

O

-N

H

H

H

NH

OProtein DNA

+

Ionic

Bond

• Ionic bonds between oppositely charged groups among the molecules of life mediatemany interactions between molecules in living systems

Ionic bonds also play important roles in the molecules of life. Frequently it is favorable foran atom in DNA, RNA, proteins, or small molecules to become positively or negatively chargedby forming or breaking covalent bonds with a proton (H+). When this happens, the resultingcharged atom will frequently form an ionic bond with an oppositely charged atom from asecond molecule. For example, the oxygen atoms bonded to the phosphorus atoms in DNAand RNA are negatively charged under most conditions. Even though these charges are oftendrawn as naked negative charges when depicting the structure of DNA and RNA on paper, inreality these charged groups are almost always associated with a cation such as Na+ or Mg2+

through an ionic bond. In isolation an ionic bond can be quite strong, often exceeding thestrength of a single covalent bond. In water, however, ionic bonds tend to rapidly dissociateas the component ions form favorable interactions with water; these interactions are the basisof water’s ability to rapidly dissolve most salts. A water-based liquid containing dissolvedmatter (such as salt, DNA, RNA, proteins, or small molecules) is called an aqueous solution.


Electronegativity & PolarityElectronegativity describes the tendency of

an atom to pull electrons towards itself...

H H

H F

Small difference inelectronegativity =

nonpolar bond

Large difference inelectronegativity =

polar bond

δ–δ+• Differences in electronegativity impart partial positive and

negative charges on bonded atoms

Electronegativity and hydrogen bonds

A third type of bond between two atoms is a hydrogen bond. Hydrogen bonds do notinvolve the sharing or transfer of electrons, although the same principle of electrostaticattraction underlies both ionic bonds and hydrogen bonds. To understand hydrogen bondingrequires understanding one of the ways in which atoms are classified: electronegativity.Electronegativity is a measure of an atom’s ability to steal electron density from other atoms towhich it is bonded. Four atoms of the periodic table are particularly electronegative: fluorine,oxygen, nitrogen, and chlorine; these four atoms have an unusual ability to draw some of theelectron density away from neighboring atoms in the same molecule. As a result, these fouratoms acquire a partial negative charge reflecting the fact that, on the average, slightly moreelectrons exist around these electronegative atoms than the number of protons in their nuclei.Robbed of some of their electron density, the neighboring atoms conversely acquire a partialpositive charge.


Hydrogen Bonding

+!

+!

!-

partialpositivecharge

partialnegativecharge

• Hydrogen bonding arises from the electrostatic attraction between partial positiveand partial negative charges created by bonding to electronegative atoms

“donor”

“acceptor”

hydrogen bond

A hydrogen bond is an electrostatic attraction between (i) an electronegative F, O, N, or Clatom (which is partially negatively charged) possessing at least one non-bonded pair ofelectrons and (ii) a hydrogen atom (which is partially positively charged) bonded to a differentF, O, N, or Cl atom. The group of atoms providing the participating hydrogen atom is calledthe hydrogen bond donor, while the group of atoms providing the participating non-bondedpair of electrons is called the hydrogen bond acceptor. Hydrogen bonds are typically muchweaker than covalent or ionic bonds, but they nevertheless play a central role in the moleculesof life. Indeed, the solvent of living systems, water, has many special properties that arisefrom the ability of water molecules to form extensive networks of hydrogen bonds with otherwater molecules.


Hydrogen Bonding in DNA

O

O

N

N

N

N

N

O

P

OO

O O

PO

O

O

O

NN

O

H

O

CH3H

H

One of the most famous examples of hydrogen bonding in living systems is the hydrogenbonding that occurs when the DNA bases pair with one another to form the double helicalstructures we saw earlier in this lecture. We will discuss in detail the molecular basis of DNAstructure later in this course.

The bonding continuum and bond polarity

In reality, chemical bonds do not always neatly fall into covalent, ionic, or hydrogen bondcategories but rather are characterized by different degrees of electron sharing (or protonsharing in the case of hydrogen bonds). Bonds between atoms of identical electronegativity(such as C–C) are entirely covalent in character, while bonds between atoms of greatlydiffering electronegativity (such as Na and Cl) are ionic in character.

In between these extremes are bonds in which electrons are shared with varying degrees ofequality. When two atoms of differing electronegativity form a bond, the more electronegativeatom tends to hoard more of the electron density between the atoms than the lesselectronegative atom. This unequal sharing results in a polar bond. Because the atomsparticipating in a polar bond have differing abilities to hoard electrons, the moreelectronegative atom acquires a partial negative charge, while the less electronegative atomacquires a partial positive charge. Sometimes these partial charges are represented with thesymbols δ+ and δ–. Conversely, bonds between atoms of similar electronegativity are non-polar bonds and do not exhibit this asymmetric distribution of electron density.

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Life Sciences 1a - Harvard Universitysites.fas.harvard.edu/~lsci1a/9-21notes.pdf · Unlike the macromolecules of life, the small molecules used by living systems are extremely diverse