Chapter 1Structure and Bonding

O i Ch i t

Structure and Bonding

Organic Chemistry, Third Edition

Janice Gorzynski SmithJanice Gorzynski SmithUniversity of Hawai’i

Prepared by Layne A MorschPrepared by Layne A. MorschThe University of Illinois - Springfield

Modified and Used for Basic Org Chem LectureModified and Used for Basic Org Chem Lectureby V.P. Hoven, as of August 2012

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.1

Atomic Structure

• The nucleus contains positively charged protons andp y g puncharged neutrons.

• The electron cloud is composed of negatively chargedp g y gelectrons.

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N fNo.of protons

In a neutral atom, protons = electrons

Isotopes are two atoms of the same element having a different number of neutrons.number of neutrons.

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The Periodic Table

Figure 1.2 Common elements seen in organic chemistry

• Elements in the same column have similar electronic andchemical properties.

• Elements in the same row are similar in size. 4

Atomic Orbitals• An s orbital has a sphere of electron density and is lower in

energy than the other orbitals of the same shell.energy than the other orbitals of the same shell.

• A p orbital has a dumbbell shape and contains a node (noelectron density) at the nucleus It is higher in energy than anelectron density) at the nucleus. It is higher in energy than ans orbital. 5

The First Row

• There is only one orbital in the first shell.y• Each shell can hold a maximum of two electrons.• Therefore, there are two elements in the first row:

H and HeH and He.

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The Second Row

Each element in the second row of the periodic table has fourporbitals available to accept additional electrons: one 2s orbital,and three 2p orbitals.

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• Since each of the four orbitals in the second shell holdst l t th i i it f i ht ltwo electrons, there is a maximum capacity of eight valenceelectrons for elements in the second row.

• The second row of the periodic table consists of eight• The second row of the periodic table consists of eight elements, obtained by adding electrons to the 2s and three 2p orbitals.

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Bonding

• Bonding is the joining of two atoms in a stableg j garrangement.

• Atoms can form either ionic or covalent bonds toAtoms can form either ionic or covalent bonds toattain a complete outer shell (octet rule for second rowelements, stable noble gas configuration).

• Ionic bonds result from the transfer ofIonic bonds result from the transfer ofelectrons from one element to another.

• Covalent bonds result from the sharing ofCovalent bonds result from the sharing ofelectrons between two nuclei.

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Ionic Bonding• An ionic bond generally occurs when elements on the far

left side of the periodic table combine with elements on thef i ht id i i blfar right side, ignoring noble gases.

• A positively charged cation formed from the element on theleft side attracts a negatively charged anion formed from thel t th i ht id ( di hl id N Cl)element on the right side (e.g., sodium chloride, NaCl).

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• Li loses its one electron to make Li+ which has no electrons• Li loses its one electron to make Li which has no electronsin second shell. However, it has a complete first shell.

• F gains one electron to make F- which has a filled valence• F gains one electron to make F which has a filled valenceshell (an octet of electrons), like neon. 11

Covalent Bonding

• Covalent bonding occurs with elements like carbon in themiddle of the table (e.g., CH4) with elements that havemiddle of the table (e.g., CH4) with elements that havesimilar electronegativity.

• Covalent bonds also occur between two of the sameelements from the same sides of the table (e.g., H2, Cl2).

• A covalent bond is a two-electron bond, and a compoundwith covalent bonds is called a molecule.

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Valence Electrons

• Second-row elements can have no more than eightelectrons around them. For neutral molecules, this haselectrons around them. For neutral molecules, this hastwo consequences:

• Atoms with one two three or four valence electronsAtoms with one, two, three, or four valence electronsform one, two, three, or four bonds, respectively, inneutral molecules (e.g., BF3, CH4).

• Atoms with five or more valence electrons form enoughbonds to give an octet (e.g., NH3). This results in thefollowing equation:following equation:

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Nonbonded Electrons

Figure 1.3Figure 1.3

• When second-row elements form fewer than fourbonds, their octets consist of both bonding (shared)and nonbonding (unshared) electrons. Unsharedelectrons are also called lone pairselectrons are also called lone pairs.

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Lewis Structures

Lewis structures are electron dot representations for molecules.

A solid line indicates a two-electron covalent bondelectron covalent bond.

General rules for drawing Lewis structures:General rules for drawing Lewis structures:1. Draw only the valence electrons.2. Give every second-row element no more than eight electrons.3. Give each hydrogen two electrons.

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How to Draw a Lewis StructureStep [1] Arrange atoms next to each other that you think are bonded together.

• Always place hydrogen and halogens on the periphery because• Always place hydrogen and halogens on the periphery because they form only one bond each.

Place no more atoms around an atom than the number of bonds• Place no more atoms around an atom than the number of bondsit usually forms.

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Step [2] Count the electrons.• Count the number of valence electrons from all atoms.

Add one electron for each negati e charge• Add one electron for each negative charge.• Subtract one electron for each positive charge.• This gives the total number of electrons that must

be used in drawing the Lewis structure.

Step [3] Arrange the electrons around the atoms.• Place a bond between every two atoms, giving two

electrons to each H and no more than eightto any second-row atom.to any second row atom.

• Use all remaining electrons to fill octets with lone pairs.• If all valence electrons are used and an atom does not

have an octet form multiple bondshave an octet, form multiple bonds.

Step [4] Assign formal charges to all atoms.

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Formal Charge

• Formal charge is the charge assigned to individual atoms in aLewis structure.

• Formal charge is calculated as follows:

• The number of electrons “owned” by an atom is determinedby its number of bonds and lone pairs.by its number of bonds and lone pairs.

• An atom “owns” all of its unshared electrons and half of itsshared electrons.

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The number of electrons “owned” by different atoms is indicated in the following examples:

Example 1

Example 3

Example 2Example 2

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Isomers

• Sometimes more than one arrangement of atoms (Lewisstructure) is possible for a given molecular formula.structure) is possible for a given molecular formula.

• These two compounds are called isomers.• Isomers are different molecules having the same molecular

formula Ethanol and dimethyl ether are constitutionalformula. Ethanol and dimethyl ether are constitutional isomers. 21

Exceptions to the Octet RuleElements in Groups 2A and 3A

p

El t i th Thi d RElements in the Third Row

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Resonance Structures• Some molecules cannot be adequately represented by a

single Lewis structure. For example:

• These structures are called resonance structures or resonance forms. A double-headed arrow is used to separate the two resonance structuresthe two resonance structures.

• Resonance structures are two Lewis structures having the same placement of atoms but a different arrangement ofsame placement of atoms but a different arrangement of electrons.

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• Neither resonance structure is an accurate representationfor (HCONH)¯ The true structure is a composite of bothfor (HCONH) . The true structure is a composite of bothresonance forms and is called a resonance hybrid.

• The hybrid shows characteristics of both structures.e yb d s o s c a acte st cs o bot st uctu es• Resonance allows certain electron pairs to be delocalized

over two or more atoms, and this delocalization addsstability.

• A molecule with two or more resonance forms is said to beresonance stabilizedresonance stabilized.

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Basic Principles of Resonance Theoryas c c p es o eso a ce eo y

• Resonance structures are not real. An individual resonancestructure does not accurately represent the structure of amolecule or ion. Only the hybrid does.

• Resonance structures are not in equilibrium with eachother. There is no movement of electrons from one form toanotheranother.

• Resonance structures are not isomers. Two isomers differin the arrangement of both atoms and electrons whereasin the arrangement of both atoms and electrons, whereasresonance structures differ only in the arrangement ofelectrons.

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Drawing Resonance Structures

Rule [1]: Two resonance structures differ in the position of multipleb d d b d d l t Th l t f t d i lbonds and nonbonded electrons. The placement of atoms and singlebonds always stays the same.

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Rule [2]: Two resonance structures must have the same number ofRule [2]: Two resonance structures must have the same number ofunpaired electrons.

Rule [3]: Resonance structures must be valid Lewis structures.[ ]Hydrogen must have two electrons and no second-row element canhave more than eight electrons.

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Curved Arrow Notation• Curved arrow notation is a convention that shows how electron

position differs between two resonance forms.

• Curved arrow notation shows the movement of an electron pair.The tail of the arrow always begins at the electron pair, either in abond or lone pair. The head points to where the electron pairbond or lone pair. The head points to where the electron pair“moves.”

Example 1:

E l 2Example 2:

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Occurrence of Resonance

Two different resonance structures can be drawn when a lone pair is located on an atom directly bonded to a double bondpair is located on an atom directly bonded to a double bond.

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Multiple resonance structures can also be drawn when an atom bearing a (+) charge is bonded either to a double bond or g ( ) gan atom with a lone pair.

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Resonance Hybrideso a ce yb d

• A resonance hybrid is a composite of all possibleeso a ce yb d s a co pos te o a poss b eresonance structures. In the resonance hybrid, theelectron pairs drawn in different locations in individualresonance forms are delocalized.

• When two resonance structures are different, thehybrid looks more like the “better” resonancestructure.

• The “better” resonance structure is called the majorcontributor to the hybrid, and all others are minorcontributors.

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A “better” resonance structure is the one that has more bondsand fewer charges.

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Molecular Geometry

Two variables define a molecule’s structure: bond lengthand bond angle.g

• Bond length decreases across a row of the periodictable as the size of the atom decreasestable as the size of the atom decreases.

• Bond length increases down a column of the periodictable as the size of an atom increasestable as the size of an atom increases.

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• The number of groups surrounding a particular atomd t i it t A i ith t ldetermines its geometry. A group is either an atom or a lonepair of electrons.

• The most stable arrangement keeps these groups as faraway from each other as possible. This is exemplified byValence Shell Electron Pair Repulsion (VSEPR) theory.

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Drawing Three-Dimensional Structuresg

• A solid line is used for a bond in the plane.• A wedge is used for a bond in front of the plane.• A dashed line is used for a bond behind the plane.

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A Nonbonded Pair of Electrons is Counted as a “Group”

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Varying Bond Angles

Methane (CH4)•In both NH3 and H2O, the bond angle is smaller than the theoretical tetrahedralthe theoretical tetrahedral bond angle because of repulsion of the lone pairs

fAmmonia (NH3) of electrons.

•The bonded atoms are d i t llcompressed into a smaller

space with a smaller bond angle.

Water (H2O)a g e

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Condensed Structures

• All atoms are drawn in, but the two-electron bond lines are generally omitted.g y

• Atoms are usually drawn next to the atoms to which they are bonded.

• Parentheses are used around similar groups bonded to the same atom.

• Lone pairs are omitted• Lone pairs are omitted.

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Skeletal Structures• Assume there is a carbon atom at the junction of any two

lines or at the end of any line.• Assume there are enough hydrogens around each carbon to

make it tetravalent.Draw in all heteroatoms and the hydrogens directly bonded to• Draw in all heteroatoms and the hydrogens directly bonded to them.

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• A charge on a carbon atom takes the place of one hydrogentatom.

• The charge determines the number of lone pairs. Negativelycharged carbon atoms have one lone pair and positivelycharged carbon atoms have one lone pair and positivelycharged carbon atoms have none.

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Orbitals and Bonding: Hybridization

• When the 1s orbital of one H atom overlaps with the 1sbit l f th H t i ( ) b d th torbital of another H atom, a sigma (σ) bond that

concentrates electron density between the two nuclei is formed.

• This bond is cylindrically symmetrical because the electrons forming the bond are distributed symmetrically g y yabout an imaginary line connecting the two nuclei.

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• To account for the bonding patterns observed in morecomplex molecules we must take a closer look at how thecomplex molecules, we must take a closer look at how the2s and 2p orbitals of atoms in the second row are utilized.

• In addition to its two core electrons carbon has four valence• In addition to its two core electrons, carbon has four valenceelectrons. In its ground state, carbon places two electrons inthe 2s orbital and one each in 2p orbitals.

Note: The lowest energy arrangement of electrons fort i ll d it d t tan atom is called its ground state.

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• Promotion of an electron from a 2s to a vacant 2p orbitalwould form four unpaired electrons for bonding This higherwould form four unpaired electrons for bonding. This higherenergy electron configuration is called an electronicallyexcited state.

• Carbon would form two different types of bonds: three withy2p orbitals and one with a 2s orbital. Experimental evidencepoints to carbon forming four identical bonds in methane.

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• To solve this dilemma, chemists have proposed that atomslike carbon do not use pure s and pure p orbitals in forminglike carbon do not use pure s and pure p orbitals in formingbonds. Instead, atoms use a set of new orbitals called hybridorbitals.

• Hybridization is the combination of two or more atomicorbitals to form the same number of hybrid orbitals, eachhaving the same shape and energy.

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Shape and Orientation of sp3 Hybrid Orbitals

• The mixing of a spherical 2s orbital and three dumbbellshaped 2p orbitals together produces four hybrid orbitals,shaped 2p orbitals together produces four hybrid orbitals,each having one large lobe and one small lobe.

• The four hybrid orbitals are oriented toward the corners of at t h d d f f i l t b dtetrahedron and form four equivalent bonds.

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Bonding Using sp3 Hybrid Orbitals

• Each bond in CH4 is formed by overlap of an sp3 hybridorbital of carbon with a 1s orbital of hydrogen.

• These four bonds point to the corners of a tetrahedron.p

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Other Hybridization Patterns• One 2s orbital + three 2p orbitals = four sp3 hybrid orbitals.• One 2s orbital + two 2p orbitals = three sp2 hybrid orbitals.• One 2s orbital + one 2p orbital = two sp hybrid orbitals.One 2s orbital + one 2p orbital two sp hybrid orbitals.

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Determining Hybridization

• Count the number of groups (atoms and nonbondedelectron pairs) around the atomelectron pairs) around the atom.

• The number of groups corresponds to the number ofatomic orbitals that must be hybridized to form theatomic orbitals that must be hybridized to form thehybrid orbitals.

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Hybridization Examples

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sp3 Hybridization and Bonding in Organic Molecules

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Rotation around C-C Bond

• Making a model of ethane illustrates one additionalf t b t it t tfeature about its structure.

• Rotation occurs around the central C–C σ bond.

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sp2 Hybridization and Bonding in Organic Molecules

Each carbon is trigonal planarEach carbon is trigonal planar.

Each carbon is sp2 hybridized.

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Unhybridized 2pz orbital (gray), which is perpendicular to the plane of the hybrid (green) orbitalsto the plane of the hybrid (green) orbitals.

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Sigma bond (s) – electron density between the 2 atoms

Pi b d ( ) l t d it b d b l l f l iPi bond (p) – electron density above and below plane of nuclei of the bonding atoms 59

No Rotation around C=C Bond

• Unlike the C–C bond in ethane, rotation about the C–C doublebond in ethylene is restricted.

• It can only occur if the π bond first breaks and then reforms, aprocess that requires considerable energy.

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sp Hybridization and Bonding in Organic Molecules

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Each carbon atom has two unhybridized 2p orbitals that are perpendicular to each otherare perpendicular to each other and to the sp hybrid orbitals.

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Summary of Covalent Bonding

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Bond Length and Bond Strength• As the number of electrons between two nuclei increases,

bonds become shorter and stronger.

• The length and strength of C–H bonds vary depending on• The length and strength of C–H bonds vary depending onthe hybridization of the carbon atom.

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Percent s-CharacterPercent s Character

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Electronegativity

Electronegativity is a measure of an atom’s attraction forelectrons in a bondelectrons in a bond.

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Bond Polarity• Electronegativity values are used are used as a guideline to indicate

whether the electrons in a bond are equally shared or unequallyh d b Wh l ll h d hshared between two atoms. When electrons are equally shared, the

bond is nonpolar.

• A carbon—carbon bond is nonpolar. The same is true wheneverdiff h i i il l i i i b d dtwo different atoms having similar electronegativities are bonded

together.

• C H bonds are considered to be nonpolar because the• C—H bonds are considered to be nonpolar because theelectronegativity difference between C and H is small. 69

• Bonding between atoms of different electronegativity valueslt i l h i f l tresults in unequal sharing of electrons.

the head of the arrowThe tail of the arrow is the head of the arrowpointing toward the moreelectronegative element

The tail of the arrow is drawn at the less electronegative element

• In the C–O bond, the electrons are pulled away from C (2.5)toward O (3.4), the element of higher electronegativity. Thebond is polar or polar covalent. The bond is said to have adipole; that is separation of chargedipole; that is, separation of charge.

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Polarity of MoleculesUse the following procedure to determine if a molecule has anet dipole:

• Use electronegativity differences to identify all of the polar bondsand the directions of the bond dipoles.

• Determine the geometry around individual atoms by counting• Determine the geometry around individual atoms by countinggroups, and decide if individual dipoles cancel or reinforce eachother in space.

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A polar molecule has either one polar bond, or two or morebond dipoles that reinforce each other.p

A nonpolar molecule has either no polar bonds, or two ormore bond dipoles that cancel.

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