1Chapter 1Bonding & Molecular Structure

Lewis Structures, Resonance Theory, Molecular Orbitals, and Hybridization

What is Organic Chemistry?

1780: Organic compounds very complex and only obtained from living sources (vitalism)

Vitalism: Belief that a magic vital force, present in plants and animals, is necessary for the synthesis of organic compounds

1789: Antoine Laurent Lavoisier observed that organic compounds are composed primarily of carbon and hydrogen

1828: Friedrich Whler synthesized an organic compound (urea) from inorganic compounds (lead cyanate and ammonium hydroxide).

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Modern Organic Chemistry is the study of Carbon Compounds

STRUCTUREDetermining the way in

which atoms are put together in space to

form complex molecules

MECHANISMUnderstanding the

reactivity of molecules: How and why chemical

reactions take place

SYNTHESISBuilding complex

molecules from simple molecules using

chemical reactions

The study and practice of organic chemistry can be subdivided into three areas:

>95% of All Known Compounds Composed of Carbon

Organic Chemistry Crucial to Our Way of Life: Dyes, Materials (Polymers), Petroleum, Pesticides, Medications, a Molecular Understanding of the Humane Body, OUR BODIES.

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ImatinibChronic myelogenous leukemia

Why is an entire subject directed towards studying a single element of the periodic table?

Why Carbon?

Carbon forms a variety of strong covalent bonds to itself and other atoms (O, N, P, S, F, Cl, Br, I).

This allows organic compounds to be structurally diverse. Carbon can bond to itself in many different configurations, including chains, rings, and branched structures of varying size and complexity.

Why Does Carbon Bond in This Way? 4

DNA Bases Hormones

Carbohydrates Amino Acids

Review of Atomic Structure

Electron cloud is ~10,000 times the diameter of its nucleus

Nucleus is made up of protons and neutronsProtons are positive chargedNeutrons are uncharged (electrically neutral)

Electrons are negatively charged and are in motion

~99.9% of the mass of an atom comes from the nucleus

Atomic Number (Z) = to the number of protons in the nucleusThe atomic number also equals the number of electrons surrounding the nucleus

Electrons surrounding the nucleus exist in shells of increasing energy and at increasing distances from the nucleus

Valence Electrons: exist in the valence shell, and are the outermost electrons that are used in making chemical bonds

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I. Review of Lewis Bonding Theory

G. N. Lewis

an atom is most stable if its outer shell is either filled or contains eight electrons, and it has no electrons of higher energy

Octet Rule: an atom will give up, accept, or share electrons in order to achieve a filled outer shell or an outer shell that contains eight electrons

Atoms form bonds with other atoms to achieve a filled valence shell (eight electrons), which gives them a stable Noble gas configurationParticularly common among second-row elements of Groups 1A-7A

Group # IA IIA IIIA IVA VA VIA VIIA VIIIA

# Valence e- 1 2 3 4 5 6 7 8

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

Full valence shell of electrons and are thus quite stable/unreactive 6

The Correlation Between Chemical Bonds and Electronegativity

Electronegativity a measure of an atoms attraction for electrons that it shares in a chemical bond with another atom

Ionic Bond electron(s) are completely transferred from one atom to another

Covalent Bond electrons are shared between two atoms

Classification of Chemical Bonds

EN Between Bonded Atoms Type of Bond

> 1.9 Ionic

0.5 to 1.9 Polar Covalent

< 0.5 Nonpolar Covalent

H2.1

Li1.0

Be1.5

B2.0

C2.5

N3.0

O3.5

F4.0

Na0.9

Mg1.2

Al1.5

Si1.8

P2.1

S2.5

Cl3.0

K0.8

Br2.8

I2.5

Increasing Electronegativity

Increasing Electronegativity

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Ionic Bonding

Between atoms of widely different electronegativity (EN>2); usually a metal and a nonmetal (results in the formation of ions, a cation(+) and an anion (-))

Atoms held together by electrostatic attraction, not electron sharing

e.g. LiF is ionic (EN: Li = 1.0 and F = 4.0)

Using Lewis Dot Structures: showing valence electrons

H = +123.6 kcal/mol(ionization potential)

H = -78.3 kcal/mol(electron affinity)

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Covalent Bonding (Electron Sharing) Very important in organic molecules! Between atoms of similar electronegativity; usually non-metallic

e.g. CH4 (methane) is nonpolar covalent (EN

Covalent Bonding (Electron Sharing)

e.g. CCl4 is polar covalent (EN 0.5 to 1.9) [EN: Cl = 3.5 and C = 2.5] .Why?

each chlorine atom shares one valence electron with carbon so that every atom has a filled octet

Polar Covalent Bond: unequal sharing of bonding electrons

more electronegative atom gains a greater fraction of the shared electrons and acquires a partial negative charge (-) The less electronegative atom has a smaller fraction of the shared electrons and acquires a partial positive charge (+)

each chlorine atom still has three lone pairs of electronsLone Pair: unshared pair of electrons; non-bonding pair of electrons

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How can you tell how many bonds and how many lone pairs an atom will have?

Count the Valence Electrons!

Neutral atom # valence es # bonds # lone pairs

H

C

N

O

F

Second row elements want to be surrounded by eight valence electrons (an octet)

e.g. Nitrogen wants three more electrons (3 bonds, one lone pair)

e.g. Carbon wants four more electrons (4 bonds)

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Multiple Covalent Bonding

Two atoms can share more than one pair of electrons to gain a filled valence shell(very common in organic molecules)

e.g. Ethylene (C2H4)

e.g. Acetylene (C2H2)

e.g. Hydrogen cyanide (HCN)

Each line represents one shared electron pairs12

Carbon Hydrogen Carbon Oxygen Carbon Oxygen Oxygen Oxygen

Group#

# non-Bonding es

(# shared es)

Formal Charge

Formal Charges and How to Calculate Them Not all atoms are neutral in a Lewis or Kekule structure Formal charges help chemists to keep track of the placement of electrons in molecules Does not indicate that all of the charge is actually localized on one atom

Formal Charge = (group #) (# non-bonding es) 1/2(# shared es)

Always indicate formal charge on problem sets and exams!

Which atom has the formal charge?

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Summary of Bonding and Formal Charges

Each line represents one shared electron pairs

The formal charges and the octet rule must coexist

Be able to recognize common bonding patterns

10 electrons in valence shell of Nitrogen

CORRECT NOT CORRECT 14

Line-Angle Formulas: Short-Hand for Chemists(Drawing Complex Molecules Quickly)

Electron-dot formulaLewis Structure

Condensed formula Line-Angle formulaDash formulaKekule drawing

Testosterone(not so easy!)

much easier!

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Brevetoxin A (!!)

Line-Angle Formulas: Short-Hand for Chemists(Drawing Complex Molecules Quickly)

Electron-dot formulaLewis Structure

CH3CH2CH2OH

Condensed formula Line-Angle formulaDash formulaKekule drawing

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Rules for Drawing Line-Angle Formulas

Bonds are represented by lines (one line = two shared electrons)

Each bend in a line or terminus of a line represents a carbon atom(unless another group is shown explicitly)

Do not draw carbon (C) atoms, except optionally at termini (for aesthetics)

Do not draw hydrogen (H) atoms, unless showing stereochemistry (more detailed explanation next slide)

Assume enough C-H bonds to give each carbon atom four bonds (an octet)(unless a charge is indicated)

Draw heteroatoms (N, O, S, P, F, Cl, Br, I)

Explicitly write hydrogen atoms bonded to heteroatoms

e.g. isopropanol: e.g. cyclohexanone

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Organic Molecules Are Not Flat, They Are Three-Dimensional!

Tetra-Substituted Carbon Is Tetrahedral (more on this later)

lines: in the plane of the paperdashes: going back into the paper (away from you)wedges: coming out of the paper (toward you) 18

Quantum and Wave Mechanics

Moving Particles Exhibit the Properties of a Wave

Each wave function () corresponds to a different energy state for an electron

Each energy state is a sublevel where one or two electrons can reside

The energy associated with the state of an electron can be calculated from the wave function

The relative probability of finding an electron in a given region of space can be calculated from the wave function

The phase sign of a wave equation indicates whether the solution is positive or negative when calculated for a given point in space relative to the nucleus

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Constructive interference occurs when wave functions with the same phase sign interact. There is a reinforcing effect and the amplitude of the wave function increases.

Destructive interference occurs when wave functions with opposite phase signs interact. There is a subtractive effect and the amplitude of the wave function goes to zero or changes sign.

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Molecular Orbital Theory

Electrons Are Waves!They Exist as 3-D Standing Waves (Orbitals)

Orbital: a region of space where the probability of finding an electron is high.

Atomic Orbitals: unhybridized orbitals on an atom (s, p, d, f)

Linear Combination of Atomic Orbitals: individual wave functions (orbitals) combine to form hybrid atomic orbitals (sp3, sp2, and sp) and molecular orbitals (, *, , *)

Hybrid Atomic Orbital: combination of atomic orbitals from the same atom

Molecular Orbital: combination of atomic orbitals from different atoms

Conservation of Orbitals: when you add orbitals together, you always end up with the same number of orbitals that you started with.

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Atomic Orbitals

The s- and p-orbitals are the most important in organic compounds.

Orbital pictures are actually electron probability cloudsThe volumes shown contain the electron 90-95% of the time

s-orbitals: spherical, electrons held close the nucleus, one sign

p-orbitals: two lobes with opposite signs, electrons further from nucleusThere are three degenerate (orbitals of equal energy) 2p orbitals that are higher in energy than the 2s orbital

Remember, the sign of the orbital does not indicate charge. It represents the sign of the wavefunction and lets us think qualitatively about whether orbital interactions are constructive (bonding) or destructive (anti-bonding)

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Electron Configuration

Rule 1. The Aufbau Principle. Orbitals fill in order of increasing energy, from lowest to highest. Orbitals fill in the order 1s, 2s, 2p, 3s, 3p, and so on.

Rule 2. The Pauli Exclusion Principle. Only two electrons can occupy an orbital and their spins must be paired. The quantum mechanical

property of spin gives an electron a tiny magnetic field

When their tiny magnetic fields are aligned N-S, the electron spins are paired

Rule 3. Hunds Rule. When orbitals of equal energy (degenerate orbitals) are available we add one electron to each orbital before a second electron is added to any one of them; and the spins of the single electrons in the degenerate orbitals should be aligned. Partially filling orbitals as much as possible minimizes electrostatic repulsion between electrons.

Energy Level Diagram 23

Molecular Orbital (MO)Represents the region of space where one or two electrons are being shared by two

different atoms (covalent bond) and are likely to be found

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The number of molecular orbitals that are created is equal to the number of atomic orbitals that combine

A bonding molecular orbital results when two orbitals of the same phase overlap

An antibonding molecular orbital results when two orbitals of opposite phase overlap

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Sigma-Bonding ()

Sigma-bonding orbitals are cylindrically symmetrical molecular orbitals. Electron density is centered along the axis of the bond. Single bonds are sigma-bonds.

e.g. H2 is the simplest sigma-bond

Bonding: (+/+ or -/-) electron density centered between nuclei

Anti-bonding: (+/-) generally has a node between nuclei

Node: area of zero electron density

In stable bonding situations, usually only the bonding orbitals (, ) are occupied

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Sigma bonds arent necessarily between two s-orbitals

s + p p + p

These are all examples of single bonds27

Pi-Bonding () Pi-bonding orbitals are not cylindrically symmetrical Electron density is located above and below the axis of the bond Double and triple bonds are pi-bonds

Double bond = +

e.g. Ethylene

To simplify drawing of orbitals 28

C C

H

H

H

H

-bond

(-orbitals overlap)

3-Dimensional View

Planar structure

Carbon with (3 + 1) bonds

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Orbital Overlap Orbitals must have the correct symmetry to overlap Orthogonal orbitals do not overlap

Good Overlap Orthogonal: No Overlap

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If all bonding occurred between simple s-and p-orbitals, then all bond angles would

be approximately 90

3-D representations of the 2px, 2py, and 2pz atomic orbitals and their orientation in space relative to one another

We know that isnt true!

Most bond angles in organic molecules are ~ 109, ~120, and ~180

How do we account for this?

Valence Shell Electron Pair Repulsion (VSEPR)

Electrons repel each other!

Lone pairs and bonds want to be a far apart as possible31

Simply

Di-substitutedlinear (180)

e.g. acetylene

Tri-substitutedtrigonal planar (120)

e.g. ethylene

Tetra-substitutedtetrahedral (109)

e.g. methane

But, how do we think about this in terms of orbitals?

In 1930, Linus Pauling introduced a theory that combines VSEPR with quantum mechanics (orbitals)

HYBRIDIZATION

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Hybridization (Linear Combination of Atomic Orbitals)

Remember Conservation of Orbitals!When you add orbitals together, you always end up

with the same number of orbitals that you started with.

Atomic orbitals on the same atom combine to formhybride atomic orbitals.

Why?

Hybrid orbitals are more directional, so they have more effective bonding interactions

Second row elements hybridize using their s- and p-orbitals (sp, sp2, sp3)

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sp3 Hybridization (Tetrahedral)

Ground State Excited State sp3-Hybridized State

Promotion of elections Hybridization

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sp2 Hybridization (Trigonal planar)

Ground State Excited State sp2-Hybridized State

Promotion of elections Hybridization

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sp2 Hybridization (Trigonal planar)

C

Three sp2-orbitals

Enhanced electron density in bonding

regions

3 sp2-orbitals1 p-orbital

For simplicity, can leave out small back lobes.

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Model of Bonding Molecular Orbitals of Ethylene (H2C=CH2)

Sigma-bonds () and lone pairs involve hybrid orbitals

Pi-bonds () involve unhybridized p-orbitals

Electrons of the bond have greater energy than electrons of the bond

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sp Hybridization (Linear)

Ground State Excited State sp2-Hybridized State

Promotion of elections Hybridization

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sp Hybridization (Linear)

Two sp-orbitals

Enhanced electron density in bonding

regions

2 sp-orbitals2 p-orbital

For simplicity, can leave out small back lobes.

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Model of Bonding Molecular Orbitals of Ethyne (HCCH)

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Drawing Hybridized Orbitals

sp3 Hybridized Orbital sp2 Hybridized Orbital sp Hybridized Orbital

Drawing Bonding Molecular Orbitals of Propene (CH3-CH=CH2)

For simplicity, draw lines connecting p-orbitals to represent -bonds

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Drawing Bonding Molecular Orbitals of Ethyne (HCCH)

For simplicity, draw lines connecting p-orbitals to represent -bonds

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Assigning Hybridization to Atoms in a Molecule

You need to be able to do this!

Count the hybrid atomic orbitals

# of hybrid orbitals =

# Hybrid orbitals Hybridization Geometry ~ bond angles

4 sp3 Tetrahedral 109

3 sp2 Trigonal planar 120

2 sp Linear 180

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Assigning Hybridization to Atoms in a Molecule

You need to be able to do this!

Bond Lengths of Ethyne, Ethylene, and Ethane

sp orbital 50% s character, 50% p character

sp2 orbital 33% s character, 66% p character

sp3 orbital 25% s character, 75% p character

The greater the s orbital character in one or both atoms, the shorter is the bond. This is because s orbitals are spherical and have more electron density closer to the nucleus than do p orbitals

The greater the p orbital character in one or both atoms, the longer is the bond. This is because p orbitals are lobe-shaped with electron density extending away from the nucleus

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Rotation of Ethane versus Ethylene

Sigma-bonds are cylindrically symmetrical: rotation does not disrupt bonding Sigma-bonds rotate freely

Ethane (much more on the energeticsof this rotation later!)

Pi-bonds require ovelap of the p-orbitals: rotation disrupts overlap Pi-bonds do not rotate

Ethylene

-overlap p-orbitals orthogonal:no overlap

This rotation does not occur!

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Representing Molecules

Lewis Structures: Represent atoms sharing electrons to form bonds

Line-Angle Structures: Simplify the drawing of complex molecular structures

Dashes and Wedges: Allow chemists to draw molecules in 3-D

BUT! These simplified structures do not accurately represent the electronic nature or reactivity of organic molecules!

It helps to think about electrons in motion

BUT HOW DO WE REPERESENT ELECTRONS IN MOTION?

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Curved Arrow Formalism (Arrow Pushing)

Chemists use arrows to represent the motion of electrons within and between molecules.

Double Arrow:2 electrons moving

Fishhook Arrow:1 electron moving

1. The tail of the arrow starts at the electrons that are moving (lone pair or bond)2. The head of the arrow shows where the electrons end up (lone pair or bond)

Always show the flow of electrons from a site of higher electron density to a site of lower electron density

NOT CORRECT! 49

Resonance: Electronic Motion Within a Molecule

In a Lewis structure, we draw a well-defined location for the electrons in a molecule.

The reactivity of a molecule is not always explained by one Lewis structureIn many molecules and ions (especially those containing bonds), more than one equivalent Lewis structure can be drawn which represent the same molecule

e.g. carbonate ion (CO3-2)

Molecules can be thought of as hybrids or weighted averages of two or more Lewis structures, each with a different placement of electrons

These structures, called resonance structures are not real or detectable, but they are a useful conceptual tool for understanding the reactivity of molecules

Contributing Resonance Structures Hydrid

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Rules for Drawing Resonance Structures

1) Only electrons move! The positions of nuclei must be the same in all contributing structures; that is, contributing structures differ only in the distribution of valence electrons. (Resonance occurs in the pi-system: conjugated lone pairs and pi-bonds)

2) Every resonance structure must be a valid Lewis structure

3) Keep track of lone pairs and formal charges. All contributing structures must have the same number of paired and unpaired electrons.

4) Use arrow-pushing formalism to interconvert and identify possible resonance structures.

5) Always use double-headed arrow () in between resonances structures. Resonance structures do not represent an equilibrium, do not use ( )

6) Resonance stabilization: the energy of the resonance hybrid is lower than the energy of any contributing structure.

7) Lower energy resonance structures contribute more to the overall structure of the molecule.

How do you predict when one resonance structure contributes more to the hybrid than another?

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Guidelines for Estimating the Relative Importance (Energies)of Contributing Resonance Structures

1) Filled Valence Shell: Structures in which all atoms have filled valence shells (completed octets for C, N, O, F) contribute more than those in which one or more valence shells are unfilled. Because C is the least electronegative, structures in which C has 6 electrons, 3 bonds, and a positive charge are possible (not possible with N, O, F)

Greater Contribution:both carbon and oxygen have

complete valence shells

Lesser Contribution:carbon has only six electrons

in its valence shells

2) Maximum Number of Covalent Bonds: Structures with a greater number of covalent bonds contribute more than those with fewer covalent bonds.

Greater Contribution:8 covalent bonds

Lesser Contribution:7 covalent bonds 52

Guidelines for Estimating the Relative Importance (Energies)of Contributing Resonance Structures

3) Least Separation of Unlike Charges: Structures that involve separation of unlike charges contribute less than those that do not involve charge separation because separation of charges costs energy.

Greater Contribution:no separation of unlike charges

Lesser Contribution:separation of unlike charges

4) Negative Charge on a More Electronegative atom: Structures that carry a negative charge on a more electronegative atom contribute more. Conversely, structures that carry a positive charge on a less electronegative atom contribute more.

Greater ContributionLesser Contribution Should not be drawn:violates all 4 guidelines 53

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Draw the Contributing Resonance Structures