Vollhardt 6e Lecture PowerPoints - Chapter 1

8/11/2019 Vollhardt 6e Lecture PowerPoints - Chapter 1

1/56

CHAPTER 1

Structure and Bonding in Organic

Molecules


2/56

The Scope of Organic Chemistry: An Overview

1-1

Functional groups determine the reactivity of organic molecules.Alkanes No functional groups, only carbon and hydrogen.(Chapter 2)

Alkane ReactionsAlkane bond strengths and reactions.(Chapter 3)

Cyclic AlkanesNew properties and changes in reactivity(Chapter 4)

StereoisomerismSame connectivity different relativepositioning of substituents in space (Chapter 5)

Haloalkanes Substitution Reactions and Elimination Reactions(Chapters 6 and 7)

Alkynes C-C triple bonds (Chapter 13)

Aldehydes and KetonesCarbonyl Compounds C=O doublebonds. (Chapters 16 and 17)


3/56


1-1

Amines Nitrogen containing functional group (Chapter 21)Tools For IdentificationSpectroscopy (Chapters 10, 11, 14 and20)

Carbohydrates and Amino AcidsMultiple Functional Groups(Chapters 24 and 26)


4/56


1-1

Synthesis is the making of new moleculesWhlers Synthesis of Urea:

SynthesisConstruct complex organic chemicals from simpler,more readily available ones (Chapter 8).

Reactions are the vocabulary, and mechanisms are the grammar oforganic chemistry

Reactants (Substrates) Starting compoundsProducts

Reaction MechanismUnderlying details of a reaction

Reaction IntermediateChemical species formed and thendestroyed on the pathway between reactants and products.


5/56

Coulomb Forces: A Simplified View of Bonding

1-2

Bondsare made by simultaneous coulombic attraction andelectron exchange.

When two atoms approach, the electrons of one are attracted bythe protons of the other and vice-versa.

2

(+) charge ( ) charge

Attracting Force = constant

distance

Energy is released as the two atoms approach each other.When the atoms get too close together the energy begins torise again due to repulsions between the two nuclei and thetwo sets of electrons.


6/56

This minimum energy is called thebond strength, and thedistance between the two nuclei at this point is called thebond length.


7/56


8/56

Ionic and Covalent Bonds: The Octet Rule1-3

The periodic table underlies the octet rule.

Electrons in atoms occupy levels orshells of fixed capacity.

The first has room for 2, the second 8, and the third 16.

Noble gaseshave 8valence electrons(Helium 2) and areparticularly stable.

Other elements lackoctetsin their outer electron shells and tendto form molecules in such a way as to create a stable octetarrangement.


9/56

In pure ionic bonds, electron octets are formed by transfer ofelectrons.

Alkali metalsreact withhalogens by the transfer of one electronfrom the alkali metal to the halogen.

Both atoms achieve a noble gas configuration: the alkali metalthat of the preceding inert gas, the halogen that of thefollowing inert gas.

IPNa= +119 kcal mol-1

EACl= -83 kcal mol-1

-LE = -120 kcal mol-1

E = -84 kcal mol-1


10/56

Valence electrons are conveniently indicated by placing dots around thesymbol for an element. The letters represent the nucleus and the coreelectrons, and the dots represent the valence electrons:

Hydrogen can either lose an electron to form an H+ion, or gain anelectron to form a H-, or hydride, ion:


11/56

Incovalent bonds, electrons are shared to achieveoctetconfigurations

Ionic bonds between identical atoms of the same element donot form.

The highionization potentialof hydrogen prevents it fromforming ionic bonds with halogens and other non-metallicelements.

Ionic bonds are also unfeasible for carbon since it would

require the loss of 4 electrons to achieve the octet of thepreceding inert gas, or the gain of 4 electrons to achieve theoctet of the following inert gas.


12/56

In these and similar cases,covalent bondingoccurs. Atomsshare electrons to achieve anoble gas configuration.

In certain cases, one atoms supplies both of the electrons in the

bond:

Often 4 electron (double) and 6 electron (triple) bonds are formed:


13/56

In most organic bonds, the electrons are not shared equally:polar covalent bonds.

Pure covalentbonds (perfect sharing of electrons) andionic

bonds(complete transfer of electrons) are two extreme types ofbonding.

Most bonds lie somewhere between these extremes and arecalledpolar covalentbonds.

Each element can be assigned anelectronegativityvalue whichrepresents its electron accepting ability when participating in achemical bond.

The larger the difference in electronegativety between twoatoms participating in a chemical bond, the more ionic is thebond.

Bonds between atoms of different electronegativity are said tobepolar bonds. A partial negative charge is found on the atomof higher electronegativity and an equal but positive charge onthe other atom.


14/56

As a rule of thumb, electronegativity differences less than 0.3represent pure covalent bonds, from 0.3 to 2.0 polar covalentbonds, and greater than 2.0 ionic bonds.

The separation of opposite charges in polar covalent moleculesresults in the formation ofdipoles:


15/56

In symmetrical molecules such as CO2and CCl4, the individualdipoles will cancel and the molecule is left with a zero dipolemoment.


16/56

Electron repulsion controls the shapes of molecules.

The shapes of molecules can be predicted using the VSEPRmethod.

Bonding and non-bonding electron pairs on the same atomwill arrange themselves in three-dimensions to be as far apartas possible.

In the case of 2 electron pairs, as in BeCl2, a lineararrangement results. For 3 electrons pairs, as is in BCl

3, a

trigonal arrangement results, and in the case of 4 electronpairs, a tetrahedral arrangement occurs:


17/56


18/56

It is often necessary to use double or triple bonds to satisfythe octet rule:


19/56

4. Assign charges to atoms in the molecule.

Charge = (# valence electrons in free, neutral atom)

- (# unshared electrons on the atom)

(# bonding electrons surrounding the atom)

In molecules such as nitric acid, charges occur on individualatoms, even though the molecule itself is neutral.


20/56

The octet rule does not always hold.

1. The molecule or ion has anodd numberof electrons.

NO, CH3, NO2

2. The central atom has a deficiencyof electrons.

CH3, BeCl2, BH3

3. Past row 2 of the periodic table, the central atom may besurrounded by more than 8 electrons (expanded octet).


21/56

Covalent bonds can be depicted by straight lines.

Bonding pairsof electrons are most often represented as straight

lines: single bonds as a single line, double bonds as two parallellines, and triple bonds as three parallel lines.

Lone pairsof electrons are either shown as dots or are omitted.

Structures of this type are calledKekulstructures.


22/56

Resonance Forms1-5

The carbonate ion has several correct Lewis structures.

Three equivalent structures must be drawn to accuratelyrepresent the carbonate ion. The only difference between thesestructures is the placement of electrons.


23/56

But what is its true structure?

The true structure can be thought of as the average of allthree structures which is called aresonance hybrid.

The 2 negative charges aredelocalizedover all three oxygenatoms.


24/56

Other examples of resonance:


25/56

Not all resonance forms are equal.

1. Structures with a maximum of octets are most important.

2. Charges should be preferentially located on atoms with

compatible electronegativity. If this conflicts with rule 1, thenrule 1 takes precedence.

3. Structures with less separation of opposite charges are moreimportant resonance contributors than those with morecharge separation.


26/56

In some cases charge separation is necessary and guideline 1takes precedence over guideline 2:

If there are two or more charge separated resonance structureswhich comply with the octet rule, the most favorable one placesthe charges on atoms of compatible electronegativity:


27/56

Atomic Orbitals: A Quantum Mechanical

Description of Electrons around the Nucleus

1-6

The electron is described by wave equations.

An electron within an atom can have only certain definiteenergies calledenergy states.

Moving particles such as electrons exhibit a wavelengthdetermined by thede Broglie relation:

h=

mv

Where h is Planks constant, m is the massof the electron in kg, and v is the velocity ofthe electron in m/s.


28/56


29/56

The wave theory of electron motion is calledquantummechanics.

The quantum mechanical equations describing the motion ofthe electrons are calledwave equations. The solutions of these

equations are calledwave functionsand are represented by theGreek letter,.

The square of the wave function, evaluated at a point in space(x,y,z) represents theprobabilityof finding the electron at thatpoint at any given time.

Each wave function corresponds to a specific discrete energyand the system is said to bequantized.


30/56


31/56


32/56

Following the 1s, 2s, and 2p orbitals are the 3s, 3p, 4s, 3d, etc.orbitals.

Organic chemistry deals primarily with the lower s and p orbitals.

Th A fb i i l i l bi l


33/56

The Aufbau principle assigns electrons to orbitals.

1. Lower energy orbitals are filled before those with higherenergy.

2. No orbital may be occupied by more than two electrons. (PauliExclusion Principle). If two electrons occupy a single orbital,they must have oppositespins. Electrons of opposite spins inthe same orbital are calledpaired electrons.

3. Degenerate orbitals must each receive a single electron of the

same spin before pairing of electrons occurs. (Hunds rule)


34/56


35/56

Atoms having a completely filled set of atomic orbitals are said tohave aclosed shell configuration. Atoms with a completely filledset are said to have anopen shell configuration.

The process of filling up the energy level diagram one electron at atime is called theAufbau process.

The d orbitalson atoms of row 3 and higher are involved in theformation ofexpanded octets(10 and 12 electrons about a centralatom).

Molec lar Orbitals and Co alent Bonding1 7


36/56

Molecular Orbitals: and Covalent Bonding1-7

The bond in the hydrogen molecule is formed by the overlapof 1s atomic orbitals.

Atomic orbitals on different atoms may overlap.The overlap of electron waves represented by the atomic orbitals mayresult in constructive (in phase) or destructive (out of phase)interference.

In phase overlap between two 1s orbitals results in a new orbital having

lower energy than either of the s orbitals. This new orbital concentratesthe electron probability between the two nuclei.

Out of phase overlap between two 1s orbitals results in a new orbitalhaving higher energy than either of the s orbitals. This new orbitalplaces most of the electron probability to the left and right of the two

nuclei.


37/56

An energy level diagramcan now be made of the two overlappingorbitals, and the Aufbau process used to determine the electronicconfigurations of H2and He2:


38/56

The overlap of atomic orbitals gives rise to sigma and pibonds.

When n atomic orbitals overlap, n newmolecular orbitalsare

formed.When n is 2, onebonding orbital and oneantibonding molecularorbital are formed.

The energy lowering of the bonding orbital and energy raising ofthe antibonding molecular orbital with respect to the atomic

orbitals is called theenergy splitting.The energy splitting indicates the strength of the bond formed.

Atomic orbitals of the same size and energy overlap to form thestrongest bonds.

Geometrical factors also affect the degree of overlap. Orbitalsexhibiting directionality in space (p orbitals) can overlap to formsigma () bonds or pi ()bonds.

All carbon-carbon single bonds contain one sigma bond. Doubleand triple bonds contain extra pi interactions.


39/56

Hybrid Orbitals: Bonding in Complex Molecules1 8


40/56

Hybrid Orbitals: Bonding in Complex Molecules1-8

Mixing of atomic orbitals from the same atom results in new atomicorbitals of different energy and directionality.

sp Hybrids produce linear structures.

An incorrect structure for BeH2is predicted if 2s and 2p orbitalsof Be are overlapped with the 1s orbitals of H:


41/56

Mixing the 2s orbital with one of the 2p orbitals of Be results intwo new hybrid sp orbitals, each made up of 50% s and 50% pcharacter. The resulting bond angle is 180owhich correspondswith the observed bond angle in the BeH2molecule.

Hybridization does not change the number of orbitals on theatom. In this case two atomic orbitals are replaced by two newhybrid orbitals. The two un-hybridized p orbitals are still

available to hold electrons.


42/56

sp2Hybrids create trigonal structures.

Hybridization of a 2s and two 2p orbitals results in three newhybrid orbitals that point to the corners of an equilateral triangle.

The remaining p orbital points up and down, perpendicular to eachof the three hybrid orbitals.

Bond angles in molecules using sp2hybridization are approximately120o

The molecule, BH3is isoelectric with the methyl cation, CH3+.

Both involve sp

2

hybridization about the central atom.


43/56

sp3Hybridizaton explains the shape of tetrahedral carboncompounds.

When the 2s and all three 2p orbitals are hybridized, four hybrid

orbitals called sp3

orbitals are formed. These orbitals point tothe corners of a regular tetrahedron.

Bond angles in molecules using sp3hybridization areapproximately 109.5o


44/56


45/56

Pi bonds are present in ethene (ethylene) and ethyne(acetylene).

Molecules containing double or triple bonds contain unhybridizedp orbitals that overlap lengthwise rather than end on.

Structures and Formulas of Organic Molecules1-9


46/56

Structures and Formulas of Organic Molecules1-9

To establish the identity of a molecule, we determine itsstructure.

The empirical formulaof a substance specifies the kinds and ratiosof elements present in the substance.

The empirical formula can be from anelemental analysisthesubstance.

More that one substance can have the same empirical formula.Each of these substances will have its own set of uniquephysicaland chemical properties, however.

Substances having the same empirical formula but differentconnectivity of atoms are calledconstitutional orstructural

isomers.


47/56

A chemist may be able to identify an unknown substance if itsproperties match those of a substance already determined.

New substances require other methods of identification such asx-ray crystallography, or various forms of spectroscopy.

Two ways of representing the structures of know molecules areball and stick models and space filling models.

Several types of drawings are used to represent


48/56

Several types of drawings are used to representmolecular structures.


49/56

Tetrahedral carbon structures can be accurately represented inthree dimensions using the dashed-wedged line notation.

The Big Picture1


50/56

The Big Picture1

1. The importance of Coulombs Law:

Atomic attraction

Relative electronegativity

Electron repulsion model for shapes of molecules

Choice of dominant resonance contributors.

2. The tendency of electrons to spread out (delocalize):

Resonance forms

Bonding overlap

3. The correlation of the valence electron count with theAufbau principle.

Associated stability of the elements in noble gas-octet-closed-shell configurations obtained by bond formation.

4. The characteristic shapes of atomic and molecular orbitals:

Provides a feeling for the location of the reacting electronsaround the nuclei.


51/56

Important Concepts1


52/56

Important Concepts1

1. Organic ChemistryChemistry of carbon and itscompounds.

2. Coulombs LawRelates attractive or repulsiveforce between charges to the distance between them.

3. Ionic BondsResult from coulombic attraction of

oppositely charged ions.4. Covalent BondsResult from electron sharing

between two atoms.

5. Bond LengthAverage distance between two

covalently bonded atoms6. Polar BondsFormed between atoms of differing

electronegativity

Important Concepts1


53/56

Important Concepts1

7. Shape of MoleculesStrongly Influenced byelectron repulsion.

8. Lewis StructuresDescribe bonding usingvalence electron dots. Hydrogen receives a duet whileother atoms receive an octet. Charge separationshould be minimized but may be enforced by the Octet

Rule.9. Resonance FormsWhen a structure is

described by two or more Lewis structures differingonly in their electron positions. The actual molecule isan average of the resonance forms. Some resonancestructures may be more important that others.

10. De Broglie RelationRelates wavelength of anelectron to its mass and velocity.

Important Concepts1


54/56

Important Concepts1

11. Wave EquationsDescribe motions of electronsabout the nucleus. Solutions are called orbitals.

These describe probabilities of finding the electrons in

particular regions of space.

12. s OrbitalSpherical. P-orbital Figure Eight. Eachorbital can hold two electrons of opposite spin. With

increasing energy, the number of nodes in an orbital

increases.

13. Aufbau PrincipleBuilding electronicconfigurations by adding one electron at a time to the

atomic orbitals, starting with those of lowest energy.

(Pauli exclusion principle, Hunds Rule).

Important Concepts1


55/56

Important Concepts1

14. Molecular OrbitalTwo overlapping atomicorbitals form either a bonding or an antibonding

molecular orbital. The number of molecular orbitals

equals the number of atomic orbitals overlapped.

15. BondsFormed when atomic orbitals overlapalong the bond axis. bonds Formed from p-orbitals

overlapping perpendicular to the bond axis.

16. Hybrid Orbitals- Formed by mixing of orbitals onthe same atom. sp: 2 linear orbitals, sp2: 3 trigonal

orbitals , sp3: 4 tetrahedral orbitals. Atomic orbitals not

hybridized remain unchanged. Hybrid orbitals cancontain either bonding or lone pair electrons.

Important Concepts1


56/56

Important Concepts1

17. Elemental AnalysisDetermines ratios of types

of atoms in a compound. Molecular Formula

Actual number of atoms of each type.

18. Constitutional Isomers(Structural Isomers)Same molecular formula but different connectivity of

atoms. Different properties.19. Condensed and Bond-Line Formulas

Abbreviated representations of molecules. Dashed-

Wedged Line DrawingsIllustrate molecules in three

dimensions.

Documents

Vollhardt 6e Lecture PowerPoints - Chapter 1