organic chemistry summary - examville.com study aids

8/9/2019 Organic Chemistry Summary - Examville.com Study Aids

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Organic ChemistryStructure and Function

Fifth EditionW. H. Freeman & CompanyNew York

K. Peter C. VollhardtNeil E. Schore


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CHAPTER 1

Structure and Bonding in OrganicMolecules


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The Scope of Organic Chemistry: An Overview1-1

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

Alkane Reactions Alkane bond strengths and reactions.(Chapter 3)

Cyclic Alkanes New properties and changes in reactivity(Chapter 4)

Stereoisomerism Same 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 Ketones Carbonyl Compounds C=O doublebonds. (Chapters 16 and 17)


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

Carbohydrates and Amino Acids Multiple Functional Groups(Chapters 24 and 26)


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Synthesis is the making of new moleculesWhlers Synthesis of Urea:

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

Reactions are the vocabulary, and mechanisms are the grammar of organic chemistry

Reactants (Substrates) Starting compounds Products

Reaction Mechanism Underlying details of a reaction

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


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


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Ionic and Covalent Bonds: The Octet Rule1-31. Covalent Bonds are based on the sharing of electrons.

If the electrons are not shared equally, a polar covalent (partially

ionic) bond is formed, otherwise a pure covalent bond isformed.

2. Ionic Bonds are based on the transfer of one or more electronsfrom one atom to another. The resulting cation and anion areelectrostatically attracted to each other.


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Ionic and Covalent Bonds: The Octet Rule1-3The periodic table underlies the octet rule.

Electrons in atoms occupy levels or shells of fixed capacity.

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

Nobel gases have 8 valence electrons (Helium 2) and areparticularly stable.

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


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

Alkali metals react with halogens 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


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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:


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In covalent bonds , electrons are shared to achieve octetconfigurations

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

The high ionization potential of 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.


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In these and similar cases, covalent bonding occurs. Atomsshare electrons to achieve a noble 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:


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

Pure covalent bonds (perfect sharing of electrons) and ionicbonds (complete transfer of electrons) are two extreme types of bonding.

Most bonds lie somewhere between these extremes and arecalled polar covalent bonds.

Each element can be assigned an electronegativity value 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 tobe polar bonds . A partial negative charge is found on the atomof higher electronegativity and an equal but positive charge onthe other atom.


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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 of dipoles :


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In symmetrical molecules such as CO 2 and CCl 4 , the individualdipoles will cancel and the molecule is left with a zero dipolemoment.


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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 BeCl 2 , 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:


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Electron-Dot Model of Bonding: Lewis Structures1-4

Lewis structures are drawn by following simple rules.

1. Draw the molecular skeleton

2. Count the number of available valence electrons

Add one electron for each negative charge, if an anion .

Subtract one electron for each positive charge, if a cation .

3. Depict all covalent bonds by two shared electrons, giving asmany atoms as possible a surrounding electron octet, except forH, which requires a duet.

Elements at the right of the periodic table (non-metals) maycontain lone pairs of electrons.

Correct Lewis Structure Incorrect Lewis Structures


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


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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.


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The octet rule does not always hold.

1. The molecule or ion has an odd number of electrons.

NO, CH 3 , NO 22. The central atom has a deficiency of electrons.

CH3 , BeCl 2 , BH 3

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


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Covalent bonds can be depicted by straight lines.

Bonding pairs of 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 pairs of electrons are either shown as dots or are omitted.

Structures of this type are called Kekul structures.


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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.


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But what is its true structure?

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

The 2 negative charges are delocalized over all three oxygenatoms.


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Other examples of resonance:


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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:


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Atomic Orbitals: A Quantum MechanicalDescription of Electrons around the Nucleus

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The electron is described by wave equations.

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

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

h =

m v

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


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The electron waves contain nodes, where the amplitude of thewave changes sign, and can interact with each other, producingeither constructive or destructive interference:


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The wave theory of electron motion is called quantummechanics.

The quantum mechanical equations describing the motion of the electrons are called wave equations . The solutions of theseequations are called wave functions and are represented by theGreek letter, .

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

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


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Atomic orbitals have characteristic shapes.

In an artists rendition of an atomic orbital, a surface is drawn which

contains most of the probability of finding the electron at a giventime.

Nodes in the a function become points or planes of 0 probability of finding the electron.

Higher energy wave functions have more nodes than do lower

energy wave functions.

The simplest atomic orbitals are spherical in shape and are called sorbitals. The lowest energy s orbital is the 1s orbital.


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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.


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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 opposite spins . Electrons of opposite spins inthe same orbital are called paired electrons .

3. Degenerate orbitals must each receive a single electron of thesame spin before pairing of electrons occurs. ( Hunds rule )


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Atoms having a completely filled set of atomic orbitals are said tohave a closed shell configuration . Atoms with a completely filledset are said to have an open shell configuration .

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

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


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An energy level diagram can now be made of the two overlappingorbitals, and the Aufbau process used to determine the electronicconfigurations of H 2 and He 2:


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

When n atomic orbitals overlap, n new molecular orbitals are

formed.When n is 2, one bonding orbital and one antibonding molecularorbital are formed.

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

orbitals is called the energy 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.


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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 180 o which correspondswith the observed bond angle in the BeH 2 molecule.

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 stillavailable to hold electrons.


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sp 2 Hybrids 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 sp 2 hybridization are approximately120 o

The molecule, BH 3 is isoelectric with the methyl cation, CH 3+ .Both involve sp 2 hybridization about the central atom.


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sp 3 Hybridizaton explains the shape of tetrahedral carboncompounds.

When the 2s and all three 2p orbitals are hybridized, four hybridorbitals called sp 3 orbitals are formed. These orbitals point tothe corners of a regular tetrahedron.

Bond angles in molecules using sp 3 hybridization areapproximately 109.5 o


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Hybrid orbitals may contain lone electron pairs: ammoniaand water.

Not all hybrid orbitals participate in bond formation. Some maycontain lone pairs of electrons.

The bond angles in ammonia are 107.3 o and that in water is

104.5o

, both close to 109.5o

. The slightly smaller bond anglesin ammonia and water are due to the slightly larger volumerequirements for lone pair electrons, which forces the remainingbonding pair electrons closer together.


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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


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Structures and Formulas of Organic Molecules1-9

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

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

The empirical formula can be from an elemental analysis thesubstance.

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

Substances having the same empirical formula but differentconnectivity of atoms are called constitutional or structural

isomers .


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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.


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

The Big Picture1


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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.


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5. The overlap model for bonding:

Allows a judgment of energetics, directions and overallfeasibility of reactions.

Important Concepts1


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Important Concepts1

1. Organic Chemistry Chemistry of carbon and itscompounds.

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

3. Ionic Bonds Result from coulombic attraction ofoppositely charged ions.

4. Covalent Bonds Result from electron sharingbetween two atoms.

5. Bond Length Average distance between twocovalently bonded atoms

6. Polar Bonds Formed between atoms of differingelectronegativity

Important Concepts1


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Important Concepts1

7. Shape of Molecules Strongly Influenced byelectron repulsion.

8. Lewis Structures Describe 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 Forms When 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 Relation Relates wavelength of anelectron to its mass and velocity.

Important Concepts1


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Important Concepts1

11. Wave Equations Describe motions of electronsabout the nucleus. Solutions are called orbitals .These describe probabilities of finding the electrons inparticular regions of space.

12. s Orbital Spherical. P-orbital Figure Eight. Eachorbital can hold two electrons of opposite spin. Withincreasing energy, the number of nodes in an orbitalincreases.

13. Aufbau Principle Building electronicconfigurations by adding one electron at a time to theatomic orbitals, starting with those of lowest energy.(Pauli exclusion principle , Hunds Rule ).

Important Concepts1


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Important Concepts1

14. Molecular Orbital Two overlapping atomicorbitals form either a bonding or an antibondingmolecular orbital. The number of molecular orbitalsequals the number of atomic orbitals overlapped.

15. Bonds Formed when atomic orbitals overlapalong the bond axis. bonds Formed from p-orbitalsoverlapping perpendicular to the bond axis.

16. Hybrid Orbitals - Formed by mixing of orbitals onthe same atom. sp: 2 linear orbitals, sp 2: 3 trigonalorbitals , sp 3: 4 tetrahedral orbitals. Atomic orbitals nothybridized remain unchanged. Hybrid orbitals cancontain either bonding or lone pair electrons.

Important Concepts1


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Important Concepts1

17. Elemental Analysis Determines 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 Drawings Illustrate molecules in three

dimensions.

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