Organic Chemistry

Organic Chemistry Structure and Function Fifth Edition W. H. Freeman & Company New York

K. Peter C. Vollhardt Neil E. Schore

CHAPTER 1 Structure and Bonding in Organic Molecules

1-1 The Scope of Organic Chemistry: An Overview •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 relative positioning 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 double bonds. (Chapters 16 and 17)

1-1 The Scope of Organic Chemistry: An Overview •Amines – Nitrogen containing functional group (Chapter 21) •Tools For Identification – Spectroscopy (Chapters 10, 11, 14 and 20) •Carbohydrates and Amino Acids – Multiple Functional Groups (Chapters 24 and 26)

1-1 The Scope of Organic Chemistry: An Overview •Synthesis is the making of new molecules •Wöhler’s 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 then destroyed on the pathway between reactants and products.

1-2 Coulomb Forces: A Simplified View of Bonding Bonds are made by simultaneous coulombic attraction and electron exchange. When two atoms approach, the electrons of one are attracted by the protons of the other and vice-versa. (+ ) ch arg e  (  ) ch arg e A ttractin g F o rce = co n stan t  2

d istan ce

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

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

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Ionic and Covalent Bonds: The Octet Rule

1. 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 is formed.

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

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Ionic and Covalent Bonds: The Octet Rule

The 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 are particularly stable. Other elements lack octets in their outer electron shells and tend to form molecules in such a way as to create a stable octet arrangement.

In pure ionic bonds, electron octets are formed by transfer of electrons. Alkali metals react with halogens by the transfer of one electron from the alkali metal to the halogen. Both atoms achieve a noble gas configuration: the alkali metal that of the preceding inert gas, the halogen that of the following inert gas.

IPNa = +119 kcal mol-1 EACl =

-83 kcal mol-1

-LE = -120 kcal mol-1 E =

-84 kcal mol-1

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

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

In covalent bonds, electrons are shared to achieve octet configurations • Ionic bonds between identical atoms of the same element do not form. • The high ionization potential of hydrogen prevents it from forming ionic bonds with halogens and other non-metallic elements. • Ionic bonds are also unfeasible for carbon since it would require the loss of 4 electrons to achieve the octet of the preceding inert gas, or the gain of 4 electrons to achieve the octet of the following inert gas.

• In these and similar cases, covalent bonding occurs. Atoms share 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:

In most organic bonds, the electrons are not shared equally: polar covalent bonds. •Pure covalent bonds (perfect sharing of electrons) and ionic bonds (complete transfer of electrons) are two extreme types of bonding. •Most bonds lie somewhere between these extremes and are called polar covalent bonds. •Each element can be assigned an electronegativity value which represents its electron accepting ability when participating in a chemical bond.

•The larger the difference in electronegativety between two atoms participating in a chemical bond, the more ionic is the bond. •Bonds between atoms of different electronegativity are said to be polar bonds. A partial negative charge is found on the atom of higher electronegativity and an equal but positive charge on the other atom.

•As a rule of thumb, electronegativity differences less than 0.3 represent pure covalent bonds, from 0.3 to 2.0 polar covalent bonds, and greater than 2.0 ionic bonds. •The separation of opposite charges in polar covalent molecules results in the formation of dipoles:

•In symmetrical molecules such as CO2 and CCl4, the individual dipoles will cancel and the molecule is left with a zero dipole moment.

Electron repulsion controls the shapes of molecules. •The shapes of molecules can be predicted using the VSEPR method. •Bonding and non-bonding electron pairs on the same atom will arrange themselves in three-dimensions to be as far apart as possible. •In the case of 2 electron pairs, as in BeCl2, a linear arrangement results. For 3 electrons pairs, as is in BCl3, a trigonal arrangement results, and in the case of 4 electron pairs, a tetrahedral arrangement occurs:

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Electron-Dot Model of Bonding: Lewis Structures

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 as many atoms as possible a surrounding electron octet, except for H, which requires a duet. • Elements at the right of the periodic table (non-metals) may contain lone pairs of electrons. Correct Lewis Structure

Incorrect Lewis Structures

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

4. Assign charges to atoms in the molecule.

Charge = (# valence electrons in free, neutral atom) – - (# unshared electrons on the atom) – ½(# bonding electrons surrounding the atom)

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

The octet rule does not always hold. 1. The molecule or ion has an odd number of electrons. NO, CH3, NO2 2. The central atom has a deficiency of electrons. CH3, BeCl2, BH3

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

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 parallel lines, 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 Kekulé structures.

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Resonance Forms The carbonate ion has several correct Lewis structures. Three equivalent structures must be drawn to accurately represent the carbonate ion. The only difference between these structures is the placement of electrons.

But what is its true structure?

The “true” structure can be thought of as the average of all three structures which is called a resonance hybrid.

The 2 negative charges are delocalized over all three oxygen atoms.

Other examples of resonance:

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, then rule 1 takes precedence. 3. Structures with less separation of opposite charges are more important resonance contributors than those with more charge separation.

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

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

1-6 Atomic Orbitals: A Quantum Mechanical Description of Electrons around the Nucleus The electron is described by wave equations. An electron within an atom can have only certain definite energies called energy states. Moving particles such as electrons exhibit a wavelength determined by the de Broglie relation:

λ=

h mv

Where h is Plank’s constant, m is the mass of the electron in kg, and v is the velocity of the electron in m/s.

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

The wave theory of electron motion is called quantum mechanics. The quantum mechanical equations describing the motion of the electrons are called wave equations. The solutions of these equations are called wave functions and are represented by the Greek letter, .

The square of the wave function, evaluated at a point in space (x,y,z) represents the probability of finding the electron at that point at any given time. Each wave function corresponds to a specific discrete energy and the system is said to be quantized.

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 given time. 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 s orbitals. The lowest energy s orbital is the 1s orbital.

The next highest energy orbital is the 2s orbital. Note that it contains a spherical nodal surface.

Of slightly higher energy than the 2s orbital are 3 degenerate 2p orbitals. These orbitals are shaped like a figure 8 and point along the 3 Cartesian axes.

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.

The Aufbau principle assigns electrons to orbitals.

1. Lower energy orbitals are filled before those with higher energy. 2. No orbital may be occupied by more than two electrons. (Pauli Exclusion Principle). If two electrons occupy a single orbital, they must have opposite spins. Electrons of opposite spins in the same orbital are called paired electrons. 3. Degenerate orbitals must each receive a single electron of the same spin before pairing of electrons occurs. (Hund’s rule)

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

The process of filling up the energy level diagram one electron at a time is called the Aufbau process. The d orbitals on atoms of row 3 and higher are involved in the formation of expanded octets (10 and 12 electrons about a central atom).

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Molecular Orbitals: and Covalent Bonding

The bond in the hydrogen molecule is formed by the overlap of 1s atomic orbitals. Atomic orbitals on different atoms may overlap. The overlap of electron waves represented by the atomic orbitals may result 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 concentrates the electron probability between the two nuclei.

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

An energy level diagram can now be made of the two overlapping orbitals, and the Aufbau process used to determine the electronic configurations of H2 and He2:

The overlap of atomic orbitals gives rise to sigma and pi bonds. When n atomic orbitals overlap, n new molecular orbitals are formed. When n is 2, one bonding orbital and one antibonding molecular orbital 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 the strongest bonds. Geometrical factors also affect the degree of overlap. Orbitals exhibiting directionality in space (p orbitals) can overlap to form sigma () bonds or pi ()bonds. All carbon-carbon single bonds contain one sigma bond. Double and triple bonds contain extra pi interactions.

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Hybrid Orbitals: Bonding in Complex Molecules

Mixing of atomic orbitals from the same atom results in new atomic orbitals of different energy and directionality. sp Hybrids produce linear structures. An incorrect structure for BeH2 is predicted if 2s and 2p orbitals of Be are overlapped with the 1s orbitals of H:

Mixing the 2s orbital with one of the 2p orbitals of Be results in two new hybrid sp orbitals, each made up of 50% s and 50% p character. The resulting bond angle is 180o which corresponds with the observed bond angle in the BeH2 molecule.

Hybridization does not change the number of orbitals on the atom. In this case two atomic orbitals are replaced by two new hybrid orbitals. The two un-hybridized p orbitals are still available to hold electrons.

sp2 Hybrids create trigonal structures. Hybridization of a 2s and two 2p orbitals results in three new hybrid orbitals that point to the corners of an equilateral triangle. The remaining p orbital points up and down, perpendicular to each of the three hybrid orbitals.

Bond angles in molecules using sp2 hybridization are approximately 120o

The molecule, BH3 is isoelectric with the methyl cation, CH3+. Both involve sp2 hybridization about the central atom.

sp3 Hybridizaton explains the shape of tetrahedral carbon compounds. When the 2s and all three 2p orbitals are hybridized, four hybrid orbitals called sp3 orbitals are formed. These orbitals point to the corners of a regular tetrahedron.

Bond angles in molecules using sp3 hybridization are approximately 109.5o

Hybrid orbitals may contain lone electron pairs: ammonia and water. Not all hybrid orbitals participate in bond formation. Some may contain lone pairs of electrons.

The bond angles in ammonia are 107.3o and that in water is 104.5o, both close to 109.5o. The slightly smaller bond angles in ammonia and water are due to the slightly larger volume requirements for lone pair electrons, which forces the remaining bonding pair electrons closer together.

Pi bonds are present in ethene (ethylene) and ethyne (acetylene). Molecules containing double or triple bonds contain unhybridized p orbitals that overlap lengthwise rather than end on.

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Structures and Formulas of Organic Molecules To establish the identity of a molecule, we determine its structure. The empirical formula of a substance specifies the kinds and ratios of elements present in the substance. The empirical formula can be from an elemental analysis the substance. More that one substance can have the same empirical formula. Each of these substances will have its own set of unique physical and chemical properties, however. Substances having the same empirical formula but different connectivity of atoms are called constitutional or structural isomers.

A chemist may be able to identify an unknown substance if its properties match those of a substance already determined. New substances require other methods of identification such as x-ray crystallography, or various forms of spectroscopy. Two ways of representing the structures of know molecules are ball and stick models and space filling models.

Several types of drawings are used to represent molecular structures.

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

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The Big Picture 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 the Aufbau principle. • Associated stability of the elements in noble gas-octet-closedshell configurations obtained by bond formation. 4. The characteristic shapes of atomic and molecular orbitals: • Provides a feeling for the location of the “reacting” electrons around the nuclei.

5. The overlap model for bonding:

• Allows a judgment of energetics, directions and overall feasibility of reactions.

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Important Concepts 1. Organic Chemistry – Chemistry of carbon and its compounds.

2. Coulomb’s Law – Relates attractive or repulsive force between charges to the distance between them.

3. Ionic Bonds – Result from coulombic attraction of oppositely charged ions.

4. Covalent Bonds – Result from electron sharing 5.

between two atoms. Bond Length – Average distance between two covalently bonded atoms

6. Polar Bonds – Formed between atoms of differing electronegativity

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Important Concepts 7. Shape of Molecules – Strongly Influenced by electron repulsion.

8. Lewis Structures – Describe bonding using valence electron dots. Hydrogen receives a duet while other atoms receive an octet. Charge separation should be minimized but may be enforced by the Octet Rule.

9. Resonance Forms – When a structure is described by two or more Lewis structures differing only in their electron positions. The actual molecule is an average of the resonance forms. Some resonance structures may be more important that others.

10. De Broglie Relation – Relates wavelength of an electron to its mass and velocity.

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Important Concepts 11. Wave Equations – Describe motions of electrons about the nucleus. Solutions are called orbitals. These describe probabilities of finding the electrons in particular regions of space.

12. s Orbital – Spherical. P-orbital – Figure Eight. Each orbital can hold two electrons of opposite spin. With increasing energy, the number of nodes in an orbital increases.

13. Aufbau Principle – Building electronic configurations by adding one electron at a time to the atomic orbitals, starting with those of lowest energy. (Pauli exclusion principle, Hunds’ Rule).

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Important Concepts 14. Molecular Orbital – Two overlapping atomic orbitals form either a bonding or an antibonding molecular orbital. The number of molecular orbitals equals the number of atomic orbitals overlapped.

15.  Bonds – Formed when atomic orbitals overlap

along the bond axis.  bonds – Formed from p-orbitals overlapping perpendicular to the bond axis.

16. Hybrid Orbitals - Formed by mixing of orbitals on the same atom. sp: 2 linear orbitals, sp2: 3 trigonal orbitals , sp3: 4 tetrahedral orbitals. Atomic orbitals not hybridized remain unchanged. Hybrid orbitals can contain either bonding or lone pair electrons.

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Important Concepts 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. DashedWedged Line Drawings – Illustrate molecules in three dimensions.