organic chemistry vollhardt/schore 6th ed. chapter 1 chapter 1

Organic ChemistryOrganic ChemistryVollhardt/SchoreVollhardt/Schore 66thth Ed.Ed.

Chapter 1Chapter 1

Version 1.1:Marc Anderson

Dept of Chemistry and BiochemistrySan Francisco State University

[[email protected]]

Warning: this is notWarning: this is nota a ““powerpointpowerpoint coursecourse”” !!!!

Chapter 1: Structure and Bonding Chapter 1: Structure and Bonding in Organic Moleculesin Organic Molecules

[1] introduction[1] introduction[2] coulomb forces[2] coulomb forces[3] ionic and covalent bonds[3] ionic and covalent bonds[4] electron dot model of bonding[4] electron dot model of bonding[5] resonance forms ([5] resonance forms (!!!!!!))[6] atomic [6] atomic orbitalsorbitals[7] molecular [7] molecular orbitalsorbitals ((~~) / covalent bonding) / covalent bonding[8] hybrid [8] hybrid orbitalsorbitals / hybridization (/ hybridization (!!!!!!))[9] structures, formulas, drawing molecules ([9] structures, formulas, drawing molecules (!!))

[1] Introduction[1] Introduction

What do I enjoy about organic chemistry?What do I enjoy about organic chemistry?It provides a core understanding for a number of

fascinating and highly relevant areas of study:• medicinal chemistry• materials: dyes, paints, fibers, polymers• biochemistry / pharmacology / neuroscience• fuel chemistry

[2] Coulomb Forces[2] Coulomb Forces

A simplified model of chemical bondingA simplified model of chemical 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.

This attraction is defined by Coulomb’s law:

2

(+) charge ( ) chargeAttracting Force = constant

distance

The attractive force releases a certain amount of energy called the bond strengthThe optimal distance between the two nuclei is called the bond length.

There is an optimal balance between these two parameters

ionic and covalent bondsionic and covalent bonds

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.

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.

[3] Ionic and Covalent Bonds: The Octet Rule[3] 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 shell has room for 2, the second 8, and the third 18.

Noble gases have 8 valence electrons (helium has 2) and are particularly stable.

Other elements lack octets in their outer electron shells and are driven to form molecules in such a way as to create a stable octet (“nobel gas”) arrangement

The periodic table underlies the octet rule.

In pure ionic bonds, electron octets can be formed by transferof electrons.

Alkali metals react with halogens: one electron is transferred and then both atoms achieve noble gas configuration

2,8,1 2,8,7 2,8 + 2,8,8

Formation of Ionic Bonds by Electron Transfer

Na + Cl [Na ] [Cl ] , or NaCl

Ionization Potential: Na = +119 kcal/mol (endothermic to lose e-)

Electron Affinity: Cl = -83 kcal/mol G (exothermic for Cl

to receive electron)

Electrostatic Attaction = -120 kcal/mol (exothermic attraction)

Eoverall = -84 kcal/mol (very exothermic!)

We can indicated the valence electron shell as “electron dots”

This is convenient when illustrating electron-transfer:

Note: hydrogen can gain or lose an electron to generate either “proton” or “hydride” (fundamentally different species!)

covalencycovalency

In covalent bonds, electrons are shared to achieve octet config.

Ionic bonds between identical atoms of the same element do not form. Instead, the electrons are shared and the bond is covalent!

Does Cl-Cl bond covalently or ionically?

Carbon does NOT form ionic bonds.Sharing requires loss or gain of 4 e- to achieve noble gas config

Carbon typically achieves neon configuration and bonds covalently

Ionic and Ionic and Covalent Bonds: Covalent Bonds:

ExamplesExamples

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

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

polar covalent bondspolar covalent bonds

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

Pure covalent bonds (perfect sharing of e-) and ionic bonds(complete transfer of e-) are two extremes.

Many bonds exist between extremes: they are polar covalent

Each element has an electronegativity value which represents its e- accepting ability. (How strongly it pulls electrons toward it)

Large difference in electronegativity between atoms will result in a highly polar covalent bond.

polar covalent bondspolar covalent bonds

The separation of charges in polar covalent molecules results in the formation of dipoles:

Symbolize this by an arrow pointing toward the more electronegative atom (with a “cross” at the other atom):

Note: in symmetrical molecules such as CO2, the individual dipoles cancel:

Btw: this picture is a computer-generated “electrostatic map”

Red indicates relative presence of electrons; blue indicates relative absence of electrons.

geometry of atoms: VSEPR theorygeometry of atoms: VSEPR theory

Electron repulsion controls the shape (“structure”) of molecules:

The shapes of molecules can be predicted using the VSEPR method.

Electron pairs arrange themselves to be as far apart as possibleIn the case of: 2 electron pairs: linear geometry

3 electrons pairs: trigonal geometry4 electron pairs: tetrahedral geometry

… and for organic molecules:

AgendaAgenda


Simple rules for Lewis Structures:[1] Draw the molecular skeleton (connect the atoms)[2] Depict all covalent bonds by a pair of shared electrons[3] Give each atom a nobel gas configuration (usually a full octet)[4] Electronegative elements may contain lone pairs of electrons

(N, O, Br, Cl).

[4] Electron[4] Electron--dot model of bonding: Lewis Structuresdot model of bonding: Lewis Structures

More examples:

formal chargesformal charges

Formal charges need to be determined for all atoms in a molecule:

FC = (# valence electrons in free, neutral atom) –- (# lone pair electrons on the atom) – ½(# covalent bonding electrons surrounding atom)

We’ll see a simpler formula in a second…

let’s try these on the board …

It is useful to memorize the valence electron counts ofB, C, N, and O. (and remember that S is under O).

three main exceptions to the octet rulethree main exceptions to the octet rule

[1] Atom has an odd # of electrons:

[2] Central atom has a free p orbital: +CH3, BeCl2, BH3

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

a simpler method of drawinga simpler method of drawingorganic molecules: organic molecules: KekulKekuléé structuresstructures

Covalent bonds can be depicted by straight lines.Bonding pairs of electrons: represented as straight lines

Lone pairs of electrons: indicated with dots (sometimes omitted)

Structures of this type are called Kekulé structures.

formal charges with formal charges with KekulKekuléé structures is simplifiedstructures is simplified

“FC = [# valence electrons] – dots – dashes”

(easier!)

Examples: you fill in the charge

Nitromethane(CH3NO2)DMSO

Some general rules: memorizing this will save you very much time, especially on the exams!!!!

“carbocation” “carbanion”

what is this molecule ???

[5] Resonance Hybrids (***)[5] Resonance Hybrids (***)

The carbonate ion has several correct Lewis structures.

Three equivalent structures must be drawn to accurately represent the carbonate ion.

The only difference: placement of electrons

Always draw arrows from a negative atom!This will beemphasized many many more times…

But what is the true structure of carbonate?

…it can be thought of as of as the average of all three structures which are each called resonance hybrids.

Each of the resonance structures contributes equally to the overall structure of this molecule.

The 2 negative charges are delocalized over all three oxygen atoms

The resonance structures areall present simultaneously; i.e. resonance is not “interconversion”

resonance hybrids: more examplesresonance hybrids: more examples

Note the direction of the arrows!

resonance hybrids: rulesresonance hybrids: rules

Not all resonance structures are equal in occurrence.

There are rules that identify prevalence of specific forms:

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

Rule 2: Charges should be located on atoms with compatible electronegativity.

Rule 3. Separation of charge should be minimized

resonance?

Example problem

Example problem

For each of the following, are they correct or incorrect? why ?

AgendaAgenda


[6] Atomic orbital theory[6] Atomic orbital theory

… is a quantum mechanical description of electrons around the nucleus …

The behavior of 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:

hλ=

mv

[ 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.]

This defines the wavelength of the electron.(which is strange: we’re not used to a particle having a wavelength)

interference in electron waves.interference in electron waves.

The electron waves contain nodes, where the amplitude of the wavechanges sign.

The waves can interact with each other, producing either constructive or destructive interference:

quantum mechanics overview (~)quantum mechanics overview (~)

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, . (psi)

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 atomic orbitalsorbitals have characteristic shapes !!have characteristic shapes !!

Knowing where electrons are tells us the geometry of the atom!!(and where bonds form).

This then tells us the possible shape of molecules!

Most importantly: quantum mechanics provides a surface, which indicates the probability of finding the electron at a given time.

The simplest example is an “s” orbital: (“spherical”)

The next highest energy orbital is the 2s orbital, which alsohas a spherical surface:

Of slightly higher energy than 2s are three degenerate 2p orbitals.

These are shaped like a“dumbell” and point along the three Cartesian axes.

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

Principles describing the asignment electrons to orbitals

[1] Lower energy orbitals are filled before those with higher energy.

(Aufbau principle).

[2] No orbital may be occupied by more than two electrons. (Pauli Exclusion Principle).

[3] Degenerate orbitals must each receive a single electron of the same spin before pairing of electrons occurs. (Hund’s rule)

Example: carbon

Examples: N, O, and F

(in their neutral forms: with 5, 6, or 7 valence electrons)

[7] Molecular [7] Molecular OrbitalsOrbitals and Covalent Bondingand Covalent Bonding

Atomic orbitals on two different atoms may overlap.

Overlap of two atomic orbitals leads to two molecular orbitals.(one bonding and one antibonding).

Example: H2 (the simplest possible molecule):

An energy level diagram can now be made of the two overlapping orbitals

Now we can use the Aufbau process to determine the electronic configurations of: H2 and He2:

AgendaAgenda


[8] Hybridized [8] Hybridized orbitalsorbitals: sp: sp33

Orbital hybridization allows us to explain more complexbonding arrangements that are typically observed

Example: sp3 hybridization explains the shape of tetrahedral carbon compounds.

One 2s and three 2p orbitals combine to make four sp3 orbitals:

hybridized hybridized orbitalsorbitals: sp: sp22

One 2s and two 2p orbitals combine to make three sp2

orbitals:

one p orbital is left unoccupied! (not clearly shown above)

One 2s and one 2p orbital combine to make two sp orbitals:(two p orbitals are left unoccupied)

hybridized hybridized orbitalsorbitals: sp: sp

water and ammonia are both spwater and ammonia are both sp33

Not all hybrid orbitals participate in bonding. Some 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

multiple bonds: multiple bonds: etheneethene and and ethyneethyne

Double and triple bonded molecules utilize two types oforbitals for bonding:

Single bond: sp2 or sp hybridized orbital.Double/triple bond: “extra” p orbital(s) (un-hybridized!)

[9] Structures and formulas in org chem[9] Structures and formulas in org chem

To establish the identity of a molecule, we determine its structure.

Empirical formula = what elements present, in what ratios. (determined via elemental analysis or mass spectrometry).

Substances having the same empirical formula but different connectivity of atoms are called constitutional or structural isomers.

[9] Structure identification: spectroscopy[9] Structure identification: spectroscopy

To fully characterize the chemical structure of an organic molecule we use techniques of spectroscopy (next semester).

There are two ways of representing 3D-structures: ball and stick and space-filling models:

Ethanol (bp= 78.5o C) Dimethyl Ether (bp= -23o C)CONSTITUTIONAL ISOMERS: both are C2H6O

1H NMR:

condensed and skeletal modelscondensed and skeletal models

Kekulé Condensed Skeletal

additional practice problems additional practice problems ……

Hint: check your answers using:

Some practice problems.1. Convert each skeletal structure to Kekule/line‐bond representation2. Determine the molecular formula

Hint: check your answers using:

Challenging practice problems.Identify examples of sp3, sp2, and sp hybridized atoms, of these complicatednatural products.

OH O

O

O

HNS

OH

OH

O

H

O

HN

O

N

O OO

O

NH

O

CH3

Latrunculol A Yanucamide A

organic chemistry vollhardt/schore 6th ed. chapter 1 chapter 1

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