Essentials of organic chemistry
, Essentials of Organic Chemistry
Table of Contents
TABLE OF CONTENTS 2
STRUCTURE DETERMINES PROPERTIES 3
ALKANES AND CYCLOALKANES: AN INTRODUCTION TO HYDROCARBONS 7
CHIRALITY 11
ALCOHOLS AND ALKYL HALIDES: INTRODUCTION TO REACTION MECHANISMS 13
NUCLEOPHILIC SUBSTITUTION 17
STRUCTURE AND PREPARATION OF ALKENES: ELIMINATION REACTIONS 21
ADDITION REACTIONS OF ALKENES 26
CONJUGATION IN ALKADIENES AND ALLYLIC SYSTEMS 29
ARENES AND AROMATICITY 31
ALCOHOLS, DIOLS, THIOLS 34
ETHERS, EPOXIDES, AND SULPHIDES 36
ALDEHYDES AND KETONES 40
CARBOXYLIC ACID 45
CARBOXYLIC ACID DERIVATIVES: NUCLEOPHILIC ACYL SUBSTITUTION 48
AMINES 54
CARBOHYDRATES 55
AMINO ACIDS, PEPTIDES, AND PROTEINS 60
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, Essentials of Organic Chemistry
Structure determines properties
Atoms
Orbitals are linked to the position of the periodic table.
➔ First row: 1s (H & He).
➔ Second row: 2s, 2px, 2py, 2pz.
The second row has to follow the octet rule.
An orbital tells you the probability of finding an electron
somewhere within an atom.
s atomic orbitals
• The s atomic orbitals begin at n=1 and are spherically
shaped.
• There is a single orbital within each ns subshell.
• Within a particular shell, s-orbitals are lower in energy than other orbitals.
• As n increases, additional sign changes (nodes) appear in the shape of the orbital (s2, s3).
p atomic orbitals
• The p-atomic orbitals begin at n=2 and are dumbbell shaped.
• Three are 3 orbitals within each np subshell, which correspond to the three Cartesian
directions.
• The np orbitals all have the same energy; according to Hund’s rule electrons are left unpaired
when filling until they must be paired.
▪ They are a little bit higher energy than s-orbitals.
• All p-orbitals contain a node at the nucleus.
Bonding
• One can imagine ionic bonds as arising from the transfer of an electron from a metal to a
non-metal.
• A neutral hydrogen atom needs one more valence electron to achieve a full valence shell.
▪ Two hydrogens sharing their valence electrons achieve a full n=1 shell.
▪ In H2, the electron configuration of each hydrogen atom is analogous to that of
helium, the first noble gas.
▪ H:H / H-H → two dots or a line between two atoms denotes the sharing of two
electrons between the atoms in a covalent bond.
• A neutral fluorine atom needs one more valence electron to achieve a full valence shell.
▪ Two fluorine’s sharing their valence electrons achieve a full n=2 shell containing 8
electrons (= octet).
▪ In F2, the electron configuration of each hydrogen atom is analogous to that of neon,
the second noble gas.
• The octet rule = in stable molecules, second row atoms share electrons until they achieve an
octet.
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, Essentials of Organic Chemistry
• In Lewis structures, electrons are represented as dots or lines. A line denotes a pair of
electrons shared in a covalent bond.
• Two atoms of different electronegativities share electrons unequally in a covalent bond →
polar covalent bond.
▪ The greater the difference in electronegativity,
the more polarized the bond.
▪ High electronegativity → stronger pull on the
share electrons → stronger unequal sharing.
→ due to the stronger electronegativity of F, the electron will be closer to F which gives it the partial
negative charge.
→ polar covalent bonds tend to be sites of reactivity in organic molecules.
Formal charge
• Formal charge = the charge assigned to an atom in a molecule, assuming that electrons in all
chemical bonds are shared equally between atoms, regardless of relative electronegativity.
• Formal charge = FEC – VEC | FC = VEC – (bonds + dots)
▪ VEC = the valence electron count of the
neutral atom.
▪ FEC = formal valence electron count of the
covalently bound atom.
• Formal charges with magnitude greater than 1 are
not encountered in organic molecules.
Levels of organic chemistry
• The composition of a molecule is captured by its molecular formula.
• The constitution of a molecule is captured by a structural formula how atoms are connected.
• The configuration of a molecule is captured by a structural drawing showing the positions of
atoms in space.
• Compounds can have identical structures on a broad level but differ on a more specific level;
such compounds are called isomers (same composition, different
constitution/configuration).
Resonance
• Resonance = the movement of an electron within a structure; delocalization and stabilization
of charge. No loss/gain of electron.
• The structure doesn’t convert into another structure, it’s the sum of the two structures.
• Electrons will always move from negative side to the positive side.
• The connectivity and positions of the atoms must remain the same in all resonance
structures → breakage of double bonds to single bonds or vice versa.
• Each contributing structure must have the same number of electrons and the same net
charge.
• Atoms of second-row elements must not violate the octet rule.
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, Essentials of Organic Chemistry
• Each contributing structures must have the same number of unpaired electrons.
▪ |O≡O|↔ ∙O=O∙ this is not a resonance because there are more unpaired electrons.
• Resonance structures containing more covalent bonds are typically more important.
• Resonance structures with minimal separation of opposite charges are most important.
• Resonance structures with negative charge on the most electronegative atom and positive
charge on the most electropositive atom are most important.
Recognizing resonance and drawing valid resonance structures are important skills because:
• Resonance indicates that molecules are stabilized due to the delocalization of charge.
• Resonance forms containing formal charges can reveal
hidden points or reactivity in organic molecules.
▪ Minor contributor K reveals that carbon in C=O can
act as a Lewis acid.
Use of arrows
• Curved arrows: to show the movement of electrons in chemical reactions.
• Full-headed arrow: implies movement of a pair of electrons; arrow starts at the source and
ends at the sink.
Acids and bases
• Acid = a species that donates a proton to a base.
• Base = a species that accepts a proton from an acid.
• Species related by the addition or removal of a single proton are called conjugates.
▪ When an acid surrenders a proton in an acid-base reaction, its conjugate base is
formed.
▪ When a base gains a proton in an acid-base reaction, its conjugate acid is formed.
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, Essentials of Organic Chemistry
[𝑯𝟑 𝑶+ ]𝒆𝒒 [𝑨− ]𝒆𝒒
𝑲𝒂 =
[𝑯𝑨]𝒆𝒒
• H2O reacts with an acid HA to from H3O and A-.
+
• Acid dissociation constant (Ka) = the equilibrium constant.
▪ The larger Ka is, the stronger the acid.
▪ Ka for an acid is related to Kb for the conjugate base of the acid.
⬧ KaKb = Kw = 1.0x10-14
• pKa: the logarithm of Ka; pKa = -log(Ka).
▪ Thanks to the negative sign, a smaller pKa is associated with a stronger acid.
There are four structural factors that exert a profound influence on the acidity of a proton:
1. The strength of the bond to the atom from which the proton is lost.
▪ As the size of the atom to which H is bonded increases, the H-A bond strength
decreases and acidity increases.
▪ Greater polarizability of larger conjugate base anions also contributes to greater
acidity.
2. The electronegativity of the atom from which the proton is lost.
▪ As the electronegativity of A increases, the polarization of the H-A bond increases
and acidity increases.
▪ The partial positive charge on H increases, increasing the likelihood that it will be
transferred as H+ (i.e., HA will act as an acid).
▪ The covalent hybrids of second-row non-metals illustrate this effect nicely. Acidity
increases moving to the right in the periodic table.
3. Inductive effects of nearby atoms.
▪ Nearby electronegative atoms can have a similar effect on the acidity of a proton. As
the number of nearby electronegative atoms increases, acidity increases.
▪ Electronegative atoms pull electron density towards
themselves, increasing the partial positive charge on H
→ this pulling is inductive effect/induction.
▪ Trifluoroethanol is a considerably stronger acid than
ethanol due to inductive effects of the fluorine atoms.
4. Electron delocalization in the conjugate base.
▪ Loss of a proton establishes a lone pair on the conjugate base.
▪ When this lone pair is involved in resonance delocalization, the conjugate base is
stabilized.
▪ The acid is more acidic when the conjugate base is resonance stabilized.
▪ For example, acetic acid (CH3CO2H, pKa 4,75) is much more acidic than ethanol
(CH3CH2OH), pKa 16).
▪ The conjugate base of (CH3CO2-) is stabilized by resonance by
the conjugate base of ethanol (CH3CH2OH) is not.
Lewis theory of acids and bases
• Acid: a species that accepts an electron pair from a base.
• Base: a species that donates an electron pair to an acid.
• Bronsted acids and bases are subsets of Lewis acids and bases
respectively.
• Diethyl ether donates a pair of electrons from O and thus serves
as a Lewis base or nucleophile.
• Boron trifluoride accepts the pair of e- from diethyl ether at boron and thus serves as a Lewis
acid or electrophile.
• A new bond forms between O and B.
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, Essentials of Organic Chemistry
Alkanes and cycloalkanes: an introduction to hydrocarbons
Hydrocarbons are compounds containing only carbon and hydrogen.
They can be classified according to their sources:
• Aliphatic hydrocarbons: are derived from fats and oils and contain mostly carbon-carbon
single bonds.
• Aromatic hydrocarbons: are derived from essential oils and contain single and double
carbon-carbon bonds.
Hydrocarbons can also be classified according to their structures:
• Alkanes: contain only C-C bonds and C-H bonds.
• Alkenes: contain at least one C=C double bond.
• Alkynes: contain at least one C≡C triple bond.
• Arenes: contain a ring of alternating C-C and C=C bonds.
Bonding
The single bond in H2 and other organic molecules is a sigma bond with cylindrical symmetry about
the bonding axis.
• Sigma bond → overlap of 2 s-orbitals → symmetry.
• The electrons like the overlap of the orbitals.
Multiple bonds derived from side-on overlap of 2p-orbitals are called pi
bonds, asymmetric.
• Pi bonds contain a node along the axis connecting the bonding
nuclei.
Alkanes have the general formula CnH2n+2 and consist of C-C and C-H bonds. The simplest alkanes are
methane, ethane, and propane. The boiling point increases with molecular weight. All three
molecules have tetrahedral geometry at their carbon atoms.
Valence bond theory suggests that bonds are derived from
overlap of atomic orbitals. How can tetrahedral geometry be
explained by overlap of s-orbitals and p-orbitals at 90 degree
to one another.
We make use of a mathematical device that explains the
observed geometry
Hybridization = atomic orbitals on the same atom are mixed to produce a new set of hybrid atomic
orbitals
Hyberdization of atomic orbitals
The organic chemistry tutor
s1p1 = s+p orbitals 50% s-character, 50% p-character
s1p2 = s+p+p orbitals 33% s-character, 67% p-character
s1p3 = s+p+p+p orbitals 25% s-character, 75% p-character
• These indicate the location of an electron around the atom.
• C has four valence electrons → CH4 has four groups and is thus sp3 hybrid.
• Ethane H2C=CH2 → C has three groups, therefore sp2 hybrid.
• Triple bond is stronger: more s-character thus stronger and shorter.
• Degenerate orbitals: all orbitals have the same energy
• Unhybridized orbitals are orbitals that aren’t used, they’re used to form pi bonds.
• Sigma bonds are formed from the overlap of hybridized orbitals.
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, Essentials of Organic Chemistry
• The pi bonds exist above the sigma bonds.
• If it’s part of a resonant structure, you don’t count the lone pairs into the groups of the
orbitals → delocalized electrons.
• Localized electrons don’t move, they’re stuck to the atom and therefore are counted with
the groups of the orbitals.
Naming organic molecules
• Condensed formulas = textual and list atoms in order of their connectivity.
• Bond-line formulas omit labels for carbon atoms and omit hydrogens where they are
implied.
• Constitutional isomers have the same composition (molecular formula) but differ in how
their atoms are connected.
▪ Different physical and chemical properties.
▪ Methane, ethane and propane don’t have constitutional isomers. Butane has two.
⬧ In both isomers, all carbons are sp3 hybridized and tetrahedral.
▪ Pentane has three isomers, hexane has five.
• There is no general way to calculate the number of isomers associated with a molecular
formula.
• In simple cases, we can enumerate possible isomers by systematically moving branches
around and changing their composition.
• The IUPAC nomenclature system is a systematic method for naming alkanes.
▪ Linear alkane chains are named using a prefix for the number of carbons and the
suffix –ane.
• Naming branched alkanes:
▪ Identify the longest continuous carbon chain.
▪ Identify the substituents, branches attached to the parent chain.
▪ Number the carbons of the longest chain such that the substituent at the first
branching point gets the smallest number possible.
▪ Write the name of the compound.
⬧ List substituents in alphabetical order, preceded by their number location.
⬧ Write the parent chain as an alkane at the end of the name.
▪ Alkane substituents are called alkyl groups; methyl, ethyl, propyl, butyl etc.
• 2,2,6,6,7-pentamethyloctane and not 2,3,3,7,7-
pentamethyl because 2,2,6,6,7 is a lower number
than 2,3,3,7,7.
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, Essentials of Organic Chemistry
• Cycloalkanes are cyclic alkanes with the general
formula CnH2n, where n>2.
▪ Cycloalkanes are numbered so that the
substituent at the first point of difference
gets the lowest number.
▪ Simple cycloalkanes bonded to longer alkyl
chains are treated as cycloalkyl
substituents → the cyclobutane ring is
smaller than the attached pentane chain.
• Functional groups are structural units containing heteroatoms or multiple bonds.
▪ Alkyl halides contain an alkyl group bound to a halogen atom.
▪ Alcohols contain an alkyl group bound to a hydroxyl group.
Conformation
• The reactivity and properties of a compound depend to
some extent on the shape of its molecules.
• Conformation = the shape of a molecule with respect to
bond rotations and other “easy” processes.
▪ Staggered: H-atoms far away as possible, most stable.
▪ Eclipsed: H-atoms as close as possible, least stable
▪ Torsion angle: angle between two bonds in a conformation.
• Butane has two distinct staggered
conformations, unlike ethane.
▪ The terminal CH3 groups can be
gauche or anti.
▪ Steric interaction between the
gauche methyl groups is called a
gauche interaction.
⬧ Steric interaction: nonbonding interactions that influence the shape
(conformation) and reactivity of ions and molecules.
▪ In the anti-conformer, the CH3 groups do not engage in steric interactions with one
another.
• In general, the all-anti conformer is most stable.
▪ Viewed from the side, this conformer has a zig-zag appearance → it’s how it’s drawn.
• The most stable cycloalkanes should have bond angles near 109.5 degree.
• Cyclopropane is characterized by extremely severe angle strain which gives it its instability.
▪ The C-C bonds in cyclopropane are bent and derived from the angled overlap of sp3
hybrids.
• Cyclobutane is less strained than cyclopropane but still unstable.
▪ The cyclobutane ring puckers to minimize torsional or eclipsing strain.
• Cyclopentane has little angle strain, but considerable torsional strain due to eclipsing C-H
bonds.
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, Essentials of Organic Chemistry
▪ The envelope and half-chair conformers
involve less torsional strain and are of
similar energy → equilibration between
these two conformations is rapid.
• Cyclohexane adopts a chair conformation to
minimize torsional strain.
▪ In the chair conformation, all adjacent C-
H bonds are staggered.
▪ The hydrogens in a cyclohexane chair
can be divided in two groups:
▪ Axial: half are pointing up or down.
▪ Equatorial: half are aligned with the
carbons (parallel).
▪ Cyclohexane interconverts between two chair conformers: equatorial bonds become
axial and axial bonds become equatorial.
• The two chairs of methylcyclohexane are not equivalent: in one the methylgroups is axial and
in the other it’s equatorial. At equilibrium, the concentration of the CH3-eq conformer is
much higher than that of the CH3-ax conformer.
▪ The axial conformer contains significant van der Waals strain between the axial CH3
and hydrogens 3 and 5 positions. The equatorial conformer lacks destabilizing 1,3-
diaxial interactions.
Cis-trans isomerism
• Cycloalkanes containing groups on two different carbons exist as cis and trans isomers →
stereoisomers = isomers with the same connectivity but different positions of groups in
space.
▪ Cis isomer: both groups at the same side of the ring.
▪ Trans isomer: contains the two groups on opposite sides of the ring.
• To determine which chair conformer of a disubstituted cyclohexane is more stable, the
numbers and sizes of equatorial substituents must be compared.
▪ trans-1,4dimethylcyclohexane: the diequatorial conformer is most stable
▪ cis-1,4-dimethylcyclohexane: both are equal in energy.
• In disubstituted cyclohexanes with two different substituents, the more stable chair is the
one with the larger group equatorial.
▪ tert-butylcyclohexanes are essentially locked in the conformer with the –C(CH3)3
group equatorial.
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