Chapter%202%20Alkanes%20and%20Cycloalkanes:%20Introduction%20to%20Hydrocarbons

1
Chapter 2Alkanes and Cycloalkanes Introduction
to Hydrocarbons
2
Classes of Hydrocarbons
3
Classes of Hydrocarbons

• Hydrocarbons only contain carbon and hydrogen
  atoms.
• Hydrocarbons are either classed as aliphatic or
  aromatic.
• Aliphatic hydrocarbons contain three main
  groups alkanes which only have carbon-carbon
  single bonds, alkenes which have a carbon-carbon
  double bond, or alkynes which have a
  carbon-carbon triple bond.

4
Classes of Hydrocarbons

• Aromatic hydrocarbons are more complex but the
  simplest aromatic hydrocarbon is benzene.
  Aromatic hydrocarbons are called arenes.

5
Electron Waves andChemical Bonds
6
Models for Chemical Bonding
The Lewis model of chemical bonding predates the
idea that electrons have wave properties. Two
widely used theories of bonding based on the wave
nature of an electron are Valence Bond Theory,
and Molecular Orbital Theory
7
Formation of H2 from Two Hydrogen Atoms
e
e

• Which electrostatic forces are involved as two
  hydrogen
• atoms approach each other and form a H-H bond.
• These electrostatic forces are
• attractions between the electrons and the nuclei
• repulsions between the two nuclei
• repulsions between the two electrons

8
Potential Energy vs Distance Between Two
Hydrogen Atoms
weak net attraction atlong distances
Potentialenergy
H H
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Potential Energy vs Distance Between Two
Hydrogen Atoms
attractive forces increasefaster than repulsive
forcesas atoms approach each other
Potentialenergy
H H
H
H
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Potential Energy vs Distance Between Two
Hydrogen Atoms
maximum net attraction (minimum potential
energy)at 74 pm internuclear distance
74 pm
Potentialenergy
H H
H
H
-436 kJ/mol
H2
11
Potential Energy vs Distance Between Two
Hydrogen Atoms
repulsive forces increasefaster than attractive
forcesat distances closer than 74 pm
74 pm
Potentialenergy
H H
H
H
-436 kJ/mol
H2
12
Models for Chemical Bonding

• Valence Bond Theoryconstructive interference
  between two half-filled atomic
• orbitals is basis of shared-electron bond
• Molecular Orbital Theoryderive wave functions of
  moleculesby combining wave functions of atoms

13
Behavior of Waves

• Waves interactions includeConstructive
  interference when the waves are in phase and
  reinforce each otherDestructive interference
  when the waves are out of phase and oppose each
  other

14
Valence Bond Model for Bonding in Hydrogen

• Electron pair can be shared when half-filled
  orbital of one atom overlaps in phase with
  half-filled orbital of another. For example
  with overlap of two 1s orbitals of two hydrogen
  atoms shown below

15
Valence Bond Model

• The approach of the two hydrogen atoms can be
  modeled showing electrostatic potential maps.
  The high electrondensity between the nuclei is
  apparent.

Electrons feel the attractive force of the protons
Orbitals begin to overlap
Optimal distance between nuclei
High electron density between the nuclei
16
The Sigma (s) Bond
A bond in which the orbitals overlap along a line
connecting the atoms is called a sigma (s)
bond.Two perpendicular views are shown below.
17
Bonding in H2The Molecular Orbital Model

• Electrons in molecules occupy molecular orbitals
  (MOs) just as electrons in an atom occupy atomic
  orbitals (AOs).
• MOs are combinations of AOs.
• Two electrons per MO.
• The additive combination of two atomic orbitals
  generates one bonding orbital.
• The subtractive combination of the two atomic
  orbitalsgenerates an antibonding orbital.

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Molecular Orbital Model for H2

• Addition of the AOs to form the bonding MO (s)

Subtraction of the AOs to form the antibonding MO
(s)
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Molecular Orbital Digrams

• Format is AOs on the sides and MOs in the middle.
• Combination of n AOs results in n MOs.
• Bonding MOs lower in energy than antibonding MOs.
• Fill electrons in MOs the same as for AOs
  lowest first.

20
Energy-Level Diagram for H2 MOs
21
Introduction to AlkanesMethane, Ethane, and
Propane
22
Small Alkanes
General formula for alkanes is CnH2n2.
Smallest alkane is methane CH4 - also the most
abundant. Ethane (C2H6) and propane (C3H8) are
the next alkanes. Natural gas is 75 methane 10
ethane and 5 propane.These alkanes have the
lowest boiling points.
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Structures of Alkanes
All carbons in methane, ethane and propane have
four bonds. Bond angles (which are close to
109.5o) and bond lengths are
24
sp3 HybridizationandBonding in Methane
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Structure and Bonding Theory

• The dilemmaMethane has tetrahedral geometry.
• This is inconsistent with electron configuration
  of carbon of1s2, 2s2, 2px1,2py1 with only two
  unfilled orbitals.

26
sp3 Hybrid Orbitals
Linus Pauling proposed a mixing or hybridization
of the s and three p orbitals to create 4 equal
unfilled orbitals called sp3 orbitals.
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Properties of sp3 Hybrid Orbitals
All four sp3 orbitals are of equal energy. The
axes of the sp3 orbitals point toward the corners
of a tetrahedron. s Bonds involving sp3 hybrid
orbitals of carbon are stronger than those
involving unhybridized 2s or 2p orbitals.
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Bonding with sp3 Hybrid Orbitals
Bonding in methane involves orbital overlap
between each partially filled carbon sp3 orbital
and a partially filled s orbital of the hydrogen
atom.
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Bonding and Structure of Ethane

• Ethane also has tetrahedral geometry about the
  carbon atoms.
• Hybridization can be used to rationalize the
  bonding.
• The C-H bonds are formed as described for
  methane.
• The C-C bond is formed by overlap of sp3 orbitals
  on each of the carbon atoms.

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C-C Bond Formation in Ethane
Two half-filled sp3 orbitals on each C
Electrons with opposite spin
Overlap of orbitals to form a bondingorbital.
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Structure of Ethylene and sp2 Hybridization

• Ethylene is planar with bond angles close to
  120o.
• sp3 Hybridization cannot be used to explain this
  bonding.
• Three atoms are bonded to each carbon so three
  hybridorbitals are formed. Called sp2 orbitals.
• One p orbital is not hybridized.

32
sp2 Hybrid Orbitals

• The 2s and two of the 2p orbitals are mixed to
  form three sp2 orbitals with a trigonal planar
  arrangement.The 2pz orbital remains half
  filled.

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Sigma (s) Bonding in Ethylene
Form C-H bonds by overlap of sp2and s orbitals
Form C-C bond by overlap of sp2orbitals on each
carbon
These are all sigma (s) bonds. An unfilled p
orbital remains on each carbon atom.
34
Pi (p) Bonding in Ethylene
Form second C-C bond by overlap of p orbitals on
each carbon
This called a pi (p) bond and the electrons in
the bond are called p electrons.
35
Structure of Acetylene and sp Hybridization

• Acetylene is linear with bond angles of 180o.
• sp3 and sp2 Hybridization cannot explain this
  bonding.
• sp Hybridization explains this. There are two
  half filled p orbitals no hybridized.

36
sp Hybrid Orbitals

• The 2s and one of the 2p orbitals are mixed to
  form two sp orbitals with a linear arrangement.
  The 2py and 2pz orbitals remain half filled.

37
Sigma (s) Bonding in Acetylene
Form C-H bonds by overlap of spand s orbitals
Form C-C bond by overlap of sporbitals on each
carbon
These are all sigma (s) bonds. Two unfilled p
orbitals remain on each carbon atom.
38
Pi (p) Bonding in Acetylene
Form one p bond by overlap of py orbitals on each
carbon
Form second p bond by overlap of pz orbitals on
each carbon
There are two pi (p) bonds and a total of 4 p
electrons.
39
Hybridization of Carbon

• Carbons bonded to four atoms are sp3 hybridized
  with bond angles of approximately 109.5o.
• Carbons bonded to three atoms are sp2 hybridized
  withbond angles of approximately 1200 and one
  C-C p-bond.
• Carbons bonded to two atoms are sp hybridized
  withbond angles of approximately 1800 and two
  C-C p-bonds.

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Which Theory of Chemical Bonding Is Best?
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Theories of Chemical Bonding

• Approaches to chemical bonding
• Lewis model
• Orbital hybridization model
• Molecular orbital model.

42
Considerations of Chemical Bonding

• Lewis and Orbital hybridization models work
  together and success in organic depends on
  writing correct Lewis structures.
• Molecular orbital theory provides insights into
  structureand reactivity lacking in the other
  models. This model requires higher level theory
  which will not be presented.The results of MO
  theory will be used for example electrostatic
  potential maps.

43
Isomers of Butane

• There is only one isomer for each of the
  molecular formulas CH4, C2H6 and C3H8.
• For C4H10 there are two distinct connectivities
  of the carbon atoms. They are constitutional
  isomers.

Bondlineformulas
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Isomers of Butane

• The isomers have different physical properties.
• All carbon atoms are sp3 hybridized.

45
Higher n-Alkanes

• n-Alkanes are straight-chain alkanes with general
  formula CH3(CH2)nCH3. n-Pentane is
  CH3CH2CH2CH2CH3 and n-hexane is
  CH3CH2CH2CH2CH2CH3. These formulas can be
  abbreviated as CH3(CH2)3CH3 or CH3(CH2)4CH3.

46
Isomers of C5H12

• There are three isomers C5H12.

It is important to realize that these are all
representations of isopentane.
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Isomers of higher n-alkanes

• For higher n-alkanes there are many isomers and
  it is not possible to easily predict how many
  isomers can be formed.

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IUPAC Nomenclature ofUnbranched Alkanes
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IUPAC Naming

• Alkane names are the basis of the IUPAC system of
  nomenclature. The ane suffix is specific to
  alkanes.

50
The IUPAC Rules for Branched Alkanes

• Rules for naming branched alkanes
• Find the longest continuous carbon chain and its
  IUPAC name. This is the parent alkane.
• Identify the substituents on this chain.

substituent
longest chain (5 carbons)
51
The IUPAC Rules for Branched Alkanes

• Rules for naming branched alkanes
• 3. Number the longest continuous chain in the
  direction that gives the lowest number to the
  first substituent.
• 4. Write the name of the compound. The parent
  alkane is the last part of the name and is
  preceded by the names of the substituents and
  their numerical locations (locants). Hyphens
  separate the locants from the words.

2-methylpentane
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The IUPAC Rules for Branched Alkanes

• Rules for naming branched alkanes
• When the same substituent appears more than once,
  
• use the multiplying prefixes di-, tri-, tetra-,
  and so on. A separate locant for each
  substituent. Locants are separated from each
  other by commas and from the words by hyphens.

2,3-dimethylbutane
2,2-dimethylbutane
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Alkyl Groups
Alkyl groups are substituents derived from
alkanes.They lack one hydrogen at the point of
attachment. The alkyl group is named from the
alkane by replacing the-ane suffix with yl.
For example a CH3CH2CH2CH2- substituent is a
butyl group.
54
Classification of Carbon Atoms
Carbon atoms are defined as primary, secondary,
tertiary or quaternary.A primary carbon is
directly attached to one other carbon.A
secondary carbon is directly attached to two
other carbons. A tertiary carbon to 3 and a
quaternary carbon to 4.
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Complex Alkyl Groups (Substituents)
Secondary and tertiary groups may have common
names and IUPAC names.The base name of these
groups is the longest chain including the
attachment carbon form and the substituents are
located on this chain.
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Naming Highly Branched Alkanes
When two or more different substituents are
present number from the end closest to the first
point of difference. When two or more different
substituents are present, they are listed in
alphabetical order in the name. Prefixes such as
di-, tri-, and tetra- are used but ignored when
alphabetizing.tert-Butyl precedes isobutyl.
sec-Butyl precedes tert-butyl.
4-ethyl-3,5-dimethyloctane
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Naming Highly Branched Alkanes
When two or more different substituents are
present number from the end closest to the first
point of difference. If the first substituent is
located an equal distance from eachend then the
second substituent becomes the first
potentialpoint of difference and so on.
58
Naming Cycloalkanes
Cycloalkanes contain a ring of carbons and have
general formula CnH2n. Add the prefix cyclo- to
the name of the corresponding alkane.
59
Naming Cycloalkanes
Identify and name substituents as before. For
one substituent no numbers are used.
60
Naming Cycloalkanes
For multiple substituents the locations must be
specified. Number the carbon atoms of the ring
in the direction that gives the lowest number to
the substituents at the first point of
difference. First substituent is on C1 by
default.
61
Naming Cycloalkanes
If the ring has fewer carbons than the alkyl
group attached to it then the ring is the
substituent.
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Sources of Alkanes and Cycloalkanes
Natural is mainly methane with ethane and
propane. Petroleum is a liquid mixture
containing approximately 150 hydrocarbons. Half
of these are alkanes or cycloalkanes. Distillatio
n of crude oil gives fractions based on boiling
point.
63
Petroleum Refining
The yield of the more useful petroleum fraction
used as automotive fuel is increased by two
processes Cracking. Cracking is the cleavage of
carboncarbon bonds in high molecular weight
alkanes induced by heat (thermal cracking) or
with catalysts (catalytic cracking). Reforming.
Reforming converts the hydrocarbons in petroleum
to aromatic hydrocarbons and highly branched
alkanes, both of which are better automotive
fuels than unbranched alkanes and cycloalkanes.
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Other Natural Sources of Alkanes
Solid n-alkanes are waxy and coat the outer
surface of many living things to prevent loss of
water. Examples includePentacosane
(CH3(CH2)23CH3 is found in the waxy outer layer
of many insects. Hentriacontane is a component of
beeswax and the outer layer of leaves of tobacco,
peach trees and others. Hopanes are found in
petroleumand geologic sediments.
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Physical Properties of Alkanesand Cycloalkanes
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Boiling Point
Boiling points of n-alkanes increase with
increasing molecular weight (number of
carbons).Branched alkanes generally have lower
boiling points thanunbranched alkanes with the
same number of carbons.
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Intermolecular Forces and Boiling Point
Attractive forces between molecules in the liquid
phase affect the boiling point of the
liquid.These Intermolecular forces are van der
Waals forces and may be divided into three
typesDipole-dipole (including hydrogen
bonding)Induced dipole-dipole orInduced
dipole-induced dipole.
68
Intermolecular Forces and Alkanes
Alkanes have no dipole so the van der Waals
forces are the temporary induced dipole-induced
dipole.This interaction is dynamic and
fluctuates.
69
Intermolecular Forces and Alkanes
Long chain alkanes have more induced
dipole-induced dipole interactions so the boiling
point increases with increasing chain length.
70
Intermolecular Forces and Alkanes
Branched alkanes have lower surface area than
isomeric n-alkanes and therefore have lower
boiling points.
71
Melting Point
Solid n-alkanes are soft low melting solids. The
same intermolecular forces hold the molecules
together in the solid state.
72
Solubility of Alkanes in Water
Alkanes (and all hydrocarbons) are virtually
insoluble in water and are said to be
hydrophobic. The densities of most alkanes are
in the range 0.6-0.8 g/mL therefore alkanes float
on the surface of water.
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Chemical Properties of Alkanesand Cycloalkanes
74
Acidity of Hydrocarbons
Hydrocarbons are very weak acids. Alkynes have
the lowest pKa.
75
Combustion of Hydrocarbons
Combustion of hydrocarbons is exothermic
generating CO2 and water.
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Combustion of Relative Stability
All isomers of C8H18 generate 8 molecules of CO2
and 9 of H2O yet different amounts of energy.
This energy difference must be directly related
to the relative energies of the isomers.
Least stable isomer
Most stable isomer
Least energyreleased
77
Oxidation and Reduction in Organic Chemistry
Assuming the oxidation state of H is 1 and O is
-2 it is possible to calculate the oxidation
state of C in compounds containing C, H and O.
78
Oxidation and Reduction in Organic Chemistry
Oxidation of carbon corresponds to an increase in
the number of bonds between carbon and oxygen or
to a decrease in the number of carbonhydrogen
bonds. Reduction corresponds to an increase in
the number of carbonhydrogen bonds or to a
decrease in the number of carbonoxygen bonds.
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Oxidation and Reduction in Organic Chemistry
Any element more electronegative than C has the
same effect as O on the oxidation state of C.
Oxidation state of C is 2 in CH3Cl and
CH3OH. Any element less electronegative than C
has the same effect as H on the oxidation state
of C.Oxidation state of C is -4 in CH4 and
CH3Li.