Chapter 2 Alkanes and Cycloalkanes: Introduction to ydrocarbons Copyright The McGraw-ill Companies, Inc. Permission required for reproduction or display.
2.1 Classes of ydrocarbons
ydrocarbons Aliphatic Aromatic
ydrocarbons Aliphatic Aromatic Alkanes Alkenes Alkynes
ydrocarbons Aliphatic Alkanes are hydrocarbons in which all of the bonds are single bonds. Alkanes C C
ydrocarbons Aliphatic Alkenes are hydrocarbons that contain a carboncarbon double bond. Alkenes C C
ydrocarbons Aliphatic Alkynes are hydrocarbons that contain a carboncarbon triple bond. Alkynes C C
ydrocarbons The most common aromatic hydrocarbons are those that contain a benzene ring. Aromatic
2.2 Electron Waves and Chemical Bonds
Models for Chemical Bonding The Lewis model of chemical bonding predates the idea that electrons have wave properties. There are two other widely used theories of bonding that are based on the wave nature of an electron. Valence Bond Theory Molecular Orbital Theory
Formation of 2 from Two ydrogen Atoms + e + e Examine how the electrostatic forces change as two hydrogen atoms are brought together. These electrostatic forces are: attractions between the electrons and the nuclei repulsions between the two nuclei repulsions between the two electrons
Figure 2.1 weak net attraction at long distances Potential energy + Internuclear distance
Figure 2.1 Potential energy attractive forces increase faster than repulsive forces as atoms approach each other + Internuclear distance
Figure 2.1 Potential energy 74 pm maximum net attraction (minimum potential energy) at 74 pm internuclear distance + -436 kj/mol 2 Internuclear distance
1s 1s 2 atoms: each electron "feels" attractive force of one proton 2 molecule: each electron "feels" attractive force of both protons
Figure 2.1 Potential energy 74 pm repulsive forces increase faster than attractive forces at distances closer than 74 pm + -436 kj/mol 2 Internuclear distance
Models for Chemical Bonding Valence Bond Theory constructive interference between electron waves of two half-filled atomic orbitals is basis of shared-electron bond Molecular Orbital Theory derive wave functions of molecules by combining wave functions of atoms
2.3 Bonding in 2 : The Valence Bond Model
Valence Bond Model Electron pair can be shared when half-filled orbital of one atom overlaps in phase with half-filled orbital of another.
Valence Bond Model 1s 1s in-phase overlap of two half-filled hydrogen 1s orbitals σ bond of 2
Valence Bond Model σ Bond: orbitals overlap along internuclear axis Cross section of orbital perpendicular to internuclear axis is circular.
Valence Bond Model of 2 Figure 2.4(a) The 1s orbitals of two separated hydrogen atoms are far apart. There is essentially no interaction. Each electron is associated with a single proton.
Valence Bond Model of 2 Figure 2.4(b) As the hydrogen atoms approach each other, their 1s orbitals begin to overlap and each electron begins to feel the attractive force of both protons.
Valence Bond Model of 2 Figure 2.4(c) The hydrogen atoms are close enough so that appreciable overlap of the two 1s orbitals occurs. The concentration of electron density in the region between the two protons is more readily apparent.
Valence Bond Model of 2 Figure 2.4(d) A molecule of 2. The two hydrogen 1s orbitals have been replaced by a new orbital that encompasses both hydrogens and contains both electrons.
2.4 Bonding in 2 : The Molecular Orbital Model
Main Ideas Electrons in a molecule occupy molecular orbitals (MOs) just as electrons in an atom occupy atomic orbitals (AOs). Two electrons per MO, just as two electrons per AO. Express MOs as combinations of AOs.
MO Picture of Bonding in 2 Linear combination of atomic orbitals method expresses wave functions of molecular orbitals as sums and differences of wave functions of atomic orbitals. Two AOs yield two MOs Bonding combination ψ MO = ψ() 1s + ψ(') 1s Antibonding combination ψ' MO = ψ() 1s - ψ(') 1s
Fig. 2.6: Energy-Level Diagram for 2 MOs 1s AO AO 1s
Fig. 2.6: Energy-Level Diagram for 2 MOs MO σ* antibonding σ bonding MO
C n 2n+2 2.5 Introduction to Alkanes: Methane, Ethane, and Propane Copyright The McGraw-ill Companies, Inc. Permission required for reproduction or display.
The Simplest Alkanes Methane (C 4 ) C 4 Ethane (C 2 6 ) C 3 C 3 Propane (C 3 8 ) C 3 C 2 C 3 bp -160 C bp -89 C bp -42 C
2.6 sp 3 ybridization and Bonding in Methane
Structure of Methane tetrahedral bond angles = 109.5 bond distances = 110 pm but structure seems inconsistent with electron configuration of carbon
Electron configuration of carbon 2p only two unpaired electrons should form σ bonds to only two hydrogen atoms 2s bonds should be at right angles to one another
sp 3 Orbital ybridization 2p Promote an electron from the 2s to the 2p orbital 2s
sp 3 Orbital ybridization 2p 2p 2s 2s
sp 3 Orbital ybridization 2p Mix together (hybridize) the 2s orbital and the three 2p orbitals 2s
sp 3 Orbital ybridization 2p 2 sp 3 4 equivalent half-filled orbitals are consistent with four bonds and tetrahedral geometry 2s
sp 3 Orbital ybridization
Nodal properties of orbitals p + s +
Shape of sp 3 hybrid orbitals p + take the s orbital and place it on top of the p orbital s +
Shape of sp 3 hybrid orbitals s + p + + reinforcement of electron wave in regions where sign is the same destructive interference in regions of opposite sign
Shape of sp 3 hybrid orbitals sp hybrid + orbital shown is sp hybrid analogous procedure using three s orbitals and one p orbital gives sp 3 hybrid shape of sp 3 hybrid is similar
Shape of sp 3 hybrid orbitals sp 3 hybrid + The hybrid orbital is not symmetrical. The higher probability of finding an electron on one side of the nucleus than the other leads to stronger bonds.
The C σ Bond in Methane In-phase overlap of a half-filled 1s orbital of hydrogen with a half-filled sp 3 hybrid orbital of carbon: s + + C sp 3 gives a σ bond. C σ + C
Justification for Orbital ybridization consistent with structure of methane allows for formation of 4 bonds rather than 2 bonds involving sp 3 hybrid orbitals are stronger than those involving s-s overlap or p-p overlap
2.7 Bonding in Ethane
Structure of Ethane C 2 6 C 3 C 3 tetrahedral geometry at each carbon C bond distance = 110 pm C C bond distance = 153 pm
The C C σ Bond in Ethane In-phase overlap of half-filled sp 3 hybrid orbital of one carbon with half-filled sp 3 hybrid orbital of another. Overlap is along internuclear axis to give a σ bond.
The C C σ Bond in Ethane In-phase overlap of half-filled sp 3 hybrid orbital of one carbon with half-filled sp 3 hybrid orbital of another. Overlap is along internuclear axis to give a σ bond.
C 4 10 2.8 Isomeric Alkanes: The Butanes
n-butane C 3 C 2 C 2 C 3 Isobutane (C 3 ) 3 C bp -0.4 C bp -10.2 C
2.9 igher n-alkanes
C 3 C 2 C 2 C 2 C 3 n-pentane C 3 C 2 C 2 C 2 C 2 C 3 n-exane C 3 C 2 C 2 C 2 C 2 C 2 C 3 n-eptane
2.10 The C 5 12 Isomers
C 5 12 C 3 C 2 C 2 C 2 C 3 n-pentane (C 3 ) 2 CC 2 C 3 Isopentane (C 3 ) 4 C Neopentane
ow many isomers? The number of isomeric alkanes increases as the number of carbons increase. There is no simple way to predict how many isomers there are for a particular molecular formula.
Table 2.1 Number of Constitutionally Isomeric Alkanes C 4 1 C 2 6 1 C 3 8 1 C 4 10 2 C 5 12 3 C 6 14 5 C 7 16 9
Table 2.1 Number of Constitutionally Isomeric Alkanes C 4 1 C 8 18 18 C 2 6 1 C 9 20 35 C 3 8 1 C 10 22 75 C 4 10 2 C 15 32 4,347 C 5 12 3 C 20 42 366,319 C 6 14 5 C 40 82 62,491,178,805,831 C 7 16 9
2.11 IUPAC Nomenclature of Unbranched Alkanes Copyright The McGraw-ill Companies, Inc. Permission required for reproduction or display.
Table 2.2 IUPAC Names of Unbranched Alkanes Retained: methane C 4 ethane C 3 C 3 propane C 3 C 2 C 3 butane C 3 C 2 C 2 C 3
Table 2.2 IUPAC Names of Unbranched Alkanes Note: n-prefix is not part of IUPAC name of any alkane. For example: n-butane is "common name" for C 3 C 2 C 2 C 3 ; butane is "IUPAC name." Others: Latin or Greek prefix for number of carbons + ane suffix
Table 2.2 IUPAC Names of Unbranched Alkanes Number of carbons Name Structure 5 pentane C 3 (C 2 ) 3 C 3 6 hexane C 3 (C 2 ) 4 C 3 7 heptane C 3 (C 2 ) 5 C 3 8 octane C 3 (C 2 ) 6 C 3 9 nonane C 3 (C 2 ) 7 C 3 10 decane C 3 (C 2 ) 8 C 3
Table 2.2 IUPAC Names of Unbranched Alkanes Number of carbons Name Structure 11 undecane C 3 (C 2 ) 9 C 3 12 dodecane C 3 (C 2 ) 10 C 3 13 tridecane C 3 (C 2 ) 11 C 3 14 tetradecane C 3 (C 2 ) 12 C 3 15 pentadecane C 3 (C 2 ) 13 C 3 16 hexadecane C 3 (C 2 ) 14 C 3
Table 2.2 IUPAC Names of Unbranched Alkanes Number of carbons Name Structure 17 heptadecane C 3 (C 2 ) 15 C 3 18 octadecane C 3 (C 2 ) 16 C 3 19 nonadecane C 3 (C 2 ) 17 C 3 20 icosane C 3 (C 2 ) 18 C 3 25 pentacosane C 3 (C 2 ) 23 C 3 30 triacontane C 3 (C 2 ) 28 C 3
2.12 Applying the IUPAC Rules: The Names of the C 6 14 Isomers
The C 6 14 Isomers C 3 C 2 C 2 C 2 C 2 C 3 (C 3 ) 2 CC 2 C 2 C 3 C 3 C 2 C(C 3 )C 2 C 3 (C 3 ) 2 CC(C 3 ) 2 (C 3 ) 3 CC 2 C 3
The C 6 14 Isomers C 3 C 2 C 2 C 2 C 2 C 3 exane The IUPAC name of the unbranched alkane with a chain of 6 carbons is hexane.
IUPAC Nomenclature of Branched Alkanes (Table 2.5) Step 1) Find the longest continuous carbon chain and use the IUPAC name of the unbranched alkane as the basis. Step 2) Add name of substituent as a prefix. Step 3) Number the chain from the end nearest the substituent, and identify the carbon to which the substituent is attached by number.
The C 6 14 Isomers 2-Methylpentane (C 3 ) 2 CC 2 C 2 C 3 3-Methylpentane C 3 C 2 C(C 3 )C 2 C 3
The C 6 14 Isomers 2,3-Dimethylbutane (C 3 ) 2 CC(C 3 ) 2 2,2-Dimethylbutane (C 3 ) 3 CC 2 C 3 Use replicating prefixes (di-, tri-, tetra-, etc.) according to the number of identical substituents attached to the main chain.
2.13 Alkyl Groups
Methyl and Ethyl Groups Methyl C or C 3 Ethyl C C or C 3 C 2
R Unbranched Alkyl Groups If potential point of attachment is at the end of the chain, take the IUPAC name of the corresponding unbranched alkane and replace the -ane ending with -yl. R
Unbranched Alkyl Groups If potential point of attachment is at the end of the chain, take the IUPAC name of the corresponding unbranched alkane and replace the -ane ending with -yl. R R C C C C or C 3 C 2 C 2 C 2 Butyl
Unbranched Alkyl Groups C 3 (C 2 ) 4 C 2 C 3 (C 2 ) 5 C 2 exyl eptyl C 3 (C 2 ) 16 C 2 Octadecyl
The C 3 7 Alkyl Groups C C C or C 3 C 2 C 2 and C C C or C 3 CC 3
The C 3 7 Alkyl Groups C C C or C 3 C 2 C 2 IUPAC name: Propyl Common name: n-propyl
Naming Alkyl Groups (Table 2.6) Step 1: Identify longest continuous chain starting at point of attachment. Step 2: Drop -ane ending from name of unbranched alkane having same number of carbons as longest continuous chain and replace with -yl. Step 3: Identify substituents on longest continuous chain. Step 4: Chain is always numbered starting at point of attachment.
The C 3 7 Alkyl Groups C C C or C 3 CC 3 IUPAC name: 1-Methylethyl Common name: Isopropyl
The C 3 7 Alkyl Groups C C C or C 3 C 2 C 2 Classification: Primary alkyl group Alkyl groups are classified according to the degree of substitution at the carbon that bears the point of attachment. A carbon that is directly attached to one other carbon is a primary carbon.
The C 3 7 Alkyl Groups C C C or C 3 CC 3 Classification: Secondary alkyl group Alkyl groups are classified according to the degree of substitution at the carbon that bears the point of attachment. A carbon that is directly attached to two other carbons is a secondary carbon.
The C 4 9 Alkyl Groups C C C C or C 3 C 2 C 2 C 2 IUPAC name: Butyl Common name: n-butyl Classification: Primary alkyl group
The C 4 9 Alkyl Groups C C C C or 1 2 3 C 3 CC 2 C 3 IUPAC name: 1-Methylpropyl Common name: sec-butyl Classification: Secondary alkyl group
The C 4 9 Alkyl Groups 3 C 3 2 C 3 C 1 C 2 IUPAC name: 2-Methylpropyl Common name: Isobutyl Classification: Primary alkyl group
The C 4 9 Alkyl Groups C 3 C 3 1 2 C C 3 IUPAC name: 1,1-Dimethylethyl Common name: tert-butyl Classification: Tertiary alkyl group
2.14 IUPAC Names of ighly Branched Alkanes
Branched Alkanes Octane
Branched Alkanes 4-Ethyloctane
Branched Alkanes 4-Ethyl-3-methyloctane List substituents in alphabetical order.
Branched Alkanes 4-Ethyl-3,5-dimethyloctane List substituents in alphabetical order. But don't alphabetize di-, tri-, tetra-, etc.
First Point of Difference Rule 1 2 3 4 5 6 7 8 8 7 6 5 4 3 2 1 What is the correct name? 2,3,3,7,7-Pentamethyloctane? 2,2,6,6,7-Pentamethyloctane? The chain is numbered in the direction that gives the lower locant to the substituent at the first point of difference in the names. Don't add locants!
First Point of Difference Rule 8 7 6 5 4 3 2 1 What is the correct name? 2,2,6,6,7-Pentamethyloctane? The chain is numbered in the direction that gives the lower locant to the substituent at the first point of difference in the names. Don't add locants!
C n 2n 2.15 Cycloalkane Nomenclature
Cycloalkanes Cycloalkanes are alkanes that contain a ring of three or more carbons. Count the number of carbons in the ring, and add the prefix cyclo to the IUPAC name of the unbranched alkane that has that number of carbons. Cyclopentane Cyclohexane
Cycloalkanes Name any alkyl groups on the ring in the usual way. C 2 C 3 Ethylcyclopentane
Cycloalkanes Name any alkyl groups on the ring in the usual way. List substituents in alphabetical order and count in the direction that gives the lowest numerical locant at the first point of difference. 3 C C 3 C 2 C 3 3-Ethyl-1,1-dimethylcyclohexane
2.16 Sources of Alkanes and Cycloalkanes Copyright The McGraw-ill Companies, Inc. Permission required for reproduction or display.
Crude oil
Naphtha (bp 95-150 C) C 5 -C 12 Kerosene (bp: 150-230 C) C 12 -C 15 Light gasoline (bp: 25-95 C) Crude oil C 15 -C 25 Refinery gas Gas oil (bp: 230-340 C) C 1 -C 4 Residue
Petroleum Refining Cracking converts high molecular weight hydrocarbons to more useful, low molecular weight ones Reforming increases branching of hydrocarbon chains branched hydrocarbons have better burning characteristics for automobile engines
2.17 Physical Properties of Alkanes and Cycloalkanes
Boiling Points of Alkanes Boiling points are governed by strength of intermolecular attractive forces. Alkanes are nonpolar, so dipole-dipole and dipole-induced dipole forces are absent. Only forces of intermolecular attraction are induced dipole-induced dipole forces.
Induced Dipole-Induced Dipole Attractive Forces + + two nonpolar molecules center of positive charge and center of negative charge coincide in each
Induced dipole-induced dipole Attractive Forces + + movement of electrons creates an instantaneous dipole in one molecule (left)
Induced dipole-induced dipole Attractive Forces + + temporary dipole in one molecule (left) induces a complementary dipole in other molecule (right)
Induced dipole-induced dipole Attractive Forces + + temporary dipole in one molecule (left) induces a complementary dipole in other molecule (right)
Induced dipole-induced dipole Attractive Forces + + the result is a small attractive force between the two molecules
Induced dipole-induced dipole Attractive Forces + + the result is a small attractive force between the two molecules
Boiling Points Increase with increasing number of carbons more atoms, more electrons, more opportunities for induced dipole-induced dipole forces Decrease with chain branching branched molecules are more compact with smaller surface area fewer points of contact with other molecules
Boiling Points Increase with increasing number of carbons more atoms, more electrons, more opportunities for induced dipole-induced dipole forces eptane bp 98 C Octane bp 125 C Nonane bp 150 C
Boiling Points Decrease with chain branching branched molecules are more compact with smaller surface area fewer points of contact with other molecules Octane: bp 125 C 2-Methylheptane: bp 118 C 2,2,3,3-Tetramethylbutane: bp 107 C
2.18 Chemical Properties: Combustion of Alkanes All alkanes burn in air to give carbon dioxide and water.
eats of Combustion Increase with increasing number of carbons more moles of O 2 consumed, more moles of CO 2 and 2 O formed
eats of Combustion eptane Octane Nonane 4817 kj/mol 5471 kj/mol 6125 kj/mol 654 kj/mol 654 kj/mol
eats of Combustion Increase with increasing number of carbons more moles of O 2 consumed, more moles of CO 2 and 2 O formed Decrease with chain branching branched molecules are more stable (have less potential energy) than their unbranched isomers
eats of Combustion 5471 kj/mol 5466 kj/mol 5458 kj/mol 5452 kj/mol 5 kj/mol 8 kj/mol 6 kj/mol
Isomers can differ in respect to their stability. Equivalent statement: Important Point Isomers differ in respect to their potential energy. Differences in potential energy can be measured by comparing heats of combustion.
Figure 2.17 5471 kj/mol + 25 2 O 2 5466 kj/mol 5458 kj/mol + 25 O 2 2 + 25 O 2 2 + 25 O 2 2 5452 kj/mol 8CO 2 + 9 2 O
2.19 Oxidation-Reduction in Organic Chemistry Oxidation of carbon corresponds to an increase in the number of bonds between carbon and oxygen and/or a decrease in the number of carbon-hydrogen bonds.
O increasing oxidation state of carbon O O C O O C O C C O C -4-2 0 +2 +4
increasing oxidation state of carbon C C C C C C -3-2 -1
But most compounds contain several (or many) carbons, and these can be in different oxidation states. C 3 C 2 O C 2 6 O
But most compounds contain several (or many) carbons, and these can be in different oxidation states. Working from the molecular formula gives the average oxidation state. C 3 C 2 O C 2 6 O Average oxidation state of C = -2
ow can we calculate the oxidation state of each carbon in a molecule that contains carbons in different oxidation states? C 3 C 2 O C 2 6 O Average oxidation state of C = -2
ow to Calculate Oxidation Numbers 1. Write the Lewis structure and include unshared electron pairs. C C O
ow to Calculate Oxidation Numbers 2. Assign the electrons in a covalent bond between two atoms to the more electronegative partner. C C O
ow to Calculate Oxidation Numbers 3. For a bond between two atoms of the same element, assign the electrons in the bond equally. C C O
ow to Calculate Oxidation Numbers 3. For a bond between two atoms of the same element, assign the electrons in the bond equally. C C O
ow to Calculate Oxidation Numbers 4. Count the number of electrons assigned to each atom and subtract that number from the number of valence electrons in the neutral atom; the result is the oxidation number. C C O Each = +1 C of C 3 = -3 C of C 2 O = -1 O = -2
Fortunately, we rarely need to calculate the oxidation state of individual carbons in a molecule. We often have to decide whether a process is an oxidation or a reduction.
Generalization Oxidation of carbon occurs when a bond between carbon and an atom which is less electronegative than carbon is replaced by a bond to an atom that is more electronegative than carbon. The reverse process is reduction. C X oxidation reduction C Y X less electronegative than carbon Y more electronegative than carbon
Examples Oxidation C 4 + Cl 2 C 3 Cl + Cl Reduction C 3 Cl + 2Li C 3 Li + LiCl
2.20 sp 2 ybridization and Bonding in Ethylene Copyright The McGraw-ill Companies, Inc. Permission required for reproduction or display.
Structure of Ethylene C 2 4 2 C=C 2 planar bond angles: close to 120 bond distances: C = 110 pm C=C = 134 pm
sp 2 Orbital ybridization 2p Promote an electron from the 2s to the 2p orbital 2s
sp 2 Orbital ybridization 2p 2p 2s 2s
sp 2 Orbital ybridization 2p Mix together (hybridize) the 2s orbital and two of the three 2p orbitals 2s
sp 2 Orbital ybridization 2p 2 sp 2 3 equivalent half-filled sp 2 hybrid orbitals plus 1 p orbital left unhybridized 2s
sp 2 Orbital ybridization
sp 2 Orbital ybridization p σ σ 2 sp 2 σ σ σ
π Bonding in Ethylene p the unhybridized p orbital of carbon is involved in π bonding to the other carbon 2 sp 2
π Bonding in Ethylene π Bonding in Ethylene p 2 sp 2 each carbon has an unhybridized 2 2p orbital axis of orbital is perpendicular to the plane of the σ bonds
π Bonding in Ethylene π Bonding in Ethylene p 2 sp 2 side-by-side overlap of half-filled p orbitals gives a π bond double bond in ethylene has a σ component and a π component
2.21 sp ybridization and Bonding in Acetylene
Structure of Acetylene C 2 2 C C linear bond angles: 180 bond distances: C = 106 pm CC = 120 pm
sp Orbital ybridization 2p Promote an electron from the 2s to the 2p orbital 2s
sp Orbital ybridization 2p 2p 2s 2s
sp Orbital ybridization 2p Mix together (hybridize) the 2s orbital and one of the three 2p orbitals 2s
sp Orbital ybridization 2 p 2p 2 sp 2 equivalent half-filled sp hybrid orbitals plus 2 p orbitals left unhybridized 2s
sp Orbital ybridization
sp Orbital ybridization 2 p σ σ 2 sp σ
π Bonding in Acetylene 2 p the unhybridized p orbitals of carbon are involved in separate π bonds to the other carbon 2 sp
π Bonding in Acetylene π Bonding in Acetylene 2 p 2 sp one π bond involves one of the p orbitals on each carbon there is a second π bond perpendicular to this one
π Bonding in Acetylene π Bonding in Acetylene 2 p 2 sp
π Bonding in Acetylene π Bonding in Acetylene 2 p 2 sp
2.22 Bonding in Water and Ammonia: ybridization of Oxygen and Nitrogen
Previously: Table 1.7 Ammonia trigonal pyramidal geometry N angle = 107 N : but notice the tetrahedral arrangement of electron pairs sp 3
Previously: Table 1.7 Water bent geometry O angle = 105 O : sp 3 but notice the tetrahedral arrangement of electron pairs..
sp 3 ybridization of Nitrogen 2p Mix together (hybridize) the 2s orbital and the three 2p orbitals 2s
sp 3 ybridization of Nitrogen 2p 2 sp 3 3 equivalent half-filled orbitals and a lone pair are consistent with three bonds and tetrahedral geometry 2s
sp 3 ybridization of Nitrogen nitrogen nitrogen atom Get Fig. 2.24
2.23 Which Theory of Chemical Bonding is Best?
Lewis Three Models most familiar easiest to apply Valence-Bond (Orbital ybridization) provides more insight than Lewis model ability to connect structure and reactivity to hybridization develops with practice Molecular Orbital potentially the most powerful method but is the most abstract requires the most experience to use effectively