Chemistry Lecture Notes: Structure and Bonding of Alkanes and Conformational Isomerism, Lecture notes of Chemistry

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University of Melbourne

CHEM 10003

CHEMISTRY 1

LECTURE NOTES

TABLE OF CONTENTS

Topic 1: Organic Chemistry

1.1 Alkanes Pages 1 - 11 1.11 Structure and bonding alkanes 1.12 Sigma bonding 1.13 sp^3 Hybridisation 1.14 Nomenclature of saturated hydrocarbons 1.15 Structural isomerism 1.16 Conformational isomerism (torsional and steric strain) 1.17 Stereochemistry (chirality, stereogenic centre, enantiomers) 1.18 Optical activity (polarimeter, optical rotation, racemic mixtures) 1.19 Diastereoisomers and meso compounds 1.2 Cycloalkanes Pages 12 - 15 1.21 Nomenclature of cycloalkanes 1.22 Conformational isomerism of cyclohexane (axial and equatorial bonds) 1.23 Ring inversion 1.24 Conformations of substituted cycloalkanes (1,3-diaxial interactions) 1.25 Comparison of cis-trans isomerism 1.3 Alkenes Pages 15 - 19 1.31 sp^2 Hybridisation and π bonding 1.32 Nomenclature of alkenes (E/Z isomerism) 1.33 Conjugated alkenes 1.34 Benzene and derivatives (Kekulé structures and resonance) 1.35 Nomenclature of substituted benzenes 1.4 Alkynes Pages 20 - 21 1.41 sp Hybridisation 1.42 Nomenclature of alkynes 1.5 Functional Groups Pages 22 - 28 1.51 Haloalkanes (alkyl halides, aryl halides, vinyl halides) 1.52 Alcohols (diols, triols, phenols, metal alkoxides) 1.53 Ethers 1.54 Amines, thiols and nitriles 1.55 Carbonyl compounds (aldehyde, ketone, carboxylic acid, amide, ester) 1.6 Spectroscopy Pages 28 - 39 1.61 Electromagnetic spectrum 1.62 Mass spectrometry 1.63 Infrared spectroscopy 1.64 Nuclear magnetic resonance (NMR) spectroscopy

3.3 Lewis Structures Pages 77 - 86 3.31 Lewis structures (Octet rule) 3.32 Resonance structures 3.33 Dative bonds (sub-octet) 3.34 Radicals 3.35 Electron pair geometry and molecular shape 3.36 Bond polarity 3.37 Valence bond structures 3.38 Molecular orbital model (Schrodinger equation, bond order) 3.39 Paramagnetism and diamagnetism in molecules 3.4 Intermolecular Forces Pages 86 - 90 3.41 Dispersion forces 3.42 Dipole-dipole interactions 3.43 Hydrogen bonding 3.44 Polarisability 3.45 Lattice enthalpy (Born-Haber cycle) 3.46 Solubility product (common ion effect) 3.5 Solids Pages 90 - 97 3.51 Metallic bonding (x-ray diffraction) 3.52 Hexagonal close packing, cubic close packing, unit cells 3.53 Body centred cubic (BCC), alloys 3.54 Ionic solids (interstices, octahedral and tetrahedral sites) 3.55 Zinc blende, wurtzite 3.6 Network Solids Pages 98 - 100 3.61 Overview of network solids 3.62 Silicates (oxyanions and pyroxenes) 3.63 Rings and sheets 3.7 Structure of Elements Pages 100 - 107 3.71 Noble gases 3.72 Allotropic species (carbon - diamond, buckminsterfullerene) 3.73 Overview of carbon nanotubes and graphite 3.74 Silicon, phosphorous, oxygen (diradicals) 3.75 Main group chemistry (oxygen, sulphur, phosphorous, nitrogen) 3.76 Oxidation states of main elements

CHEMISTRY 1 - LECTURE NOTES

Structure and Bonding Alkanes

• (^) HYDROCARBONS are compounds of hydrogen and carbon.
• (^) Hydrocarbons that contain only single bonds are called alkanes.
• (^) The simplest alkane is methane, CH 4 ; the next is ethane, C 2 H 6 and the third member is propane, C 3 H 8. Covalent Bonding (H 2 ), Sigma Bonding and sp^3 Hybridisation
• (^) COVALENT BONDING is the sharing of electrons between atoms.
• (^) Hydrogen - atomic number: 1, mass: 1.01, 1p+, 1e-.
• (^) Two hydrogen atoms overlap atomic orbitals 1s orbitals and share a pair of electrons. The bond that forms has a circular cross section and is known as a sigma bond.
• (^) The most stable configuration for any system is the lowest energy one. As such, at the hydrogen covalent bond length of 0.74 Å, the bond energy is 436 kJ/mol. That is, to break the bond of 1 mole of H 2 molecules, 436 kJ are required.
• (^) The type of covalent bond formed depends on the shape of the overlapping orbitals.
• (^) The INTERNUCLEAR AXIS is the imaginary straight line that connects the nuclei of atoms bonded to each other in a molecule (equivalent to x -axis).
• (^) Overlap of orbitals which have circular cross section normal to the internuclear axis forms a SIGMA BOND (sharing of two electrons between two atoms).
• (^) Overlap of orbitals which have concentrations of electrons density which do not include the axis forms a PI BOND.
• (^) If two orbitals do not have matching symmetries, no bond can be formed.

• (^) All four sp^3 hybrid orbitals on each carbon atom are all now involved in sigma bonds. Nomenclature of Saturated Hydrocarbons
• (^) ALKANES are compounds of carbon and hydrogen only, in which each C atom has four tetrahedrally arranged sigma bonds.
• (^) These hydrocarbons are called SATURATED HYDROCARBONS.
• (^) The suffix -ane always indicates a saturated hydrocarbon. The formula always has the form: CnH2n+
• (^) Prefix-Parent-Suffix, where the prefix refers to the substituents, the parent refers to the number of carbons and the suffix refers to the family (alkanes = ane).
• (^) Steps to name alkanes:
  1. Name the longest linear carbon chain (with the most substituents if there are two the same).
  2. Number the atoms in the main chain from the end nearest the first branch point.
  3. Name the substituents in alphabetical order using numbers to locate on carbon chain (separated by hyphen).
  4. Name the substituents, listed below.

• (^) A simple representation of the structure of an alkane is with a LINE BOND STRUCTURE where every sigma bond is represented by a straight line.
• (^) However, line bond structures are quite cumbersome. The diagram below shows the line structure and the equivalent CONDENSED STRUCTURE for an alkane.
• (^) Many alkanes have large molecules and even the condensed structures are unwieldy.
• (^) A SKELETAL STRUCTURE is often used in which a straight line represents two C atoms joined by a sigma bond. H atoms are omitted altogether.
• (^) Each vertex or terminus represents a carbon atom, and it is assumed that enough hydrogen atoms are attached to give each carbon atom the required four bonds.
• (^) The systematic name of an alkane is determined by the length of the longest possible chain of C atoms.
• (^) Thus. although molecules I and II below, both with the same formula, could be referred to as ‘butanes’, systematic naming requires that since the longest C chain in II is only 3 carbons long it be called a propane. The full systematic name of II is 2-methylpropane.
• (^) Many of the chemical properties of simple organic molecules are determined by groups of electronegative atoms that can take part in characteristic chemical reactions.
• (^) In organic molecules, the halogens F, Cl, Br and I can replace an H atom forming a sigma bond with carbon, forming HALOALKANES (alkyl halides).
• (^) Each one can form one covalent bond to carbon (7 valence electrons).
• (^) In naming organic compounds, it is always alphabetical.

• (^) The position when the H-atoms of the two methyl groups are as far from one another as possible is called the STAGGERED CONFORMATION.
• (^) A staggered conformation is correspondent to the lowest possible energy.
• (^) The position when the H-atoms are at their closest point is called the ECLIPSED CONFORMATION.
• (^) An eclipsed conformation is less stable by around 12 kJ/mol.
• (^) There are three equivalent staggered and three equivalent eclipsed conformations in a full rotation.
• (^) A NEWMAN PROJECTION visualises chemical conformations of a carbon-carbon chemical bond from front to back, with the front carbon represented by a dot and the back carbon as a circle.
• (^) The front carbon atom is called PROXIMAL , while the back atom is called DISTAL.
• (^) Although there is always quite rapid rotation around the C-C sigma bond in ethane, the energy of the eclipsed conformation is about 12 kJ/mol higher than the energy of the staggered conformation due to repulsion between hydrogen atoms bonded to adjacent carbon atoms.
• (^) This is known as TORSIONAL STRAIN , which is the increase in potential energy of a molecule due to repulsion between electrons in bonds that do not share an atom.
• (^) As a result, this means that the molecule spends a higher proportion of its time in a staggered conformation.

• (^) Butane has a larger and more complex set of conformations associated with its constitution than does ethane.
• (^) There are two energy minima, the gauche and anti forms, which are both staggered and thus have no torsional strain.
• (^) The SYN-PERIPLANAR form is the absolute energy maximum, in which the two CH3 groups are as close as possible to one another. As the two CH3 groups cannot occupy the same space, they repel one another and this increases the energy of conformation.
• (^) ANTICLINAL has less steric and torsional strain for eclipsed conformations.
• (^) The ANTI-PERIPLANAR form is the absolute energy minimum, since the SYNCLINAL (GAUCHE) form has a small steric interaction between the two methyl groups.
• (^) At a dihedral angle of 60 degrees, one hydrogen of each of the methyl groups is relatively close to a hydrogen of the other methyl groups (van der Waals repulsion).
• (^) There are also two energy maxima, both of which are eclipsed and thus torsionally strained. The higher energy conformation also has steric strain.
• (^) STERIC STRAIN is the increase in potential energy of a molecule due to repulsion between electrons in atoms that are not directly bonded to each other (close than their radii allow).
• (^) As two groups cannot occupy the same space, they repel one another and this increases the energy of the conformation.