Chapter 1 Carbon and Its Compounds Copyright 2018 by Nelson Education Limited 1
1.2 Organic Molecules from the Inside Out I: The Modelling of Atoms Copyright 2018 by Nelson Education Limited 2
s orbitals: p orbitals: The colour of each lobe (blue vs purple) represents its phase. Copyright 2018 by Nelson Education Limited 3
Recall: The Pauli Exclusion Principle Orbital energies: Copyright 2018 by Nelson Education Limited 4
Ground-state electron configuration for carbon: Valence orbitals: occupied orbitals of highest energy (and any accompanying empty orbitals of similar energy) Valence electrons: electrons found in valence orbitals Copyright 2018 by Nelson Education Limited 5
1.3 Organic Molecules from the Inside Out II: Bonding Copyright 2018 by Nelson Education Limited 6
Ionic bonds: occur when electrons transferred from one atom to another electrostatic attraction between opposite charges Ionic compounds (salts) Covalent bonds: arise due to sharing of electrons between atoms each bonding atom usually contributes one electron to a bond (opposite spin) Copyright 2018 by Nelson Education Limited 7
1.4 Organic Molecules Represented as Lewis Structures Copyright 2018 by Nelson Education Limited 8
Lewis structures: Formal charge: a deficit or excess of electrons Copyright 2018 by Nelson Education Limited 9
1.4.1 Formal charge method of drawing Lewis structures 1. Count total number of valence electrons in the structure based on group number of each atom. Example: Copyright 2018 by Nelson Education Limited 10
2. Charged molecules or groups: Add one electron for each negative charge and subtract one for each positive charge. Example: Example: Copyright 2018 by Nelson Education Limited 11
3. Draw connected atoms in a structure using single bonds only (do not exceed four bonds for second-row elements). Example: Copyright 2018 by Nelson Education Limited 12
4. Count the total number of bonds drawn. Ø multiply by 2 to get the total number of bonding electrons Ø subtract this number from the total number of electrons to get the number of non-bonded electrons (30 total valence electrons) (20 bonded electrons) = (10 non-bonded electrons) Copyright 2018 by Nelson Education Limited 13
5. Add non-bonded electrons to structure. Ø start with most electronegative atoms Ø continue until octet is filled, then move on to next atom until all electrons have been distributed 6. Calculate formal charges. Copyright 2018 by Nelson Education Limited 14
7. Use electron pairs from negative atoms to make extra bonds with adjacent positive atoms that do not have filled octets. Ø recalculate formal charges Ø try to generate a structure with the fewest number of charges possible Copyright 2018 by Nelson Education Limited 15
1.4.2 Exceptions to the octet rule Incomplete octets: common for first-row elements Expanded octets: common for elements in second row and beyond each has 10 valence electrons Copyright 2018 by Nelson Education Limited 16
1.5 Covalent Bonds: Overlap of Valence Atomic Orbitals Copyright 2018 by Nelson Education Limited 17
Covalent bonds are the result of overlap of atomic orbitals. Direct (head-on) overlap results in a s (sigma) bond. Copyright 2018 by Nelson Education Limited 18
Indirect (side-by-side) overlap results in a p (pi) bond. Because a p bond involves indirect orbital overlap, it is weaker than a s bond. Copyright 2018 by Nelson Education Limited 19
1.5.1 Unequal sharing of electrons: Electron-rich and electron-deficient atoms Electronegativity: the ability of an atom to pull electrons toward itself from the surrounding atoms to which it is bonded Copyright 2018 by Nelson Education Limited 20
A difference in electronegativity between two covalently bonded atoms leads to unequal sharing of electrons. The result is a bond dipole: partial charges (denoted with the symbol d) resulting from unequal electron sharing Copyright 2018 by Nelson Education Limited 21
1.6 The Shapes of Atoms in Organic Molecules Copyright 2018 by Nelson Education Limited 22
1.6.1 Three-dimensional distribution of electrons around atoms Tetrahedral geometry Example: Methane (CH 4 ) Copyright 2018 by Nelson Education Limited 23
Trigonal planar geometry Example: Formaldehyde (CH 2 O) Copyright 2018 by Nelson Education Limited 24
Linear geometry Example: Acetylene (C 2 H 2 ) Copyright 2018 by Nelson Education Limited 25
1.6.2 Predicting shape using VSEPR theory VSEPR: Valence Shell Electron Repulsion Theory most stable structure for a molecule is the one in which valence electron pairs are as far apart as possible minimizes electron-electron repulsion Copyright 2018 by Nelson Education Limited 26
Methane (CH 4 ) Copyright 2018 by Nelson Education Limited 27
Dash-wedge representation Copyright 2018 by Nelson Education Limited 28
Formaldehyde (CH 2 O) Copyright 2018 by Nelson Education Limited 29
Acetylene (C 2 H 2 ) Copyright 2018 by Nelson Education Limited 30
1.7 The Valence Bond Approach to Electron Sharing Copyright 2018 by Nelson Education Limited 31
Valence Bond Model: a bond involves the sharing of spin-paired electrons resulting bonds are localized bonds Ø confined to the region between the bonded atoms better overlap = stronger bond Copyright 2018 by Nelson Education Limited 32
1.7.1 Tetrahedral geometries require sp 3 hybrid orbitals Hybrid orbitals: atomic orbitals that form a bonding geometry by mixing 2s and 2p orbitals of the atom Copyright 2018 by Nelson Education Limited 33
sp 3 hybrid orbitals result from mixing one 2s and three 2p orbitals to form four sp 3 orbitals. Copyright 2018 by Nelson Education Limited 34
sp 3 geometry Copyright 2018 by Nelson Education Limited 35
Example: Methane (CH 4 ) Copyright 2018 by Nelson Education Limited 36
Example: Ammonia (NH 3 ) Copyright 2018 by Nelson Education Limited 37
1.7.2 Trigonal planar geometries require sp 2 hybridization Copyright 2018 by Nelson Education Limited 38
sp 2 hybrid orbitals result from mixing one 2s and two 2p orbitals to form three sp 2 orbitals Copyright 2018 by Nelson Education Limited 39
sp 2 geometry Copyright 2018 by Nelson Education Limited 40
Example: Formaldehyde (CH 2 O) Copyright 2018 by Nelson Education Limited 41
1.7.3 Linear geometries require sp orbitals Copyright 2018 by Nelson Education Limited 42
sp hybrid orbitals result from mixing one 2s and one 2p orbitals to form two sp orbitals Copyright 2018 by Nelson Education Limited 43
sp geometry Copyright 2018 by Nelson Education Limited 44
1.8 Resonance Forms: Molecules Represented by More than One Lewis Structure Copyright 2018 by Nelson Education Limited 45
Resonance: a tool to describe electron delocalization in a molecule Example: Allyl cation Delocalized electrons: Electrons not associated with a single atom; they are shared between multiple atoms Copyright 2018 by Nelson Education Limited 46
Resonance structures describe electron delocalization: resonance structures Double-headed arrow denotes resonance structures. The true structure is a hybrid of all contributing resonance structures: Copyright 2018 by Nelson Education Limited 47
Example: Allyl anion resonance structures delocalization Copyright 2018 by Nelson Education Limited 48
1.8.1 Ways to draw simple resonance forms from Lewis structures Example: Allyl cation 1. Draw Lewis structure using previously described formal charge procedure. 2. Use only single bonds, alternate positive and negative charges. Copyright 2018 by Nelson Education Limited 49
3. Construct double and triple bonds in the usual way. Include all possible structures. Copyright 2018 by Nelson Education Limited 50
1.9 Molecular Orbital Approach to Electron Sharing Copyright 2018 by Nelson Education Limited 51
Molecular orbital approach: an alternative model for describing bonding (contrast with valence bonding approach) electrons are found in molecular orbitals instead of atomic orbitals Molecular orbital: a region around the nuclei of an entire molecule in which there is a high probability of finding an electron molecular orbitals for the allyl cation Copyright 2018 by Nelson Education Limited 52
Example: Carbon monoxide (CO) Molecular orbitals result from the mixing of atomic orbitals: Anti-bonding molecular orbitals (*) result from out-of-phase overlap. Bonding molecular orbitals result from in-phase overlap. s molecular orbitals have amplitude concentrated along bonding axis. Copyright 2018 by Nelson Education Limited 53
Electrons in bonding orbitals stabilize a molecule. Electrons in anti-bonding orbitals destabilize a molecule. CO: Two electrons in s orbital vs. zero electrons in s* means there is a net stabilization of the molecule (i.e., a stable bond exists). Copyright 2018 by Nelson Education Limited 54
1.10 Other Representations of Organic Molecules Copyright 2018 by Nelson Education Limited 55
1.10.1 Condensed structure Lewis structures can be cluttered and time-consuming to draw. Example: Alternative: Condensed structure HONC Rule: Copyright 2018 by Nelson Education Limited 56
1.10.2 Line structure 1. All bonds between atoms (other than hydrogen) are drawn as solid lines in a zig-zag pattern. 2. The atomic symbols of the carbon atoms are not shown. Copyright 2018 by Nelson Education Limited 57
3. The elemental symbols of heteroatoms are shown, but lone pairs of electrons are not shown. 4. Hydrogens are not shown unless the hydrogen is connected to an explicitly drawn atom. Copyright 2018 by Nelson Education Limited 58
5. Number of hydrogens connected to each carbon is implied (number of electron pairs at each carbon must total four). 6. Lone pairs of electrons are not usually included in line structure. Wedged and hashed bonds in line structures: Projects forward (out of page) Projects backward (into page) Copyright 2018 by Nelson Education Limited 59
Chapter summary Covalent bonds in organic molecules can be classified as s or p bonds. Bonding in organic molecules is predicted using three models: Ø Lewis model Ø valence bonding model Ø molecular orbital model Resonance is a model used to illustrate electron delocalization. Organic molecules can be represented in a variety of ways: Ø Lewis structures Ø condensed structures Ø line structures Copyright 2018 by Nelson Education Limited 60