2. Structural Theory Central Premises
Valency: atoms in organic compounds form a fixed number of bonds
Carbon can form one or more bonds to other carbons
Chapter 1

3. Chapter 1

4. Dot Structure IA 4A 5A 6A 7A

5. Octet rule
Ar: 1s2 2s2 2p6 3s2 3p6
Chapter 1

6. The electron configuration of atoms
K: 1s2 2s2 2p6 3s2 3p6 4s1 K: [Ar] 4s1 K+: 1s2 2s2 2p6 3s2 3p6 = Ar
Cl : 1s2 2s2 2p6 3s2 3p5
Cl- : 1s2 2s2 2p6 3s2 3p6 = Ar
Ar: 1s2 2s2 2p6 3s2 3p6

7. (1 e- in outermost shell) (7 e- in outermost shell)
give 1 e- to Na
1s2 2s2 2p6 3s1 Na
1s2 2s2 2p6 3s1 Cl
1s2 2s2 2p6 3s2 3p5 Cl
1s2 2s2 2p6 3s2 3p5
(1 e- in outermost shell)
(7 e- in outermost shell)
ionic bonding
Na+
1s2 2s2 2p6
Na+
1s2 2s2 2p6
Cl–
1s2 2s2 2p6 3s2 3p6
Cl–
1s2 2s2 2p6 3s2 3p6 8 8

8. Chemical Bonds: The Octet Rule
Atoms form bonds to produce the electron configuration of a noble gas (because the electronic configuration of noble gases is particularly stable)
For most atoms of interest this means achieving a valence shell configuration of 8 electrons corresponding to that of the nearest noble gas
Atoms close to helium achieve a valence shell configuration of 2 electrons
Atoms can form either ionic or covalent bonds to satisfy the octet rule
Chapter 1

9. Ionic Bonds
atoms gain or lose electrons to achieve the electronic configuration of the nearest noble gas
IONS are formed
Lithium → helium
Fluoride → neon
Chapter 1

10. Ionic Bonds oppositely charged ions attract and form ionic bonds
strong ionic bond (actually in a crystalline lattice)
between atoms of widely different electronegativities
ELECTRO-
NEGATIVITY
Chapter 1

11. Covalent Bonds
between atoms of similar electronegativity (close to each other in the periodic table)
Atoms achieve octets by sharing of valence electrons
Lewis structure
Chapter 1

12. Ions, themselves, may contain covalent bonds
Ions, themselves, may contain covalent bonds. Consider, as an example, the ammonium ion

13. EN diff < 0.5 EN diff= 0.5-2 EN Diff > 2
Chapter 1

14. Chapter 1

15. Electronegativity
Electronegativity is the ability of an atom to attract electrons
It increases from left to right and from bottom to top in the periodic table (noble gases excluded)
Fluorine is the most electronegative atom and can stabilize excess electron density the best
Chapter 1

16. Explaining the patterns in electronegativity
The attraction that a bonding pair of electrons feels for a particular nucleus depends on:
the number of protons in the nucleus;
the distance from the nucleus;
the amount of screening by inner electrons.
Chapter 1

17. ELECTRONEGATIVITY DECREASES AS YOU GO DOWN A GROUP
As the halogen atoms get bigger, any bonding pair gets further and further away from the halogen nucleus, and so is less strongly attracted towards it.
The larger pull from the closer fluorine nucleus is why fluorine is more electronegative than chlorine is.
As you go down a group, electronegativity decreases because the bonding pair of electrons is increasingly distant from the attraction of the nucleus.
Chapter 1

18. ELECTRONEGATIVITY I NCREASES ACROSS A PERIOD
Consider sodium at the beginning of period 3 and chlorine at the end (ignoring the noble gas, argon).
Both have their bonding electrons in the 3-level. The electron pair is screened from both nuclei by the 1s, 2s and 2p electrons
but the chlorine nucleus has 6 more protons in it.
Therefore, the electron pair gets dragged towards the chlorine….
Electronegativity increases across a period because the number of charges on the nucleus increases.
Atomic numbers: Na: 11 Cl: 17
Chapter 1

19. Writing Lewis Structures
To construct molecules the atoms are assembled with the correct number of valence electrons
The number of valence electrons = group number of the atom
Carbon is in group 4A and has 4 valence electrons
Hydrogen is in group 1A and has 1 valence electron
If the molecule is an ion, electrons are added or subtracted to give it the proper charge
The structure is written to satisfy the octet rule for each atom and to give the correct charge
If necessary, multiple bonds can be used to satisfy the octet rule for each atom
Chapter 1

20. General Rules
* Hydrogen is NEVER the central atom * The atom with the LOWEST electronegativity will be the central atom in the molecule or ion. (Element present in the least amount). * Place one pair of electrons between each bonded atoms. . * Place lone pairs of electrons about each atom (except H) to satisfy the octet rule. If electron pairs remain put them on the central atom. * If the central atom is not yet surrounded by four electron pairs, convert one of the pairs of valence electrons available on one of the surrounding atoms to or more pi bonds.
Chapter 1

21. Example Write the Lewis structure for the chlorate ion (ClO3-)
The total number of valence electrons including the electron for the negative charge is calculated
Three pairs of electrons are used to bond the chlorine to the oxygens
The remaining 20 electrons are added to give each atom an octet
Chapter 1

22. Examples
Lewis structure of CH3Br
Total number of all valence electrons: C H Br
x = 14

23. : : : 8 valence electrons remaining 6 valence electrons H H C Br H H H
: : Br : 8
valence
electrons H
remaining 6 valence electrons

24. . The carbonate ion (CO32-): _
C:4 e
3xO: 18 e
Charge: 2 e
Total: 24 valance e
The organic molecules ethene (C2H4) and ethyne (C2H2) must also use multiple bonds to satisfy the octet rule for each atom . _
Chapter 1

25. NO2- HClO - N:5 e 2xO: 12 e Charge: 1 e Total: 18 valance e H:1 e
Cl: 7 e
Total: 14 valance e
Chapter 1

26. Exceptions to the Octet Rule
The octet rule applies only to atoms in the second row of the periodic table (C, O, N, F) which are limited to valence electrons in the 2s and 2p orbitals
In second row elements fewer electrons are possible
Chapter 1

27. Exceptions to the Octet Rule
In higher rows other orbitals are accessible and more than 8 electrons around an atom are possible
Example: PCl5 , SF6 , SF4
Chapter 1

28. # of valance electrons in free atom # bonding electrons/2
Formal charge
A formal charge is a positive or negative charge on an individual atom
The sum of formal charges on individual atoms is the total charge of the molecule or ion
The formal charge is calculated by subtracting the assigned electrons on the atom in the molecule from the electrons in the neutral atom
Electrons in bonds are evenly split between the two atoms; one to each atom
Lone pair electrons belong to the atom itself
# of valance electrons in free atom
# bonding electrons/2
(#nonbonding electrons =
Formal
Charge
Chapter 1

29. Examples
Ammonium ion (NH4)+
Nitrate ion (NO3)-
Double bond
Chapter 1

30. Chapter 1

31. A Summary of Formal Charges
+NH3CH2COO-
Chapter 1

32. H2S HCN
Chapter 1

33. Resonance 4e- (from C) + 3x6e- (from 3 O) + 2e- = 24 e-
all bonds are equal in length and the charge is spread equally over all three oxygens
The real structure is a resonance hybrid or mixture of all three Lewis structures
Chapter 1

34.  Structures 1–3, although not identical on paper, are equivalent; all of its carbon–oxygen bonds are of equal length

35. curved arrows which show the movement of electrons
all structures differ only in position of electrons
electron density is spread equally among the three oxygens
electrostatic potential map
Chapter 1

36. N=5 valance-3 bonding -2 nonbonded=0 C=4 valance-4 bonding=0H H H H
N=5 valance-4 bonding = +1
C=4 valance-3 bonding- 2 nonbonding = -1

37. How to Interpret and Write Structural Formulas

38. 17A. Dash Structural Formulas
Atoms joined by single bonds can rotate relatively freely with respect to one another

39. Dot formulas are more cumbersome to draw than dash formulas and condensed formulas
Lone-pair electrons are often (but not always) drawn in, especially when they are crucial to the chemistry being discussed
Chapter 1

40. 17B. Condensed Structural Formulas

41. 17C. Bond-Line Formulas

43. Three-Dimensional Formulas
Bonds that lie in the plane of the paper = simple line
Bonds that come forward out of the plane = solid wedge
Bonds that go back out of the plane of the paper = dashed wedge
Generally to represent a tetrahedral atom:
Two of the bonds are drawn in the plane of the paper about 109o apart
The other two bonds are drawn in the opposite direction to the in- plane bonds but right next to each other
109°
in plane
Chapter 1

45. Rules for Resonance:
Individual resonance structures exist only on paper
The real molecule is a hybrid (average) of all contributing forms
Only electrons are allowed to move between resonance structures
The position of nuclei must remain the same
Only electrons in multiple bonds and nonbonding electrons can be moved
All structures must be proper Lewis structures
Chapter 1

46. The lowering of energy is called resonance stabilization
The energy of the actual molecule is lower than the energy of any single contributing form
The lowering of energy is called resonance stabilization
Equivalent resonance forms make equal contributions to the structure of the real molecule
Structures with equivalent resonance forms tend to be greatly stabilized
Chapter 1

47. Unequal resonance structures contribute based on their relative stabilities
More stable resonance forms contribute more to the structure of the real molecule
Chapter 1

48. Rules to Assign Relative Importance of Resonance Form
A resonance form with more covalent bonds is more important than one with less
Example: 6 is more stable and more important because it has more total covalent bonds
Chapter 1

49. 2. Resonance forms in which all atoms have a complete valence shell of electrons are more important
Example: 10 is more important because all atoms (except hydrogen) have complete octets
Chapter 1

50. Resonance forms with separation of charge are less important
Separation of charge cost energy and results in a less stable resonance contributor
less important
Forms with negative charge on highly electronegative atoms are more important
Those with positive charge on less electronegative atoms are also more important
Chapter 1

51. NITRATE ION All N-O bond lengths are equal
the negative charge spread over all three atoms equally
three equivalent resonance forms
Curved arrows show the movement of electrons between forms
When these forms are hybridized (averaged) the true structure of the nitrate ion is obtained
Use carboxylate for resonance
Most
electronegative
Chapter 1

52. Examples . . . . . . . . . . . . . . . . . . . . . .

53. Chapter 1

54. Quantum Mechanics
A mathematical description of bonding that takes into account the wave nature of electrons
A wave equation is solved to yield a series of wave functions for the atom
The wave functions psi (Y) describe a series of states with different energies for each electron
Wave Equations are used to calculate:
The energy associated with the state of the electron
The probability of finding the electron in a particular state
Chapter 1

55. The phase sign of a wave equation indicates whether the solution is positive or negative when calculated for a given point in space relative to the nucleus
The sign of the wave function does not indicate a greater or lesser probability of finding an electron in that location
Chapter 1

56. Constructive interference occurs when wave functions with the same phase sign interact. There is a reinforcing effect and the amplitude of the wave function increases
Destructive interference occurs when wave functions with opposite phase signs interact. There is a subtractive effect and the amplitude of the wave function goes to zero or changes sign

58. Atomic Orbitals (AOs)
The physical reality of Y is that when squared (Y 2) it gives the probability of finding an electron in a particular location in space
Plots of Y 2 in three dimensions generate the shape of s, p, d and f orbitals
Orbital: a region in space where the probability of finding an electron is large
volumes which contain the electron 90-95% of the time
Chapter 1

59. Atomic Orbitals and Electron Configuration
The greater the number of nodes in an orbital the higher its energy
2s and 2p orbitals each have one node and are higher in energy than the 1s orbital which has no nodes

60. s orbitals
Chapter 1

61. p Orbitals Chapter 1 Value of psi squared
Dots represent electron probability in a plane passing through the nucleus
Electron probability and charge density represented in three dimensions. Note that it is not spherically symmetric.
Chapter 1

62. d Orbitals
Chapter 1

63. Aufbau Principle: The lowest energy orbitals are filled first
Atoms can be assigned electronic configuration using the following rules:
Aufbau Principle: The lowest energy orbitals are filled first
Pauli Exclusion Principle: A maximum of two spin paired electrons may be placed in each orbital
Hund’s Rule: One electron is added to each degenerate (equal energy orbital) before a second electron is added
Chapter 1

64. Molecular Orbitals (MOs)
A simple model of bonding is illustrated by forming molecular H2 from H atoms and varying distance:
Region I: The total energy of two isolated atoms
Region II: The nucleus of one atom starts attracting the electrons of the other; the energy of the system is lowered
Region III: at 0.74 Å the attraction of electrons and nuclei exactly balances repulsion of the two nuclei; this is the bond length of H2
Region IV: energy of system rises as the repulsion of the two nuclei predominates
Chapter 1

65. An atomic orbital represents the region of space where one or two electrons of an isolated atom are likely to be found
A molecular orbital (MO) represents the region of space where one or two electrons of a molecule are likely to be found

66. An orbital (atomic or molecular) can contain a maximum of two spin-paired electrons (Pauli exclusion principle)
When atomic orbitals combine to form molecular orbitals, the number of molecular orbitals that result always equals the number of atomic orbitals that combine

67. A bonding molecular orbital (ymolec) results when two orbitals of the same phase overlap

68. An antibonding molecular orbital (y
An antibonding molecular orbital (y*molec) results when two orbitals of opposite phase overlap

70. Molecular orbitals of H2
Chapter 1

71. sp3 Hybridization
The structure of methane with its four identical tetrahedral bonds cannot be adequately explained using the electronic configuration of carbon
Hybridization of the valence orbitals (2s and 2p) provides four new identical orbitals which can be used for the bonding in methane
Orbital hybridization is a mathematical combination of the 2s and 2p wave functions to obtain wave functions for the new orbitals
Chapter 1

72. Structure of Methane
When one 2s orbital and three 2p orbitals are hybridized four new and identical sp3 orbitals are obtained
Each new orbital has one part s character and 3 parts p character
The four identical orbitals are oriented in a tetrahedral arrangements
The four sp3 orbitals are then combined with the 1s orbitals of four hydrogens to give the molecular orbitals of methane
Each new molecular orbital can accommodate 2 electrons
Chapter 1

73. Hybridization
sp3
sp3
hybridized
carbon
covalent
bond

74. Hybridization
sp3

75. 12A. The Structure of Methane

79. Chapter 1

80. Ethane (C2H6)
The carbon-carbon bond is made from overlap of two sp3 orbitals to form a s bond
The molecule is approximately tetrahedral around each carbon
Chapter 1

81. Generally there is relatively free rotation about s bonds
Very little energy (13-26 kcal/mol) is required to rotate around the carbon-carbon bond of ethane
Chapter 1

82. sp3 hybridized N
Chapter 1

83. The Structure of Ethene (Ethylene): sp2 Hybridization

84. sp2

85. One p orbital is left unhybridized
The three sp2 hybridized orbitals come from mixing one s and two p orbitals
One p orbital is left unhybridized
The sp2 orbitals are arranged in a trigonal planar arrangement
The p orbital is perpendicular (orthoganol) to the plane
Chapter 1

89. Ethylene
Chapter 1

90. The s orbital is lower in energy than the p orbital
The ground state electronic configuration of ethene is shown
Chapter 1

91. The Structure of Ethene (Ethylene) : sp2 Hybridization
An alkene
The geometry around each carbon is called trigonal planar
All atoms directly connected to each carbon are in a plane
The bonds point towards the corners of a regular triangle
The bond angle are approximately 120o
Chapter 1

92. Restricted Rotation and the Double Bond
There is a large energy barrier to rotation (about 264 kJ/mol) around the double bond
This corresponds to the strength of a p bond
The rotational barrier of a carbon-carbon single bond is kJ/mol
This rotational barrier results because the p orbitals must be well aligned for maximum overlap and formation of the p bond
Chapter 1

93. 13B. Cis–Trans Isomerism
Stereochemistry of double bonds

94. (trans)
(cis)
Restricted rotation of C=C

95. Cis-Trans System
Useful for 1,2 disubstituted alkenes

97. Orbitals in Boron
Chapter 1

98. The Structure of Ethyne (Acetylene): sp Hybridization

102. sp orbital
50% s character, 50% p character
sp2 orbital
33% s character, 66% p character
sp3 orbital
25% s character, 75% p character

103. sp3 sp sp2 Bond Lengths of Ethyne, Ethene and Ethane
The carbon-carbon bond length is shorter as more bonds hold the carbons together
With more electron density between the carbons, there is more “glue” to hold the nuclei of the carbons together
The carbon-hydrogen bond lengths also get shorter with more s character of the bond
2s orbitals are held more closely to the nucleus than 2p orbitals
A hybridized orbital with more percent s character is held more closely to the nucleus than an orbital with less s character sp
sp2
sp3
Chapter 1

104. Orbitals in Beryllium
Chapter 1

105. Molecular Geometry: The Valence Shell Electron Pair Repulsion (VSEPR) Model
This is a simple theory to predict the geometry of molecules
All sets of valence electrons are considered including:
Bonding pairs involved in single or multiple bonds
Non-bonding pairs which are unshared
Electron pairs repel each other and tend to be as far apart as possible from each other
Non-bonding electron pairs tend to repel other electrons more than bonding pairs do (i.e. they are “larger”)
The geometry of the molecule is determined by the number of sets of electrons by using geometrical principles
Chapter 1

106. The valence shell of methane contains four pairs or sets of electrons
To be as far apart from each other as possible they adopt a tetrahedral arrangement (bond angle 109.5o)
Chapter 1

107. Ammonia
When the bonding and nonbonding electrons are considered there are 4 sets of electrons
The molecule is essentially tetrahedral but the actual shape of the bonded atoms is considered to be trigonal planar
The bond angles are about 107o and not 109.5o because the non-bonding electrons in effect are larger and compress the nitrogen-hydrogen bond
Chapter 1

108. Water
There are four sets of electrons including 2 bonding pairs and 2 non-bonding pairs
geometry is essentially tetrahedral but the actual shape of the atoms is considered to be an angular arrangement
The bond angle is about 105o because the two “larger” nonbonding pairs compress the electrons in the oxygen-hydrogen bonds
Chapter 1

109. Boron Trifluoride
Three sets of bonding electrons are farthest apart in a trigonal planar arrangement (bond angle 120o)
Chapter 1

110. Beryllium Hydride
Two sets of bonding electrons are farthest apart in a linear arrangement (bond angles 180o)

111. Carbon Dioxide
180o
There are only two sets of electrons around the central carbon and so the molecule is linear (bond angle 180o)
Electrons in multiple bonds are considered as one set of electrons in total

112. A summary of the results also includes the geometry of other simple molecules
Chapter 1

113. sp3d and sp3d2 Hybridization
5 e- pairs
6 e- pairs
Chapter 1

114. Bond angles should be approximately 120o
Trigonal planar arrangements of atoms can be drawn in 3-dimensions in the plane of the paper
Bond angles should be approximately 120o
These can also be drawn side-on with the central bond in the plane of the paper, one bond forward and one bond back
Linear arrangements of atoms are always best drawn in the plane of the paper
Chapter 1

115. Three-Dimensional Formulas
Bonds that lie in the plane of the paper = simple line
Bonds that come forward out of the plane = solid wedge
Bonds that go back out of the plane of the paper = dashed wedge
Generally to represent a tetrahedral atom:
Two of the bonds are drawn in the plane of the paper about 109o apart
The other two bonds are drawn in the opposite direction to the in- plane bonds but right next to each other
109°
in plane
Chapter 1

117. Isomers are different molecules with the same molecular formula
Many types of isomers exist
Example
Molecular formula C2H6O
They have different structures
The two compounds differ in the connectivity of their atoms
Chapter 1

118. Constitutional Isomers
Constitutional isomers are one type of isomer
They are different compounds that have the same molecular formula but different connectivity of atoms
They often differ in physical properties (e.g. boiling point, melting point, density) and chemical properties
Chapter 1

119. Three Dimensional Shape of Molecules
It was proposed in 1874 by van’t Hoff and le Bel that the four bonds around carbon where not all in a plane but rather in a tetrahedral arrangement i.e. the four C-H bonds point towards the corners of a regular tetrahedron
Chapter 1

120. Cis-trans isomers Result of restricted rotation about double bonds
same connectivity of atoms
different arrangement of atoms in space
Stereoisomers
The molecules below do not superpose on each other
One molecule is designated cis (groups on same side) and the other is trans (groups on opposite side)
Cis-trans isomerism is not possible if one carbon of the double bond has two identical groups
Chapter 1