R-1
Review of Important Material
from Previous Courses
R-2
Previous Topics of CHEM 2311
•• Nomenclature (alkanes, alkenes, alkynes, alcohols, ethers)
•• Interpretation of 1H and 13C NMR spectra
•• Stereochemistry (chirality, isomerism, optical activity)
•• Reaction chemistry: nucleophilic substitution (SN1, SN2), elimination reactions (E1, E2), and addition reactions)
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
R-3
Models of Bonding: An Overview
•• Lewis structures
•• Valence-shell electron pair repulsion rules (VSEPR)
•• Molecular orbital theory
•• Valence bond model: Hybridization
•• Combined valence bond/molecular orbital model
1.3-4
R-4
Lewis Theory: The Octet Rule
Lewis theory of bonding (1916) was one of the earliest models to have success
Octet Rule: Main group elements are surrounded by 8 valence electrons when forming
covalent compounds.
Works well with second period elements, but runs into problems with all others!
provides no insights into molecular shapes, orbitals, or distributions of electrons
CHEM 2312 Spring 2016
F
F
Single bond: bond order =
O
O
Double bond: bond order =
N
N
Triple bond: bond order =
Notes: C.J. Fahrni
1.5
R-5
Formal Charges
Formal charges are calculated by dividing shared electrons equally between the bonded
Formal charge = valence shell electrons –– unshared electrons –– 1/2 shared electrons
How realistic are formal charges?
(CH3)3OBF4 (Meerwein reagent):
CH 3
O
CH 3
H 3C
––> Formal charges should be seen primarily as a book-keeping tool!
R-6
Oxidation Number
Oxidation numbers are calculated by assigning shared electrons to the more
electronegative atom
Oxidation number = number of electrons in valence shell
–– number of assigned electrons
General rules for assigning oxidation numbers in molecules:
–– Group 1 metals: –– Group 2 metals: –– Al is only Al3+ –– F is always F––
–– O is always O2––
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
1.7
R-7
Lewis Structures of Organic Molecules
For convenience, C-H bonds and lone pair electrons are typically not explicitly drawn in
organic structures. Note: Be careful when working with formal charges!
1.16 and Atkins/Jones: 3.1.-3.3
R-8
The Shape of Molecules: VSEPR Theory
Valence Shell Electron Pair Repulsion Theory is useful for predicting the geometry of main
group compounds
Molecules adopt the geometry for which the repulsion between electron pairs (bonding or nonbonding) are as small as possible:
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
R-9
Problem 1: Use VSEPR to predict the shape of the following molecules. Draw Lewis
structures which include lone pair electrons and formal charges and determine
the oxidation number of the central atom.
POCl 3
(CH 3)2CO
(CH 3)2SO
R-10
Atomic Structure: The Schrödinger Equation
The probability distribution and energy levels for electrons
in atoms and molecules can be calculated using the
Schrödinger equation:
H = E
H: Hamilton Operator (Energy)
E: Energy of solution : wavefunction
Erwin Schrödinger
1887-1961
Each solution of the equation corresponds to a different electron probability
distribution with a distinct energy E (The probability of nding an electron at some
point is proportional to *; with * being the complex conjugate of ).
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
Atkins/Jones: 1.9
R-11
Obital Shapes and Signs
Note: The orbital represents a probability function [(r) (r)* or 2],
whose sign is always positive. Only (r) or is a signed quantity.
For a comprehensive overview of all atomic orbitals, check out the orbitron
web-page by Mark Winter at url:
http://winter.group.shef.ac.uk/orbitron/
R-12
The Five d Orbitals
As for the p orbitals, only one d orbital (3dz2) corresponds directly to the value of ml (= 0).
The wavefunction solutions with ml = ±1 and ±2 are exponential imaginary functions. A real
wavefunction is again obtained by linear combinations of these solutions. CHEM 2312 Spring 2016
Notes: C.J. Fahrni
1.10-1.11
R-13
Molecular Orbital Theory
In principle, the electronic structure of molecules can be worked out in the same way as for
atoms:
––> solve the Schrödinger equation!
This gives molecular orbitals rather than atomic orbitals
––> compared to valence bond theory, electrons are not conned to the bonding region
between two atoms but highly delocalized
Challenge: It is difcult to solve the Schrödinger equation for molecular species (only through
approximation!)
But: Approximate MOs can be also constructed through linear combination of AOs or
group orbitals
=> qualitative molecular orbital theory (QMOT)
R-14
Diatomic Molecules
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
R-15
energy
Atkins/Jones: 3.8-10
R-16
Molecular Orbital Theory: Methane
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
1.12 and Atkins/Jones: 3.4-6
R-17
Valence Bond Theory: Hybridization
The valence bond description utilizes four equal
hybrid orbitals to describe the C-H bonds in
methane
Hybrid Orbital Description of Covalent Bonding:
Valence bond approach:
––> Bonds described as localized interactions of TWO electrons
Bonding between two atoms is described as overlap of two hybrid orbitals, which
represent the correct valence geometries
––> A hybrid orbital is a linear combination of AOs of a SINGLE atom
-> different linear combinations will result in different geometries
Molecular Orbital Theory:
Electrons are distributed over entire molecule (including core shell electrons)
1.12
R-18
Forming Hybrid Orbitals
sp hybrid orbitals
sp2 hybrid orbitals
sp3 hybrid orbitals
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
1.12
R-19
Hybridization of Hypervalent Atoms
Atkins/Jones: 3.7
R-20
Combined Valence Bond/MOTheory
Combination of hybrid orbitals provide a localized description of the bonding
interaction, each with an in-phase (bonding) and out-of-phase (anti-bonding)
interaction:
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
1.13
R-21
Example: Ethene (H2CCH2):
The remaining two 2px orbitals (Figure c) are perpendicular to the three hybrid orbitals
(Figure b) and contain one electron each
––> Delocalization yields one additional orbital with -symmetry: In-phase (bonding) overlap
1.14
R-22
Example: HCN (Hydrogen cyanide):
The two remaining p-orbitals (Figure c) contain one electron each
––> Delocalization of these two electrons yields two orbitals with -symmetry
(and a 90° angle spacing):
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
R-23
Problem 2: Assign the hybridization of all non-hydrogen atoms in the following structures.
C O
POCl 3
N
R-24
Polar Bonds: Electronegativity
Pauling (1930): Electronegativity = the ability of an atom to attract electron density towards itself in a
molecule
Based on differences in bond strength:
––> For a polar molecule A––B the strength of the A––B bond is greater than the average of the
strengths of the A––A and B––B bonds (due to an ionic contribution to the bonding)
––> This difference in bond strength, , was related to the difference in electronegativity
using the expression:
––1
= 96.49( A B )2
( in kJmol )
E is the electronegativity of element E
––> The larger the difference between the
electronegativities for a pair of atoms in
a bond, the more the bond is polarized:
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
1.3A
R-25
Periodic Trends in Electronegativity
R-26
Electronegativity and Hybridization
Radial distribution of the 2s and 2p wavefunctions:
4r2R(r)2
2s wavefunction offers better stabilization of
electrons due to increased nuclear penetration
compared to 2p
2p
2s
Within a molecule, the electronegativity of a bonded atom increases with increasing s character of the hybridized atom:
sp3 < sp2 < sp
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
R-27
Polar Bonds
C C
3.4
C O
R-28
Reactions between Electrophiles and Nucleophiles
If the wave functions of the nucleophile (typically the HOMO) and electrophile (typically the
LUMO) are matched in energy and symmetry, a reaction can occur that yields a new bond:
The nucleophile ALWAYS provides the two electrons for formation of the new bond!
CHEM 2312 Spring 2016
Notes: C.J. Fahrni
R-29
Nucleophiles and Electrophiles
Nucleophiles: contain an electron pair of relatively high energy (HOMO)
Lone pairs:
Sigma-bonds:
Pi-bonds:
ROH, RO––, R3N, R2N––, X––, RSH, R2S
typically highly polarized M-C or M-H bond: RLi, R2CuLi, RMgBr,BH4––, AlH4–– typically pi-electrons of carbon double (or triple) bond
Electrophiles: contain an relatively low-energy empty orbital (LUMO) for bond-making
Lewis acids:
Sigma-bonds:
Pi-bonds:
typically lack full octet: carbocations, boron compounds,
aluminum compounds
E-X (X = leaving group)
Leaving group ability closely related to pKa: The less basic a
leaving group is, the better it is its leaving group ability. Polarized double (or triple bonds), or C=C (or triple) bonds
attached to an electrophilic atom
3.1-2
R-30
Brønsted-Lowry Acids and Bases
A Brønsted-Lowry acid is a compound that can donate a hydrogen ion
A Brønsted-Lowry base is a compound that can accept a hydrogen ion HA
+
acid
B
base
BH+
+
conjugate acid
A––
conjugate base
Example:
O
O
+ NaOH
CHEM 2312 Spring 2016
Na
+ H 2O
Notes: C.J. Fahrni
3.5-6
R-31
The Strength of Acids and Bases
Acids are classied as strong or weak depending on whether their reaction with water
to give H3O+ (aq) go to completion or reach an equilibrium:
Ka
HA + H2O
H3
O+
+
A––
[H3O+ ][A ]
Ka =
[HA]
The acidity constant Ka (also called acid dissociation constant or acid ionization
constant) is a quantitative measure of the strength of the acid in a given solvent (in
this case water)
––> the larger Ka the stronger the acid and the weaker its conjugate base A––
For convenience, acidity constants are typically written as pKa values:
pKa = ––log Ka
R-32
Problem 3: Predict the outcome for each of the following reactions:
OH
+ NaH
+ NaNH 2
N
CHEM 2312 Spring 2016
NH 2
+ HCl (aq)
Notes: C.J. Fahrni
R-33
R-34
CHEM 2312 Spring 2016
Notes: C.J. Fahrni