Dayong Sang 桑大永
Organic Chemistry 有机化学
College of Chemical Engineering & Pharmacy
sangdy@jcut.edu.cn
A-504, Building for Chemical Engineering Experiments
sangdy@jcut.edu.cn
A-504，Building for
Chemical Engineering
Experiments
An Introduction to Organic Chemistry
▪ Orbital, electron configuration and VSEPR
▪ Common organic compounds
▪ Nomenclature of organic compounds
▪ Physical properties of organic compounds
▪ Chemical properties of organic compounds
▪ Electromeric effect and Steric Effect
Organic Chemistry--Definitions
• Old: “derived from living organisms”
• New: “chemistry of carbon compounds”
• From inorganic to organic, Wöhler, 1828
4
Chapter 1
5
Atomic Structure
• protons, neutrons, and electrons
• isotopes
12
C
6
14
6
Chapter 1
C
6
Atomic Orbitals
•
An atomic orbital is a region of space in which
there is a high probability of finding an
electron
•
Atomic orbitals are assigned with letters s, p, d
or f (smart people do fine)
•
Energy sublevels correspond to a shape where
the electron is likely to be found.
Atomic Orbitals
Orbitals: area of 90~95% of probability to find
electrons
s orbital
p orbitals
d orbitals
s orbital - spherical
1s
The electron density is highest at the nucleus and drops off exponentially with
increasing distance from the nucleus in any direction.
2s
The 2s orbital has a
small region of high
electron density
close to the
nucleus, but most
of the electron
density is farther
from the nucleus,
beyond a node, or
region of zero
electron density.
p orbitals–dumb-bell shaped
2p
There are three
2p orbitals,
oriented at
right angles to
each other.
Each is labeled
according to its
orientation
along the x, y,
or z axis.
d orbitals
f-orbitals
Energy Levels, Sublevels, and Orbitals
◼
◼
Principal energy levels – n, assigned values 1-7
(Like floors in a hotel)
Energy sublevels- s, p, d, f (Type of suite in a
hotel)
▪
▪
▪
▪
◼
s sublevel – 1 orbital
p sublevel – 3 orbitals
d sublevel – 5 orbitals
f sublevel – 7 orbitals
(Orbitals are like the number
of rooms in a suite)
Orbitals – Two electrons per orbital
(Two people per room)
Periodic Table and Electron Configurations
• Build-up order given by position on periodic table; row by row.
• Elements in same column will have the same outer shell
electron configuration.
Energy Diagram for Sublevels
◼
◼
◼
Each orbital can only
contain a maximum of 2
electrons.
Electron shells at a higher
energy level have more
orbitals.
Order of filling energy
levels: 1s2 2s2 2p6 3s2 3p6
4s2 3d10 4p6
Order of filling
Electrons in an atom:
▪ fill the lowest energy
level and orbitals
first,
▪ fill orbitals in a
particular sublevel
with one electron
each until all orbitals
are half full, and
then
▪ fill each orbital using
electrons with
opposite spins.
Electron Configurations
◼
◼
◼
➢
➢
➢
Electron configuration – the arrangement of
electrons in an atom.
Example Sodium (Na) – 1s22s22p63s1
Three rules determine electron configurations:
The Aufbau Principle The electrons occupy the lowest
energy orbitals available. The “Ground State” for an atom
is when every electron is in its lowest energy orbital.
The Pauli Exclusion Principle Each orbital can be occupied
by no more than two electrons.
Hund’s Rule When more than one orbital exists of the
same energy (p, d, and f orbitals), place one electron in
each orbital.
Writing Orbital Diagrams
The orbital diagram for
carbon has 6 electrons:
▪ 2 electrons are used to fill
the 1s orbital.
▪ 2 more electrons are used
to fill the 2s orbital.
▪ 2 electron are used in two
of the 2p orbitals so they
are half-filled, leaving the
third p orbital empty.
Electron
arrangements
in orbitals in
energy levels
1 and 2.
Orbital Diagrams
An orbital diagram shows
▪ orbitals as boxes in each sublevel.
▪ electrons in orbitals as vertical arrows.
▪ electrons in the same orbital with opposite spins (up
and down vertical arrows).
Example: orbital diagram for Li
1s2
2s1
filled half-filled
2p
empty
Electron Configuration
An electron configuration
▪ lists the filled and partially filled energy levels
in order of increasing energy.
▪ lists the sublevels filling with electrons in order
of increasing energy.
▪ uses superscripts to show the number of
electrons in each sublevel.
▪ for neon is as follows: number of electrons=10
2
2
6
1s 2s 2p
Period 1 Configurations
In Period 1, the first two electrons enter the 1s orbital.
Abbreviated Configurations
In an abbreviated configuration,
▪ the symbol of the noble gas is in brackets, representing
completed sublevels.
▪ the remaining electrons are listed in order of their sublevels.
Example: Chlorine has the following configuration:
1s22s22p63s23p5
[Ne]
The abbreviated configuration for chlorine is
[Ne]3s23p5.
Period 2 Configurations
Period 3 Configurations
Anomalous Electron Configurations
◼
A few exceptions to the Aufbau principles exist.
Stable configuration:
◼
◼
◼
half-filled d shell:
1
5
❖ Cu has [Ar]4s 3d
1
5
❖ Mo has [Kr] 5s 4d
filled d subshell:
1
10
❖ Cr has [Ar]4s 3d
1
10
❖ Ag has [Kr]5s 4d
1 14
10
❖ Au has [Xe]6s 4f 5d
Exceptions occur with larger elements where
orbital energies are similar.
Valence electrons
◼
◼
◼
Valence electrons are electrons in the outermost
orbital
For A group elements the group number corresponds
to number of valence electrons.
Electron-dot structures – Element’s symbol
surrounded by dots representing the valence electrons
8A
s and p orbitals
Bonding
Ionic Bonds
Compounds are held together by chemical bonds.
In ionic bonds, charged particles called ions are held
together because they are attracted to one another by
opposite charges via electrostatic interactions.
Bond strength / bond
dissociation: energy
required to break a
bond or energy
released to form a
bond.
Ionic bonds are often
4-7 kcal/mol in
strength.
Covalent Bonds
In covalent bonds atoms are held together because they
share electrons.
Covalent Bonds are the strongest chemical bonds, and
the energy of a typical single covalent bond is ~80
kcal/mol.
However, this bond energy can vary from ~50 to ~110
kcal/mol depending on the elements involved.
Bonding and antibonding molecular orbital
out-of-phase overlap
forms an
antibonding MO
In-phase overlap
forms a
bonding MO
sigma (s) bond
sigma (σ) bond is the strongest type of covalent
chemical bond. It is formed by head-on overlapping
between atomic orbitals.
σ bond formed by overlapping of two p orbitals
H-H bond formation
Electron distribution in covalent bonds
Electron cloud
Dipole moment
The bond dipole moment uses the
idea of electric dipole moment to
measure the polarity of a chemical bond
within a molecule.
It occurs whenever there is a separation
of positive and negative charges.
The bond dipole μ is given by:
.
Dipole moment (D) = μ = qd
q: magnitude of the charge on the atom
d: distance between the two charges
Unit of D: 1 debye = 3.33556⨯10-30 C∙m
1.0 debye results from an electron and a proton separated by
0.208 Å (Angstrom)
Dipole Moments of covalent bonds
debye
Electronegativity
Electronegativity is a measure
of the tendency of an atom to
attract a bonding pair of
electrons.
The Pauling scale is the most
commonly used.
Fluorine (the most
electronegative element) is
assigned a value of 4.0, and
values range down to cesium
and francium which are the
least electronegative at 0.7,
and calcium, 1.0 .
Fluorine
Calcium
Cesium
Electronegativity
By Physchim62 - Own work, CC BY-SA 3.0,
https://commons.wikimedia.org/w/index.php?curid=1984818
Comparison of covalent and ionic bonds
Octet rule
▪ The octet rule is a bonding theory used to predict the
molecular structure of covalently bonded molecules.
Each atom will share (covalent compounds), gain or
lose (ionic compounds) electrons in order to fill outer
electron shells with eight electrons.
▪ The octet rule is a chemical rule of thumb that states
that atoms of low atomic number (3-20) tend to
combine in such a way that they each have eight
electrons in their valence shells, giving them the
same electronic configuration as a noble gas.
▪ Example: CO2
Exceptions to the Octet rule
Lewis Structure and Kekulé formula
Lewis structures, also known as Lewis dot diagrams, Lewis dot
formulas, Lewis dot structures, electron dot structures, or Lewis
electron dot structures (LEDS), are diagrams that show the
bonding between atoms of a molecule and the lone pairs of
electrons that may exist in the molecule.
These structures can be simplified (Kekulé formula):
Formal charge
number of valence electrons
– number of lone pair electrons
– 1/2 number of bonding electrons
= Formal charge
Important Bond Numbers
(formal charge = 0)
one bond
two bonds
three bonds
four bonds
H
F
Cl
Br
O
N
C
I
Shapes of organic compounds
◼
Carbon adopts several bonding formats:
109.5o
◼
◼
120o
180o
Carbon’s electron configuration is 1s22s22p2. Its valence
orbitals are the 2s and 2p orbitals.
This model can’t explain the observed bond angles in molecules
like the ones shown above (shapes and orientations of the
valence orbitals are incorrect):
Valence Bond Theory
To explain bonding in these cases, a new model is used (called
“Valence Bond Theory”) in which atomic orbitals (2s, 2p, etc.) are
mixed to produce hybrid orbitals, which have directions that
depend on the number of atomic orbitals mixed.
Hybrid orbitals point in the same directions as electron groups in
Valence shell electron pair repulsion (VSEPR) theory.
Hybridization of orbitals
Methane (CH4)
▪ The orbitals used in bond
formation determine the bond
angles;
▪ Tetrahedral bond angle:
109.5°;
▪ Electron pairs spread
themselves into space as far
from each other as possible.
Hybrid Orbitals of Ethane (CH3CH3)
In C-C single bonds, the bond is created by the overlap of orbitals
in a head-on fashion. The situation is similar to what occurs when
two H-atoms bond (or H and Cl-atoms) is formed.
Hybrid orbitals and bonding-sp2
For a trigonal planar carbon, three atomic orbitals are
combined to make three, sp2-hybrid orbitals.
120o
An sp2-Hybridized Carbon
• The bond angle in the sp2 carbon is 120°
• The sp2 carbon is the trigonal planar carbon
Bonding in Ethene: C=C Double Bond
H
p bonds are created by the
sideways overlap of parallel,
atomic p orbitals
H
H
p bond formed by sideway overlap of two parallel p orbitals
Hybrid orbitals and bonding: H2C=O
In a molecule that contains a double bond, like
formaldehyde (H2CO):
double bond = s-bond + p-bond
sp2-hybrid orbitals are used to create the trigonal planar
molecular geometry; and
the unhybridized p orbital is used to make the p-bond
sp
180o
Hybrid orbitals and bonding-sp
Two atomic orbitals are combined to make a new
hybrid orbital set (two sp-hybrid orbitals):
180o
Bonding in Ethyne: A Triple Bond
▪ A triple bond consists of
one s bond and two p
bonds
▪ Bond angle of the sp
carbon: 180°
Bonding in the Methyl Cation, radical and anion
CH3+
CH3·
CH3-
Stabilities of carbocation, anion and radical
Most stable
Tertiary carbanion
Least stable
Secondary carbanion
Primary carbanion
Methyl anion
VSEPR theory prediction of molecular geometry
dash-wedge-line
structure
Steric number (SN)
Steric number is the number of
atoms bonded to a central atom
of a molecule plus the number of
lone pairs attached to the central
atom.
SN
Predicted Shape
Bond angle
2
Linear
180°
3
Trigonal Planar
120°
4
Tetrahedral
109.5°
5
Trigonal Bipyramidal
90°, 120°
6
Octahedral
90°
Examples and exceptions
▪ Methane (CH4) consists of carbon bonded to 4 H atoms and 0 lone pairs. SN= 4.
▪ Water (H2O) has two H atoms bonded to oxygen and also 2 lone pairs. SN= 4.
▪ Ammonia (NH3) also has a steric number of 4 because it has 3 hydrogen atoms
bonded to nitrogen and 1 lone electron pair.
▪ Ethylene (C2H4) has 3 bonded atoms and no lone pairs. Note the carbon double
bond. SN = 3.
▪ Acetylene (C2H2) – The carbons are bonded by a triple bond. There are 2 bonded
atoms and no lone pairs. SN = 2.
▪ Carbon Dioxide (CO2) is an example of a compound that contains 2 sets of double
bonds. There are 2 O atoms bonded to carbon, with no lone pairs, so the SN is 2.
Exceptions to VSEPR Theory
Valence Shell Electron Pair Repulsion theory does not always predict the correct
geometry of molecules. Examples of exceptions include:
transition metal molecules (e.g., CrO3 is trigonal bipyramidal, TiCl4 is tetrahedral)
odd-electron molecules (CH3 is planar rather than trigonal pyramidal)
some AX2E0 molecules (e.g., CaF2 has a bond angle of 145°)
some AX2E2 molecules (e.g., Li2O is linear rather than bent)
some AX6E1 molecules (e.g., XeF6 is octahedral rather than pentagonal pyramidal)
some AX8E1 molecules
s character, electronegativity and bond length
▪ The carbon-carbon bond distance decreases as the s
character of the hybrid orbitals increases.
▪ An sp3 orbital has a 25% s character, an sp2 has 33%
s character, and an sp orbital a 50% s character.
▪ The s character of a carbon is proportional to its
electronegativity.
Summary-VSEPR
▪
The greater the electron density in the region of orbital overlap,
the stronger is the bond;
▪
The more s character, the shorter and stronger is the bond:
s (100%), sp (50%), sp2 (33%), sp3 (25%)
▪
The more s character, the larger is the bond angle;
▪
A p bond is formed by two unhybridized p orbitals, and is
weaker than a s bond.