Covalent
Bonding &
Shapes of
Molecules
Unit 1
1-1
Organic Chemistry
 The
study of the compounds of carbon
 Over 10 million compounds have been identified
• about 1000 new ones are identified each day!
 C is a small atom
• it forms single, double, and triple bonds
• it is intermediate in electronegativity (2.5)
• it forms strong bonds with C, H, O, N, and
some metals
1-2
History of Organic Chemistry

Organic Chemistry – the chemistry of carbon compounds
• Swedish chemist Torbern Bergman (in 1770) was the
first to distinguish organic compounds (compounds
coming from living things) from inorganic compounds.
• Organic compounds were believed to have a “vital
force”.
• Michel Chevreul (in 1816) showed that organic
compounds, without a vital force, could be turned into
other organic compounds when making soap.
• Friedrich Wöhler (in 1828) isolated an “organic”
compound from “inorganic” material, accidentally!
1-3
Usefulness of Organic Chemistry
 Why
is organic useful to you?
• You are organic (DNA, proteins,
carbohydrates, lipids, etc.)
• Biological sciences (medicine, pharmacy,
etc.)
• Polymer chemistry
• Food chemistry
• Nanotechnology
• …survive college organic and do well on the
MCATs
1-4
What makes carbon unique?



There are more carbon compounds than compounds of all
the other elements combined.
Carbon forms very strong (high bond energy) covalent bonds
with many different types of atoms allowing long chains to
form.
Carbon can be found as many “allotropes” or coaxed into
different arrangements (Note: Not all compounds listed below are
considered organic.)
•
•
•
•
•
•
Rings (cyclo- compounds)
Graphite (pure C)
Diamond (pure C)
Buckeyballs (spheres of pure C)
Nanotubes (tubes of pure C)
Polymers (Styrofoam®)
1-5
Schematic View of an Atom
• a small dense nucleus,
diameter 10-14 - 10-15 m,
which contains positively
charged protons and
most of the mass of the
atom
• an extranuclear space,
diameter 10-10 m, which
contains negatively
charged electrons
1-6
Atomic Structure
Structure of the atom:
• made of protons, neutrons and electrons
• atomic number (Z) – number of p+
• mass number (M) – number of p+ + no
example: 12C
example: hydrogen-2
• atoms are neutral, so p+ = e-
1-7
Electron Configuration of Atoms
 Electrons
are confined to regions of space called
principle energy levels (shells)
• each shell can hold 2n2 electrons (n = 1,2,3,4......)
N u mb er of
Relative En ergies
Electrons S hell
of Electrons
S hell
Can Hold
in Thes e Shells
h igh er
4
32
3
18
2
8
1
2
low er
1-8
Electron Configuration of Atoms
 Shells
are divided into subshells called orbitals,
which are designated by the letters s, p, d, f,........
• s (one per shell)
• p (set of three per shell 2 and higher)
• d (set of five per shell 3 and higher) .....
S hell
O rb itals Contain ed in Th at S hell
3
3s , 3p x , 3p y , 3p z, p lu s five 3d orbitals
2
2s , 2p x , 2p y , 2p z
1
1s
1-9
Electron Configuration of Atoms
 Aufbau
Principle:
• orbitals fill in order of increasing energy from lowest
energy to highest energy
 Pauli
Exclusion Principle:
• only two electrons can occupy an orbital and their
spins must be paired
 Hund’s
Rule:
• when orbitals of equal energy are available but there
are not enough electrons to fill all of them, one
electron is added to each orbital before a second
electron is added to any one of them
1-10
Lewis Dot Structures
 Gilbert
N. Lewis
 Valence shell:
• the outermost occupied electron shell of an atom
 Valence
electrons:
• electrons in the valence shell of an atom; these
electrons are used to form chemical bonds and in
chemical reactions
 Lewis
dot structure:
• the symbol of an element represents the nucleus and
all inner shell electrons
• dots represent valence electrons
1-11
Lewis Dot Structures

Lewis Dot Structures for Elements 1-18
1A
2A
3A
4A
5A 6A
7A
H.
He :
:N e :
.
: :
Al:
:F :
:
.
.
:Cl : :A r :
:
N a. M g :
B:
.
.
.
C : N : : O. :
.
.
.
.
Si : P : :S. :
:
Be :
.
:
.
Li .
8A
1-12
Lewis Model of Bonding
 Atoms
bond together so that each atom acquires
an electron configuration the same as that of the
noble gas nearest it in atomic number
• an atom that gains electrons becomes an anion
• an atom that loses electrons becomes a cation
• the attraction of anions and cations leads to the
formation of ionic solids
• an atom may share electrons with one or more atoms
to complete its valence shell; a chemical bond formed
by sharing electrons is called a covalent bond
• bonds may be partially ionic or partially covalent;
these bonds are called polar covalent bonds
1-13
Electronegativity
 Electronegativity:
• a measure of an atom’s attraction for the electrons it
shares with another atom in a chemical bond
 Pauling
scale
• generally increases left to right in a row
• generally increases bottom to top in a column
1-14
Electronegativity Table
1-15
Formation of Ions
A
rough guideline:
• ions will form if the difference in electronegativity
between interacting atoms is 1.9 or greater
• example: sodium (EN 0.9) and fluorine (EN 4.0)
• we use a single-headed (barbed) curved arrow to show
the transfer of one electron from Na to F
+
F
••
+
Na
••
-
•
•
•
•
Na
•
•
••
F
••
• in forming Na+F-, the single 3s electron from Na is
transferred to the partially filled valence shell of F
Na(1s22s 22p 63s1 ) + F(1s 22s2 2p5 )
Na+(1s2 2s22p 6) + F-(1s2 2s2 2p6 )
1-16
Covalent Bonds
 The
simplest covalent bond is that in H2
• the single electrons from each atom combine to form
an electron pair
H•
+
•H
H-H
H 0 = -435 kJ (-104 kcal)/mol
• the shared pair functions in two ways simultaneously;
it is shared by the two atoms and fills the valence shell
of each atom
 The
number of shared pairs
• one shared pair forms a single bond
• two shared pairs form a double bond
• three shared pairs form a triple bond
1-17
Polar and Nonpolar Covalent Bonds
 Although
all covalent bonds involve sharing of
electrons, they differ widely in the degree of
sharing
 We divide covalent bonds into nonpolar covalent
bonds & polar covalent bonds
 Bonds are polarized along a continuum from
covalent to ionic depending on the difference in
electronegativity. A polar covalent bond has
some amount of partial + and – charges at either
end.
A-A
δ+A-B δA+B|____________________________________________________________|
Covalent
Polar Covalent
Ionic
1-18
Polar and Nonpolar Covalent Bonds
• an example of a polar covalent bond is that of H-Cl
• the difference in electronegativity between Cl and H is
3.0 - 2.1 = 0.9
• we show polarity by using the symbols d+ and d-, or by
using an arrow with the arrowhead pointing toward the
negative end and a plus sign on the tail of the arrow at
the positive end
d+
H
dCl
H
Cl
1-19
Polar Covalent Bonds
 Bond
dipole moment (m):
• a measure of the polarity of a covalent bond
• the product of the charge on either atom of a polar
bond times the distance between the nuclei
• Table 1.7 shows average bond dipole moments of
selected covalent bonds
Bond
Dipole
Bond (D )
Bond
Dipole
Bond (D )
Bond
D ipole
Bond (D)
H-C
H-N
H-O
H-S
C-F
C-Cl
C-Br
C-I
C-O
C=O
C-N
-C=N
0.3
1.3
1.5
0.7
1.4
1.5
1.4
1.2
0.7
2.3
0.2
3.5
1-20
Lewis Structures
 To
write a Lewis structure
•
•
•
•
determine the number of valence electrons
determine the arrangement of atoms
connect the atoms by single bonds
arrange the remaining electrons so that each atom has
a complete valence shell
• show a bonding pair of electrons as a single line
• show a nonbonding pair of electrons as a pair of dots
• in a single bond atoms share one pair of electrons, in a
double bond they share two pairs of electrons, and in a
triple bond they share three pairs of electrons
1-21
General Rules for Drawing Lewis Structures
 Add
up total number of valence electrons
(the ones that bond)
 Determine the central atom
 Determine bonding scheme (HONC-1234)
 Distribute remaining electrons to follow the
octet rule (lone pairs)
1-22
Lewis Structures
H-O-H
H 2 O (8)
Water
H
H
H
C2 H 4 (12)
Ethylene
•
•
•
•
•
H-N-H
H
N H 3 (8)
Ammonia
H
C C
 In
H
H-C-H
H
CH 4 (8)
Meth ane
H-Cl
HCl (8)
Hyd rogen ch loride
O
H
H-C C-H
C O
C2H 2 (10)
Acetylen e
H
CH 2O (12)
Formald ehyde
H
O
C
O
H
H 2CO 3 (24)
Carbonic acid
neutral molecules
hydrogen has one bond
carbon has 4 bonds and no lone pairs
nitrogen has 3 bonds and 1 lone pair
oxygen has 2 bonds and 2 lone pairs
halogens have 1 bond and 3 lone pairs
1-23
HONC - 1234
1-24
Formal Charge
 Formal
charge: the charge on an atom in a
molecule or a polyatomic ion
 To derive formal charge
1. write a correct Lewis structure for the molecule or ion
2. assign each atom all its unshared (nonbonding)
electrons and one-half its shared (bonding) electrons
3. compare this number with the number of valence
electrons in the neutral, unbonded atom
Formal
charge
N umber of
= valence electrons
in th e neutral,
un bonded atom
All
One h alf of
un shared + all sh ared
electrons
electrons
1-25
Formal Charges
charges – these charges do not imply the
presence of actual ionic charges …instead they
give insight into the distribution of electrons
 calculating the formal charges of each atom in a
molecule will help you determine the best, most
favorable structure (lowest energy)
 formal
1-26
General Rules of Stability
Lewis structures that approximate the actual
molecule most closely are those that have:
 maximum number of covalent bonds
 minimum separation of unlike charges
 formal charges of zero are ideal
 placement of any negative charges on the most
electronegative atom (or any positive charge on
the most electropositive atom)
• Ex. Oxygen would rather 1- then 1+
1-27
Lewis Structures – Single Bonding
 In
a Lewis structure, each valence electron is
symbolized by a dot, bonding electrons are
symbolized by lines (non-bonding electrons
are drawn as a pair of dots).
• Try to arrange ALL of the following compounds so
they have a noble gas configuration (full octet).
Draw the Lewis structures for
CH4
CH3NH2
CH3CH2OH
CH3Cl
1-28
Lewis Structures – Multiple Bonds


The sharing of one pair of electrons is called a single
bond.
• Sharing of 2 pairs is a double bond.
• Sharing of 3 pairs is a triple bond.
In some compounds, the ONLY way to satisfy all
element’s octets is to use multiple bonds.
 Ex.
Draw Lewis structures for C2H4, CH2O
and C2H2
1-29
Lewis Structures – Single Bonding
 Try these:
Ex. CH3CH2COCH2NH2
Ex. CH3CO2CH=CHCH3
Ex. CH3CHOHCH2CONHCH2CH3
1-30
Lewis Structures – More Complex
Ex. CH3(CH2)3COOH
Ex. CH3CHOHCH2CONHCH2CH3
Ex. CH3CO2C(CH3)=CHCH3
1-31
Formal Charge

Example: Draw Lewis structures, and show which atom in
each bears the formal charge
(a) NH2
(b) HCO3
(c) CO3
2-
+
(d) CH3 NH3
(e) HCOO
(f) CH3 COO
1-32
Exceptions to the Octet Rule
 Molecules
containing atoms of Group 3A
elements, particularly boron and aluminum
B
Boron trifluoride
: Cl
Al
: Cl :
:
:
:F:
: Cl :
: :
: :
:F
6 electrons in the
valence shells of boron
and aluminum
:
:
: F:
Aluminum chloride
1-33
Exceptions to the Octet Rule
 Atoms
of third-period elements have 3d orbitals
and may expand their valence shells to contain
more than 8 electrons
• phosphorus may have up to 10
:
H- O-P- O-H
:
Cl :
:
Cl :
:O:
: :
: Cl
P
: : : :
O-H
Phosphoric
acid
:
CH3
Trimethylphosphine
: : : :
:
CH3 -P- CH3
: Cl
: Cl :
Phosphorus
pentachloride
1-34
Exceptions to the Octet Rule
• sulfur, another third-period element, forms compounds
in which its valence shell contains 8, 10, or 12
electrons
H-O- S-O-H
:
CH 3 -S-CH 3
:
:
:
:
H-S- H
:
: O:
:
: O:
:O :
Hydrogen
sulfi de
Dimethyl
sulfoxide
Sul furi c
acid
1-35
VSEPR
 Based
on the twin concepts that
• atoms are surrounded by regions of electron density
• regions of electron density repel each other
H
C
H
H
H
H
H
C
O
N
H
: :
: :
O C
C
H
O
H
H
C
H
H
H
:
H
2 region s of e - d ensity
(lin ear, 180°)
H H
H
H
C
N
: :
3 region s of e - d ensity
(trigon al planar, 120°)
:
4 region s of e - d ensity
(tetrahed ral, 109.5°)
O H C C H H C N
1-36
Polar and Nonpolar Molecules
 To
determine if a molecule is polar, we need to
determine
• if the molecule has polar bonds
• the arrangement of these bonds in space
 Molecular
dipole moment (m): the vector sum of
the individual bond dipole moments in a
molecule
• reported in debyes (D)
1-37
Polar and Nonpolar Molecules
 these
molecules have polar bonds, but each has
a zero dipole moment
Cl
F
O
C
B
O
F
F
Carbon dioxide
m=0D
Boron trifluoride
m=0D
C
Cl
Cl
Cl
Carbon tetrachloride
m=0D
1-38
Polar and Nonpolar Molecules
 these
molecules have polar bonds and are polar
molecules
direction
of dip ole
moment
N
O
H
H
Water
m = 1.85D
H
H
Ammonia
m = 1.47D
H
direction
of dip ole
moment
1-39
Polar and Nonpolar Molecules
• formaldehyde has polar bonds and is a polar molecule
direction
of dip ole
moment
H
O
C
H
Formaldehyde
m = 2.33 D
1-40
Resonance
 For
many molecules and ions, no single Lewis
structure provides a truly accurate
representation
O
O
H3 C C
and
O
-
H3 C C
O
Ethan oate ion
(acetate ion)
1-41
Resonance
 Linus
Pauling - 1930s
• many molecules and ions are best described by
writing two or more Lewis structures
• individual Lewis structures are called contributing
structures
• connect individual contributing structures by doubleheaded (resonance) arrows
• the molecule or ion is a hybrid of the various
contributing structures
1-42
Resonance
 Curved
arrow: a symbol used to show the
redistribution of valence electrons
 In using curved arrows, there are only two
allowed types of electron redistribution:
• from a bond to an adjacent atom
• from an atom to an adjacent bond
1-43
Resonance

Examples: equivalent contributing structures
CH3
O:
CH3
O:
C
: O :-
:
N itrite ion
(equivalent con trib uting
s tru ctures)
C
:
:O :-
:
:
O:
: O :-
O:
:N
:
:N
:
:
:
:O: -
A cetate ion
(equ ivalen t contributin g
s tru ctures)
1-44
Resonance Structures
 Draw
resonance structures for NO3_
_
O
_
O
O
O
N
N
N
O
O
O
O
O
 The
“real” structure is a resonance hybrid
 Each oxygen has a partial negative charge
1-45
Resonance Structures

The “real” structure is said to have its electrons
delocalized and is represented by a dotted bond.
Remember, different resonance structures are not
always equivalent.
Why would the one on the left be least influential?

1-46
Let’s consider methane (CH4)






How many valence electrons?
4
In which orbitals?
2s22p2
How many unpaired electrons
(from orbital diagram)?
2 – indicating the possibility for 2 bonds
But I thought carbon always formed 4 bonds???
So, both the 2s and 2p orbitals are used to form bonds
How many bonds does carbon
form?
4
How do we explain this?
Hybridization
1-47
Hybridization
 The
s and p orbitals of the C atom combine with
each other to form hybrid orbitals before they
combine with orbitals of another atom to form a
covalent bond.
 Lone pairs MUST be accounted for in
hybridization of orbitals since they are also
negatively charged regions causing repulsion
(regions of electron density).
1-48
Resonance
 All
contributing structures must
1. have the same number of valence electrons
2. obey the rules of covalent bonding
• no more than 2 electrons in the valence shell of H
• no more than 8 electrons in the valence shell of a 2nd
period element
• 3rd period elements, such as P and S, may have up to
12 electrons in their valence shells
3. differ only in distribution of valence electrons; the
position of all nuclei must be the same
4. have the same number of paired and unpaired electrons
1-49
Electronic Structure of the Atom
 Electrons
are found in orbitals.
• An orbital is a region of space where
there is a high probability of finding an
electron.
• It is designated by ψ2 (ψ is a wave with
a + and – sign, so we use ψ2 so we’re
always dealing in the positive region).
1-50
Region of Electron Density
 Regions
of electron density spread out threedimensionally around an atom due to
repulsions.
 Types
•
•
•
•
of regions of electron density
single bond
double bond
triple bond
lone pair
1-51
Orbitals – the 1s and 2s
Recall quantum numbers:
• The first two energy levels
can hold 2 and 8 electrons,
respectively
• First level: 1s-orbital
• Electron density is a function
of distance from the nucleus.
• Highest density is at the
nucleus.
• Second level: 2s-orbital, 2porbital
• Region of space where there is
no electron density is a node.
• Most of the density is farther
away, so 2s is higher in energy
than the 1s.
1-52
Shapes of 1s and 2s Orbitals
distribution (2) for 1s and 2s orbitals
showing an arbitrary boundary surface
containing about 95% of the electron density
 Probability
1-53
Shapes of a Set of 2p Atomic Orbitals
 Three-dimensional
shapes of 2p atomic orbitals
1-54
Hybrid Orbitals
 The
Problem:
• bonding by 2s and 2p atomic orbitals would give bond
angles of approximately 90°
• instead we observe bond angles of approximately
109.5°, 120°, and 180°
A
Solution
• hybridization of atomic orbitals
• 2nd row elements use sp3, sp2, and sp hybrid orbitals
for bonding
1-55
Hybrid Orbitals
 Hybridization
of orbitals (L. Pauling)
• the combination of two or more atomic orbitals forms a
new set of atomic orbitals, called hybrid orbitals
 We
deal with three types of hybrid orbitals
sp3 (one s orbital + three p orbitals)
sp2 (one s orbital + two p orbitals)
sp (one s orbital + one p orbital)
 Overlap
of hybrid orbitals can form two types of
bonds depending on the geometry of overlap
 bonds are formed by “direct” overlap
 bonds are formed by “parallel” overlap
1-56
sp3 Hybrid Orbitals
• each sp3 hybrid orbital
has two lobes of
unequal size
• the sign of the wave
function is positive in
one lobe, negative in the
other, and zero at the
nucleus
• the four sp3 hybrid
orbitals are directed
toward the corners of a
regular tetrahedron at
angles of 109.5°
1-57
sp3 Hybrid Orbitals
• orbital overlap pictures of methane, ammonia, and
water
1-58
sp2 Hybrid Orbitals
• the axes of the three sp2 hybrid orbitals lie in a plane
and are directed toward the corners of an equilateral
triangle
• the unhybridized 2p orbital lies perpendicular to the
plane of the three hybrid orbitals
1-59
Development of Chemical Bonding Theory
August Kekulé and Archibald Couper proposed
that carbon is tetravalent (1858). Carbon
always forms 4 bonds to make a stable
compound.
Multiple bonding was proposed when Emil
Erlenmeyer showed acetylene (C2H2) to have a
triple bond (1862) and Crum Brown showed
ethylene (C2H4) to have a double bond (1864).
August Kekulé determined that carbon chains
can link end to end to become rings,
such as benzene (1865).
1-60
Development of Chemical Bonding Theory
Jacobus van’t Hoff proposed that the four bonds
of carbon are not random, but are threedimensionally arranged with specific direction
(1874). He helped determine that the hydrogens
in methane are positioned at the corners of a
tetrahedron, 109.5o.
1-61
Bond Formation: The Octet Rule
G.N. Lewis (1916) proposed theories about how
atoms form bonds
• Atoms transfer or share electrons in such a way as to
attain a filled valence shell of electrons (the Octet
Rule).
 Covalent
bonding involves the sharing of
electrons.
• Equal sharing: non-polar bond; Ex: C-C or C-H
• Unequal sharing: polar bond; Ex: C-O or O-H
 Ionic
bonding involves the loss of an electron
due to a large difference in electronegativity
(>2.0).
1-62
Formation of an sp hybrid
2p
2s
1s
Mathematically rearranged…
2p
2sp
1s
1-63
Formation of an sp hybrid
 In
the formation of acetylene (C2H2), each carbon
becomes sp hybridized
2p
2sp
1s
These electrons pair to
make 2  bonds
2p
These electrons pair
to make a  bond
2sp
1s
2
hydrogens come in and pair in the lowest
energy orbitals available (the 2sp)
1-64
Bonding in Acetylene, C2H2
1-65
Hybrid Orbitals
Group s Orb ital Predicted
Bond ed Hyb rid Bond
to Carb on ization
A ngles
4
sp 3
109.5°
Types of
Bond s
to Carbon
Examp le
N ame
4 sigma b on ds
HH
H-C-C-H
Ethan e
HH
2
2
sp
sp
2
120°
180°
3 sigma b on ds
and 1 p i bond
2 sigma b on ds
and 2 p i bonds
H
H
C
H
Eth ylen e
C
H
H-C C-H
Acetylene
1-66
Bond Length
So, why are double bonds shorter than single bonds?
Which orbital is closer to the nucleus, s or p?
What is the %s character in an sp3 hybridized orbital?
What is the %s character in an sp2 hybridized orbital?
The one with more %s character will be closer to the
nucleus, thus the bond between two of these orbitals
will be closer.
1-67
Bond Strength
So, why are double bonds higher in energy than single
bonds?
Which bond involves the higher energy orbitals  or ?
How many electrons are involved in a double bond
compared to a single bond?
2p
You would think that a double bond is twice the bond
strength as 2sp
a single
bond, but due to the poor
2
overlap of the p-orbitals ( bond), that overlap is not
a strong as the  bonds. So, the double bond is less
than double
1sthe strength of a single bond.
1-68
Bond Lengths and Bond Strengths
Orb ital
Overlap
Bond Length
(pm)
Bon d Strength
[k J (k cal)/mol]
N ame
Formula
Bond
C-C
sp 3 -sp 3
153.2
376 (90)
Ethan e
HH
H-C-C-H
C-H
sp 3 -1s
111.4
422 (101)
H
C-C
sp 2-sp 2, 2p-2p
133.9
727 (174)
H
C-H
sp 2-1s
110.0
464(111)
C-C
sp-sp, tw o 2p -2p
C-H
sp-1s
121.2
109.0
966 (231)
556 (133)
HH
H
Ethylene
C C
H
Acetylene H-C C-H
1-69