SLE133 Lecture Notes - Lecture 4: Chemical Polarity, Coordinate Covalent Bond, Covalent Bond
11 May 2018
School
Course
Professor
We will be covering chapter 4 of the McMurry textbook.
4.1 Covalent Bonds
4.2 Covalent Bonds and the Periodic Table
4.3 Multiple Covalent Bonds
4 4 Coordinate Covalent Bonds
Week 4a beginning 26 March 2018
Week 4b beginning 05 April 2018
J. Gordon and Deakin College, 2018. SLE133 trimester 1
4
.
4
Coordinate
Covalent
Bonds
4.5 Characteristics of Molecular Compounds
4.6 Molecular Formulas and Lewis Structures
4.7 Drawing Lewis Structures
4.8 The Shapes of Molecules
4.9 Polar Covalent Bonds and Electronegativity
4.10 Polar Molecules
4.11 Naming Binary Molecular Compounds
1. What is a covalent bond?
Be able to describe the nature of covalent bonds and how
they are formed.
2. How does the octet rule apply to covalent bond formation?
B blt th tt lt ditth b f
Goals
B
e a
bl
e
t
o use
th
e oc
t
e
t
ru
l
e
t
o
p
re
di
c
t
th
e num
b
ers o
f
covalent bonds formed by common main group elements.
3. What are the major differences between ionic and
molecular compounds?
Be able to compare the structures, compositions, and
properties of ionic and molecular compounds.
Chapter 4
4. How are molecular compounds represented?
Be able to interpret molecular formulas and draw Lewis
structures for molecules.
5. What is the influence of valence-shell electrons on
molecular sha
p
e?
Goals
p
Be able to use Lewis structures to predict molecular
geometry. Covered in more detail in SLE155
6. When are bonds and molecules polar?
Be able to use electronegativity and molecular geometry to
predict bond and molecular polarity.
Chapter 4
Covalent bond—A bond formed by sharing
electrons between atoms
Molecule—A group of atoms held together by
covalent bonds
Covalent Bonds
Chapter 4
Page 99
Main group elements undergo reactions that
leave them with eight valence electrons (or
two for hydrogen), so that they have a noble
gas electron configuration.
Covalent Bonds
gas
electron
configuration.
Nonmetals can achieve an electron octet by
sharing an appropriate number of electrons in
covalent bonds.
Chapter 4
Covalent bonding in hydrogen (H2):
•Spherical 1sorbitals overlap to give an egg-shaped
region.
•There are two electrons between the nuclei, providing
1
2
fi ti f h li
Covalent Bonds
1
s
2
con
fig
ura
ti
on o
f
h
e
li
um.
•H-H, H:H and H2all represent a hydrogen molecule.
Chapter 4
Page 100
Slides based on McMurry chapter 4
Molecular Compounds
Page 1 of 9
•Bond length is the optimum distance
between nuclei in a covalent bond.
•If atoms are too far apart, attractive forces are small
and no bond exists.
•If atoms are too close, the repulsive interaction
Covalent Bonds
between nuclei is so stron
g
that it
p
ushes the atoms
apart.
•There is an optimum point where net attractive
forces are maximised and where the molecule is most
stable.
•In the H2molecule, this optimum distance between
nuclei is 74 pm.
Chapter 4
Chlorine also exists as a diatomic molecule
due to the end-on overlap of 3porbitals.
Covalent Bonds
There are seven elements which exist
naturally as diatomic molecules: nitrogen,
oxygen, hydrogen, fluorine, chlorine, bromine,
and iodine. Learn them!
Chapter 4
Page 101
Covalent Bonds
Learn these because they will be useful when you write
chemical equations
Chapter 4
Figure 4.2, page 101
A molecular compound is a compound that
consists of molecules rather than ions.
Covalent Bonds and the Periodic Table
In these compounds, each atom shares enough
electrons to achieve a noble gas configuration
or filled octet (except hydrogen). Chapter 4
Page 102
Numbers of covalent bonds typically formed by main group
elements to achieve octet configurations.
Covalent Bonds and the Periodic Table
Chapter 4
Figure 4.3, page 102
Compare with your table of covalencies
from pages for study
Exceptions to the Octet Rule:
Boron has only three electrons to share, so will
form compounds with six shared electrons instead
of the usual eight.
Covalent Bonds and the Periodic Table
of
the
usual
eight.
Elements in the third row and below in the
periodic table have vacant dorbitals that can be
used for bonding.
These elements can share more than eight
electrons (covered in detail in SLE155). Chapter 4
Slides based on McMurry chapter 4
Molecular Compounds
Page 2 of 9
The bonding in some molecules cannot be explained
by the sharing of only two electrons between atoms.
Multiple Covalent Bonds
The only way these molecules’ atoms can have outer-
shell electron octets is by sharing more than two
electrons between pairs of atoms, resulting in the
formation of multiple covalent bonds.
Chapter 4
Page 104
Single bond—A covalent bond formed by
sharing one electron pair (2 electrons).
Represented by a single line: H–H
Double bond—A covalent bond formed by
hi t lt i (4lt )
Multiple Covalent Bonds
s
h
ar
i
ng
t
wo e
l
ec
t
ron pa
i
rs
(4
e
l
ec
t
rons
)
.
Represented by a double line: O=O
Triple bond—A covalent bond formed by
sharing three electron pairs (6 electrons).
Represented by a triple line: N≡N
Chapter 4
Carbon, nitrogen, and oxygen are the elements most
often present in multiple bonds.
Carbon and nitrogen can form double and triple bonds.
Oxygen forms only double bonds.
Multiple covalent bonding is particularly common in
Multiple Covalent Bonds
organic molecules, which consist mainly of the
element carbon.
Note that in compounds containing multiple bonds,
carbon still forms a total of four covalent bonds,
nitrogen still forms a total of three covalent bonds,
and oxygen still forms a total of two covalent bonds.
Chapter 4
A coordinate covalent bond is the covalent
bond that forms when both electrons are
donated by the same atom.
Coordinate Covalent Bonds
Chapter 4
Page 106
Once formed, a coordinate covalent bond is
no different from any other covalent bond.
Coordinate covalent
bonds often result in
Coordinate Covalent Bonds
unusual bonding
patterns, such as
nitrogen with four
covalent bonds, or
oxygen with three
covalent bonds (H3O+).
Chapter 4
Page 107
Note that these molecules have a charge
Ionic compounds have high melting and
boiling points because the attractive forces
between oppositely charged ions are so strong.
Ml l
lhi
Characteristics of Molecular Compounds
M
o
l
ecu
l
es are neutra
l
, so t
h
ere
i
s no strong
attraction to hold them together.
There are weaker forces between molecules,
called intermolecular forces.
Chapter 4
Slides based on McMurry chapter 4
Molecular Compounds
Page 3 of 9
Document Summary
Be able to compare the structures, compositions, and properties of ionic and molecular compounds. Be able to interpret molecular formulas and draw lewis structures for molecules: what is the influence of valence-shell electrons on molecular shape? p. Be able to use lewis structures to predict molecular geometry. Be able to use electronegativity and molecular geometry to predict bond and molecular polarity. Covalent bond a bond formed by sharing electrons between atoms. Molecule a group of atoms held together by covalent bonds. Main group elements undergo reactions that leave them with eight valence electrons (or two for hydrogen), so that they have a noble gas electron configuration. gas electron configuration. Nonmetals can achieve an electron octet by sharing an appropriate number of electrons in covalent bonds. Covalent bonding in hydrogen (h2): (cid:149) spherical 1s orbitals overlap to give an egg-shaped region. (cid:149) there are two electrons between the nuclei, providing.
