Chapter 8 “Covalent Bonding”
Bonds are…
Section 8.1 – Molecular Compounds
…forces that hold groups of atoms together
and make them function as a unit. Two types:
1)
• OBJECTIVES:
– Distinguish between the melting points
and boiling points of molecular
compounds and ionic compounds.
– Describe the information provided by a
molecular formula.
How does H2 form?
(diatomic hydrogen molecule)
• The nuclei repel each other, since they both have a positive charge
(like charges repel).
• But, the nuclei are attracted to the
electrons
• They share the electrons, and this is
called a “covalent bond”, and
involves only NONMETALS!
+
+
2)
 Many elements found in nature are in the form of molecules:
 a neutral group of atoms joined together by covalent bonds.
 For example, air contains oxygen molecules, consisting of two
oxygen atoms joined covalently
 Called a “diatomic molecule” (O2)
Covalent bonding
• Nonmetals hold on to their valence electrons.
• They don’t give away electrons to bond.
– But still want noble gas configuration.
• Get it by sharing valence electrons with each other =
covalent bonding
• By sharing, both atoms get to count the electrons toward
a noble gas configuration.


+
+


Covalent bonding
• Compounds that are bonded covalently (like in water, or carbon
dioxide) are called molecular compounds
• Molecular compounds tend to have relatively lower melting and
boiling points than ionic compounds – this is not as strong a
bond as ionic
 Fluorine has seven
valence e A second atom also
has seven
 By sharing
8 Valence
electrons…
electrons
 …both end with
full orbitals
F F
Ionic bonds – transfer of electrons (gained or lost;
makes formula unit)
Covalent bonds – sharing of electrons. The
resulting particle is called a “molecule”
Fluorine has
seven valence eA second atom
also has seven
By sharing e-…
…both end with full orbitals
F F
8 Valence
electrons
Molecular Compounds
• Thus, molecular compounds tend to be gases or liquids at
room temperature
– Ionic compounds were solids
• A molecular compound has a molecular formula:
– Shows how many atoms of each element a molecule
contains
• The formula for water is written as H2O
– The subscript “2” behind hydrogen means there are 2
atoms of hydrogen; if there is only one atom, the
subscript 1 is omitted
• Molecular formulas do not tell any information about the
structure (the arrangement of the various atoms).
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Section 8.2
- Page 215
These are some of the
different ways to represent
ammonia:
3. The ball and stick model is
the BEST, because it shows
a 3-dimensional arrangement.
1. The molecular
formula shows
how many atoms
of each element
are present
2. The structural
formula ALSO
shows the
arrangement of
these atoms!
A Single Covalent Bond is...
•
•
•
•
A sharing of two valence electrons.
Only nonmetals and hydrogen.
Different from an ionic bond because they actually form molecules.
Two specific atoms are joined.
• In an ionic solid, you can’t tell which atom the electrons moved
from or to
• Ionic compounds
organize in a
characteristic crystal
lattice of alternating
positive and negative
ions, repeated over
and over.
Lewis Structures
• a Lewis structure represents a chemical formula: the nuclei
and inner-shell e- are represented by the element’s atomic
symbol, and covalent bonds are represented by pairs of dots
or dashes
– one pair is a single covalent bond (2 e-) that is shared
– two pair are a double covalent bond (4 e-) that is shared
– three pair are a triple covalent bond (6 e-) that is shared
• valence e- are the e- in the highest occupied energy level for
any given element
• there can also be an unshared pair of e-; it is not involved in
covalent bonding, but instead belongs exclusively to one
atom
The Nature of Covalent Bonding
• OBJECTIVES:
– Describe how electrons are shared to form covalent bonds, and
identify exceptions to the octet rule.
– Demonstrate how electron dot structures represent shared
electrons.
– Describe how atoms form double or triple covalent bonds.
– Distinguish between a covalent bond and a coordinate covalent
bond, and describe how the strength of a covalent bond is
related to its bond dissociation energy.
– Describe how oxygen atoms are bonded in ozone.
H
O
Water
 Each hydrogen has 1 valence electron
- Each hydrogen wants 1 more
 The oxygen has 6 valence electrons
- The oxygen wants 2 more
 They share to make each other complete
Note the two
“unshared pairs”
of electrons
HO
H
• It takes 2 hydrogens
to provide enough e• Every atom has full
energy levels
There are 7 diatomic
elements, KNOW THEM:
H2 N2 O2 F2 Cl2 Br2 I2
How to Draw a Lewis Structure
• add up the valence e- from each element to get the total
number of e- that can be used for bonding purposes
• when drawing a Lewis structure for a molecule with more than
two atoms, use the following guidelines for arrangement
– hydrogen and halogen atoms usually bond to only one other atom in a
molecule and are usually on the outside or end of a molecule
– the atom with the smallest electronegativity is often the central atom
– when a molecule contains more atoms of one element than the
others, these atoms often surround the central atom
• using trial and error and only the number of valence eallowed, place dots so that every element satisfies the
octet rule (except hydrogen, which only needs 2 e-)
• dots that are shared should be changed to dashes, and
unshared pairs left as dots
• when more than one Lewis structure can be drawn for a given
molecule, the molecule is said to be a resonance structure
2
Bond Dissociation Energies...
• The total energy required to break
the bond between 2 covalently
bonded atoms
• High dissociation energy usually
means the chemical is relatively
unreactive, because it takes a lot
Resonance Structures
Neither ozone structure is
correct, it is actually a hybrid of
the two. To show it, draw all
varieties possible, and join them
with a double-headed arrow.
Note the different location
of the double bond
• Occur when more than one valid Lewis structure can be written for
a particular molecule (due to position of double bond)
• Polyatomic ions – note the different positions of the double bond
• Resonance in a carbonate ion (CO32-):
of energy to break it down.
Section 8.3 – Bonding Theories
• OBJECTIVES:
– Describe the relationship between atomic and
molecular orbitals.
– Describe how VSEPR theory helps predict the shapes
of molecules.
Molecular Orbitals are...
• The model for covalent bonding assumes the orbitals are
those of the individual atoms = atomic orbital
• Orbitals that apply to the overall molecule, due to atomic
orbital overlap are the molecular orbitals
VSEPR stands for...
• Valence Shell Electron Pair Repulsion
• Predicts the three dimensional shape of molecules.
• The name tells you the theory:
– Valence shell = outside electrons.
– Electron Pair repulsion = electron pairs try to get
as far away as possible from each other.
• Can determine the angles of bonds.
• Based on the number of pairs of valence electrons,
both bonded and unbonded.
• Unbonded pair also called lone pair.
– A bonding orbital is a molecular orbital that can be occupied by
two electrons of a covalent bond
Tetrahedral and Other Shapes
• Methane (CH4) has a tetrahedral shape = 109.5°
• Carbon dioxide (CO2) has a linear shape = 180o
• Ammonia (NH3) has a
trigonal pyramidal
shape = 107o
• Water (H2O) has a
bent shape = 105o
H
H
C
H
109.5º
H
Section 8.4
Polar Bonds and Molecules
• OBJECTIVES:
– Describe how electronegativity (EN) values
determine the distribution of charge in a polar
molecule.
– Describe what happens to polar molecules
when they are placed between oppositely
charged metal plates.
– Evaluate the strength of intermolecular
attractions compared with the strength of
ionic and covalent bonds.
– Identify the reason why network solids have
high melting points.
3
Bond Polarity
Bond Polarity
• Covalent bonding means shared electrons
– but, do they share equally?
• Electrons are pulled, as in a tug-of-war, between the
atoms nuclei
– In equal sharing (such as diatomic molecules), the
bond that results is called a nonpolar covalent bond
• When two different atoms bond covalently, there is an
unequal sharing; electronegativity is the ability of an
atom in a molecule to attract shared electrons to itself
– the more EN atom will have a stronger attraction, and
will acquire a slightly negative charge
– called a polar covalent bond, or simply polar bond.
• Refer to periodic table for EN values
• Consider HCl: H = EN of 2.1 and Cl = EN of 3.0
– the bond is polar
– the chlorine acquires a slight negative charge, and the
hydrogen a slight positive charge
• Only partial charges, much less than a
true 1+ or 1- as in ionic bond
d+
d• Written as:
H Cl
• the + and – signs (with the lower case
delta: d+ and d- ) denote partial charges.
• Can also be shown:
H
Cl
– the arrow points to the more EN atom
Attractions Between Molecules
Polar molecules
• They are what make solid and liquid molecular compounds
possible.
• Intermolecular attractions are weaker than either ionic or covalent
bonds.
• There are two major types: van der Waals forces (of which there
are two) and hydrogen bonding.
• The effect of polar bonds on the polarity of the entire
molecule depends on the molecule shape
– water has two polar bonds and a bent shape; the highly
electronegative oxygen pulls the e- away from H = very
polar!
OR
OR
• When polar molecules are placed between oppositely
charged plates, they tend to become oriented with
respect to the positive and negative plates.
d+
1. Dispersion forces (van der Waals force)
• the weakest of all, caused by motion of e• increases as # e- increases
• halogens start as gases; Br is liquid; I is solid – all in Group 17
2. Dipole interaction (van der Waals force) occurs when polar molecules
are attracted to each other, like in water.
• + region of one molecule attracts the - region of another
d- molecule
• occur when polar molecules are attracted to each other
• slightly stronger than dispersion forces
• opposites attract, but not completely hooked like in ionic solids
H F
Attractions Between Molecules
Hydrogen Bonding
(Shown in water)
3. Hydrogen bonding is the attractive force caused by
hydrogen bonded to N, O, F, or Cl
• N, O, F, and Cl are very electronegative, so this is a very strong
dipole.
• And, the hydrogen shares with the lone pair in the molecule
next to it.
• This is the strongest of the intermolecular forces.
Order of Intermolecular Attraction Strengths
–
–
–
–
Weakest are the dispersion forces
A little stronger are the dipole interactions
Strongest is the hydrogen bonding
All of these are weaker than ionic bonds
d+ dH O
H d+
Hydrogen bonding
allows H2O to be a
liquid at room temp
and pressure.
This hydrogen is bonded
covalently to: 1) the highly
negative oxygen, and 2) a
nearby unshared pair.
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