Transcript Chapter 8 – Covalent Bonding

Chapter 8 – Covalent Bonding
Ms. Wang
Lawndale High School
Chapter 8.1 – Molecular Compounds
In Chapter 7, we learned about electrons
being “given up” or “stolen away”
This type of “tug of war” bond between a
metal and nonmetal is called an ionic bond
In this chapter, you will learn about another
type of bond
Covalent Bond – atoms held together by
sharing electrons between two nonmetals
Molecules
We know that a metal cation and nonmetal
anion are joined together by an ionic bond
and called a SALT
A neutral group of atoms joined together by a
covalent bond is called a MOLECULE
Monatomic vs. Diatomic Molecules
Molecules can be monatomic or diatomic
Diatomic Molecule – a molecule consisting
of two atoms
There are 7 diatomic molecules on the
periodic table (SUPER 7) – N2, O2, F2, Cl2, Br2,
I2 , H 2
Properties of Molecular Compounds
Lower Melting and Boiling Points than
Ionic Compounds
Gases or liquids at room temperature
Molecular Formulas
Molecular Formula – the chemical formula of a
molecular compound
It shows how many atoms of each element a
molecule contains
Example
H2O
contains 3 atoms (2 atoms of H, 1 atom of O)
C2H6 contains 8 atoms (2 atoms of C, 6 atoms of H)
Practice
How many atoms total and of each do the
following molecular compounds contain?
1. H2
2. CO
3. CO2
4. NH3
5. C2H6O
Chapter 8.2 – Covalent Bonding
Remember that ionic compounds either gain
or lose electrons in order to attain a noble
gas electron configuration
Covalent compounds form by sharing
electrons to attain a noble gas electron
configuration
So the Octet Rule still applies to covalent
bonds
Single Covalent Bond
Single Covalent Bond - two atoms held
together by sharing one pair of electrons
Unshared Pair / Lone Pair / Nonbonding Pair
– a pair of valence electrons that is not shared
between atoms
Let’s Practice (you can use dots or lines)

F2, H2O, CH4
Double and Triple Covalent Bonds
Double Covalent Bond – a bond that
involves two shared pairs of electrons
Triple Covalent Bond – a bond that involves
three shared pairs of electrons
Let’s Practice
O2
 N2

Polyatomic Ion
Polyatomic Ion - tightly bound group of
atoms that has a positive or negative charge
and behaves as one unit

Some examples are NH4+ and SO32-
Bond Dissociation Energy
Bond Dissociation Energy - the energy
required to break the bond between two
covalently bonded atoms

A large bond dissociation energy
corresponds to a strong covalent bond

For example, carbon-carbon has a strong
bond dissociation energy so it is not very
reactive
Chapter 8.3 - Bonding Theories
So far, the orbitals we have been discussing
are atomic orbitals (s, p, d, f) for each atom
When two atoms combine, their atomic orbitals
overlap and make molecular orbitals
Molecular Orbitals – orbitals that apply to
the entire molecule instead of just one
atom
Molecular Orbitals

Just as atomic orbitals belong to a particular
atom, a molecular orbital belongs to a
molecule as a whole
Each orbital is filled with 2 electrons
Orbital – a molecular orbital that can be
occupied by two electrons of a covalent bond
Bonding
Sigma Bond ()
Sigma Bond - when 2 atomic orbitals
combine to form a molecular orbital that is
symmetrical around the axis
S orbitals overlapping
P orbitals overlapping end-to-end
Pi Bond ()
Bonding electrons likely to be found in a
sausage-shape above and below the axis
 Pi bonds tend to be weaker than sigma
bonds because pi bonds overlap less than
sigma bonds
P orbitals overlapping
side-by-side
VSEPR Theory
Explains the 3D shape of molecules

According to VSEPR theory, the repulsion
between electron pairs causes molecular
shapes to adjust so that the valence
electron pairs stay as far apart as possible
A Few VSEPR Shapes
Nine possible molecular shapes
VSEPR Theory

Unshared pairs of electrons (lone pairs)
are very important in predicting the shapes
of molecules

Examples
 Methane (CH4) - tetrahedral
 Ammonia (NH3) - pyramidal
 Water (H2O) – bent
 Carbon Dioxide (CO2) - linear
Hybrid Orbitals
VSEPR is good at describing the molecular
shapes, but not the types of bonds formed
Orbital hybridization provides information
about both molecular bonding and
molecular shape

In hybridization, several atomic orbitals
mix to form the same total number of
hybrid orbitals
Bond Hybridization
Hybridization Involving Single Bonds – sp3 orbital

Ethane (C2H6)
Hybridization Involving Double Bonds – sp2 orbital

Ethene (C2H4)
Hybridization Involving Triple Bonds – sp orbital

Ethyne (C2H2)
Chapter 7.4 – Polar Bonds and Molecules
There are two types of covalent bonds
 Polar Covalent Bonds
 Nonpolar Covalent Bonds
• Polar Covalent Bond – unequal sharing of
electrons where one atom has a slightly negative
charge and the other atom has a slightly positive
charge (HCl, H2O)
• Nonpolar Covalent Bond – equal sharing of
electrons between two atoms (Cl2, N2, O2)
Classification of Bonds
You can determine the type of bond
artificially by calculating the difference
in electronegativity between elements
Type of Bond
Nonpolar Covalent
Electronegativity
Difference
0  0.4
Polar Covalent
0.5  1.9
Ionic
2.0  4.0
Let’s Practice Together
What type of bond is HCl? (H = 2.1, Cl = 3.1)
Difference = 3.1 – 2.1 = 1.0
Therefore it is polar covalent bond.
Your Turn To Practice

N(3.0) and H(2.1)
 Al(1.5)

H(2.1) and H(2.1)

Mg(1.2) and O(3.5)

Ca(1.0) and Cl(3.0)

H(2.1) and F(4.0)
and Cl(3.0)
Dipole
• No bond is purely ionic or covalent …
they have a little bit of both characters
When there is unequal sharing of electrons a
dipole exists
Dipole is a molecule that
has two poles or regions
with opposite charges
Represented
by a
dipole arrow pointing
towards the more
negative end.
Practice Drawing Dipoles
P- Br
P = 2.1
Br = 2.8
P –Br
+
Practice
H(2.1) – S(2.5)
 C(2.5) – F(4.0)
 Si(1.8) – C(2.5)
 N(3.0) – O(3.5)

-
Attractions Between Molecules
Intermolecular attractions are weaker than
either ionic or covalent bonds
Van der Waals forces – consists of the two weakest
attractions between molecules
dipole interactions – polar molecules attracted
to one another

dispersion forces – caused by motion of
electrons (weakest of all forces)

Hydrogen Bond

Hydrogen Bonds - attractive forces in
which a hydrogen covalently bonded to a
very electronegative atom is also weakly
bonded to an unshared electron pair of
another electronegative atom
Hydrogen Bond

This other atom may be in the same
molecule or in a nearby molecule, but
always has to include hydrogen
Hydrogen Bonds have about 5% of the
strength of an average covalent bond
 Hydrogen Bond is the strongest of all
intermolecular forces

Intermolecular Attractions

A few solids that consist of molecules do
not melt until the temperature reaches
1000ºC or higher called network solids
(Example: diamond, silicon carbide)

Network Solid – solids in which all of the
atoms are covalently bonded to each other
• Melting a network solid would require breaking
covalent bonds throughout the solid
Homework
Chapter 8 Assessment Page 247
#’s 39-41, 43-46, 51, 53, 54, 57-59, 61, 65, 68,
83, 85, 86, 89, 96, 99, 100