Chapter 6
Honors Chemistry 1
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Section 6.1: Covalent Bonding
Review of IONIC BONDING:
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Ionic bonds result from TRANSFER of electrons
o In other words, the loss or gain of electrons is essential to ionic bonding.
o Cations LOSE one or more electron, anions GAIN one or more electron
Ionic bonding typically occurs between metals and nonmetals
Ionic compounds are NOT molecules!!
Example: Consider KBr
 A potassium atom LOSES an electron, and a bromine atoms GAINS an electron
 Both atoms have obtained a stable noble gas electron configuration in the process
What is a COVALENT BOND, and how is it different than an IONIC BOND?
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A covalent bond is the SHARING of an electron pair.
Covalent bonding has nothing to do with electric charge. When analyzing covalently-bonded substances,
our focus shifts to valence electrons!
Covalent bonding usually involves nonmetals
“Binary covalent compounds” include two different types of atoms that are sharing electrons
Example: Consider H2
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This simplified picture summarizes the difference between ionic compounds and molecules
What is a “Molecule?”
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A molecule is an assembly of one or more atoms tightly bound together by a covalent bond.
o A molecule can be thought of as a “package of atoms” that behaves in many ways as a single,
distinct object, just like your cell phone is a single object composed of multiple different
parts.
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A “diatomic molecule” is a molecule that is made up of two of the same atoms.
o Seven elements on the periodic table exist naturally as diatomic molecules:
 H2, N2, O2, F2, Cl2, Br2, I2
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The “Magnificent Seven” includes all of the diatomic molecules that exist naturally!
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A “molecular compound” is a compound containing more than one type of atom that possesses covalent
bonds. (e.g., H2O, CO2).
Almost all molecular substances that we will encounter contain only nonmetals!!
The “Octet Rule”
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All of the second row nonmetals (including carbon) want to obtain a stable electron configuration, and
they need to obtain 8 valence electrons to do so.
The Octet Rule requires that these elements must be surrounded by 8 electrons
Nonmetals can form stable molecules by SHARING electrons
o Remember, ionic compounds form from the loss or gain of electrons, but molecules form when
electrons are SHARED between atoms!
Noble gases don’t form molecules because they already have a stable octet…in other words, they are
already equipped with 8 valence electrons.
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Chapter 6
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Visualizing Covalent Bonds:
You can almost imagine each atom as having a hand with an electron in it…a covalent bond is when the two
atoms hold hands, with two electrons wrapped inside of their grip.
The molecular orbital shown to the right is a region of high probability of where the electron will be…don’t
forget that electrons behave as waves…the molecular orbital is a region in between the two atoms where the
electrons are likely to be.
Electronegativity in Molecules:
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Recall that electronegativity is an atoms tendency to pull electrons closer to itself in a molecular bond.
Fluorine is the KING of electronegativity!! In other words, it doesn’t like to share!
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Chapter 6
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CLASSIFY EACH OF THE FOLLOWING COMPOUNDS ACCORDING TO BOND TYPE (USE THE
INFORMATION ABOVE TO DO SO!)
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NO
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CO
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HF
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NaCl
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HBr
•
NaI
Important Keys to molecular bonding:
o If the electronegativity difference between two atoms is 0.0-0.5, the molecule will be a nonpolar
covalent molecule (e.g., like gasoline, which is made of carbon and hydrogen atoms)
 The electrons are shared equally in a nonpolar covalent molecule
o If the electronegativity difference is between 0.5-2.1, the molecule will be polar covalent (e.g.,
WATER!!!)
 The electrons are NOT shared equally in a polar covalent bond…the atom with the higher
electronegativity holds the electrons closer to itself
o Electronegativity differences greater than 2.1 indicate ionic compound
 There is no electron sharing, there was a TRANSFER OF ELECTRONS!!
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Chapter 6
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Polar molecules:
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Polar molecules (e.g., water) have positive and negative ends, known as “dipoles”
o How can we know this? Simply look at the electronegativity difference between the atoms!
In a polar molecule, the more electronegative atom pulls the shared electron pair much closer to itself.
Example: Show the dipoles of the following polar molecules
1) HF
2) HCl
Molecular bonds can vary in strength
Based on the information shown above, as electronegativity decreases, bond energy for a given molecule
______________________
Suppose that a molecule had an electronegativity difference of 1.6. What would you predict as the bond
energy associated with that molecule? ______________
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Chapter 6
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Section 6.2: Drawing and naming molecules
Drawing Lewis Structures:
Never begin pairing dots until all four sides (north, south, east, and west) have a dot!!
Examples: Draw the Lewis Structure for each of the atoms listed below
1) Potassium
2) Sulfur
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3) Iodine
4) Argon
5) Barium
6) Aluminum
7) Silicon
8) Sodium
Octet Rule and Molecular Bonds:
Octet Rule:
Single Bonding Examples:
1) Draw a Lewis structure for F2
2) Draw a Lewis structure for Cl2
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3) Draw a Lewis structure for H2O
4) Draw a Lewis structure for CH4
5) Draw a Lewis structure for NH3
6) Draw a Lewis structure for HF
7) Draw a Lewis structure for CF4
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8) Draw a Lewis structure for CH3I
9) Draw a Lewis structure for CH2Cl2
10) Draw a Lewis structure for NO+
11) Draw a Lewis structure for the sulfate ion, SO42-
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Multiple Bond Examples:
1) Draw a Lewis structure for O2
2) Draw a Lewis structure for N2
3) Draw a Lewis structure for CO2
4) Draw a Lewis structure for HCN
5) Draw a Lewis structure for formaldehyde, CH2O
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Resonance Structures:
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Some molecules, such as ozone, O3, cannot be represented by just one single Lewis structure.
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When a molecule has two or more possible Lewis structures, the two structures are called resonance
structures.
Draw the two resonance structures for ozone, O3
Structure 1
Structure 2
Draw the THREE resonance structures for the Nitrate ion, NO3-
Structure 1
Structure 2
Structure 3
Naming Molecular Compounds:
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Chapter 6
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Honors Chemistry 1
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The first listed element is the LESS-ELECTRONEGATIVE element
The second element’s name will end in –ide
For molecular compounds, you will use the prefixes shown below to indicate the number of each atom
present in the molecule
A DETAILED EXAMPLE:
Examples:
1) N2O
2) NO2
3) H2O
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4) CCl4
5) Li3N
6) PF5
Bond Lengths and Bond Types
Bond
C≡C
C=C
C–C
Bond
Length (A)
1.20
1.34
1.54
Bond
C≡N
C=N
C–N
Bond
Length (A)
1.16
1.38
1.43
Bond
C≡O
C=O
C–O
Bond
Length (A)
1.13
1.23
1.43
Bond
N≡N
N=N
N–N
Bond
Length (A)
1.10
1.24
1.47
In general, as the number of bonds between two particular atoms increases, the bond length ________________
Which have shorter bonds, C-C or N-N? __________
Arrange the following bond lengths from shortest to longest: C = O, C ≡ C, N = N, C – N, C ≡ N
Which bond would you suspect to be stronger, a C ≡ C or a C – C ? Why?
Section 6.3: Molecular Shapes
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Chapter 6
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The three-dimensional shape of a molecule is important in determining the molecule’s physical and
chemical properties.
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You can predict the shape of a molecule by examining the Lewis structure of the molecule.
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The __________________________________________________________________ is a theory that
predicts some molecular shapes based on the idea that pairs of valence electrons surrounding an atom
repel each other.
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According to the VSEPR theory, the shape of a molecule is determined by the valence electrons
surrounding the central atom.
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Electron pairs are negative, so they repel each other.
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Therefore, the shared pairs that form different bonds repel each other and remain as far apart as
possible.
Example: Consider the BF3 molecule
Example: Consider the CH4 molecule
Determining Molecular Shape (most common cases):
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Chapter 6
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A Complete Perspective (you only need to know up to five electron domains):
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Chapter 6
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HOW TO APPLY THE VSEPR MODEL:
1) Draw the Lewis Structure (the single atom is typically the center atom)
2) Count the electron pairs and arrange them in ways that minimizes the repulsion (in other words, as far
away as possible)
3) Determine the positions of the atoms from the way the electron pairs are shared
4) Determine the name of the molecular structure from the positions of the atoms
Examples: Predict the molecular geometry of the following using the VSEPR Theory
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Chapter 6
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1) PBr3
2) PCl5
3) SO3
4) CH3+
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Chapter 6
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5) O3
6) SeCl2
7) SnCl3-
8) SF4
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Chapter 6
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9) IF5
10) CO32-
Extra Practice Problems and Test Review
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Chapter 6
Honors Chemistry 1
Wilmington High School
1) Classify the bonding in each of the following molecules as ionic, polar covalent, or nonpolar covalent:
a. H2
b. K3P
c. NaI
d. SO2
e. HF
f. CCl4
g. CF4
h. K2S
2) Write the Lewis Structures that obey the octet rule for each of the following molecules: (the central atom
is always the first atom listed)
a. NCl3
b. ICl
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c. CO
d. SCl2
e. POCl3
f. PO43-
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Chapter 6
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h. ClO3-
i. ClO2-
j. SO2Cl2
3) Write Lewis Structures and predict the molecular geometries for the following molecules:
a. OCl2
b. BeH2
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Chapter 6
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c. SO2
d. SO3
e. NF3
f. IF3
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