Name ___________________________________ Date ______________________ Class ______________________
Chapter 8 – Covalent Bonding
Guided Reading
Section 1 – Molecular Compounds - This section explains how to distinguish between ionic and molecular
compounds.
As you read Chapter 8 Section 1 define the following words:
Covalent bond –
Molecule –
Diatomic molecule –
Molecular compound -
Molecular formula –
As you read Chapter 8 Section 1 complete the following questions:
How is a covalent bond different from an ionic bond?
Which have higher melting points, ionic or covalent? Why?
What do molecular formulas tell us?
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What is the molecular formula for ethane?
In Figure 8.4 what information about the ethane molecule given by its formula C2H6 is also given by the drawing?
Label the following pictures with the correct labels: space-filling molecular model, ball-and-stick molecular model,
perspective drawing, molecular formula, structural formula
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In Figure 8.6 which of these molecules has the greatest number of oxygen atoms?
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Section 2 – The Nature of Covalent Bonding – This section uses electron dot structure to show the formation of single,
double, and triple covalent bonds. It also describes and gives examples of coordinate covalent bonding, resonance
structures, and exceptions to the octet rule.
As you read Chapter 8 Section 2 define the following words:
Single covalent bond –
Structural formula –
Unshared pair –
Double covalent bond –
Triple covalent bond –
Coordinate covalent bond –
Polyatomic ion –
Bond disassociation energy –
Resonance structure –
As you read Chapter 8 Section 2 complete the following questions:
Why is the octet rule important in covalent bonding?
What does a structural formula represent?
What is the electron dot structure of a methane molecule?
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In a structural formula how are double bonds represented? Triple bonds?
How do we represent a coordinate covalent bond?
How can bond disassociation energy help us?
Use the bond dissociation energies of H2 (435 kJ/mol) and of a typical carbon-carbon bond (347 kJ/mol) to decide which
bond is stronger. Explain your reasoning.
Why might resonance structures be important?
What notation is used to show that the two covalent bonds in O3 are the same?
List three ways in which the octet rule can sometimes fail to be obeyed?
Give two examples of exceptions to the octet rule.
In Figure 8.12 how many valence electrons does the sulfur in sulfur hexafluoride (SF6) have for the structure shown in
the figure?
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Section 3 – Bonding Theories – This section describes the molecular orbital theory of covalent bonding, including orbital
hybridization. It also explains the use of the VSEPR theory to predict the shapes of some molecules.
As you read Chapter 8 Section 3 define the following words:
Molecular orbitals –
Bonding orbital –
Sigma (δ) bond –
Pi (π) bond –
Tetrahedral angle –
VSEPR Theory –
Hybridization –
As you read Chapter 8 Section 3 complete the following questions:
How are molecular orbitals different from atomic orbitals?
Tetrahedral comes from the Greek tetra-, meaning "four" and hedra, meaning "face." The Greek pod means "foot."
What do you think tatrapod means?
In Figure 8.16 how do the resulting H-N-H bond angles compare to the tetrahedral angle?
List the nine most common molecular shapes.
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Explain how VSEPR theory can be used to predict the shapes of molecles?
Why are the bonds formed by the 2s and the 2p electrons of carbon the same in methane?
In Figure 8.20 what region of space does the pi bond occupy relative to the carbon atoms?
In Figure 8.21 how many pi bonds are formed in an ethyne molecule?
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Section 4 – Polar Bonds and Molecules – This section explains the use of electronegativity values to classify a bond as
nonpolar covalent, polar covalent, or ionic. It also names and describes the weak attractive force that holds groups of
molecules together.
As you read Chapter 8 Section 4 define the following words:
Nonpolar covalent bond –
Polar covalent bond –
Polar bond –
Polar molecule –
Dipole –
Van Der Waals forces –
Dipole interactions –
Dispersion forces –
Hydrogen bonds –
Network solids –
As you read Chapter 8 Section 4 complete the following questions:
Fill in the missing information from Table 8.3 below.
Electronegativity difference range
Most probable type of bond
0.0-0.4
Example
Nonpolar covalent
0.4-1.0
1.0-2.0
Very polar covalent
≥2.0
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In Figure 8.23 which atom is more electronegative, a chlorine atom or a hydrogen atom?
How is bond polarity determined?
Compare and contrast polar and nonpolar bonds.
In Figure 8.24 what would happen if, instead, carbon dioxide molecules were placed between the plates? why?
Why are intermolecular forces important?
Which are stronger intermolecular forces or ionic bonds?
List the intermolecular forces from weakest to strongest?
What are hydrogen bonds?
What is one example of a network solid?
Read the Technology & Society on the Chemistry of Adhesives. What are two things you found interesting.
What information is presented in Table 8.4?
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