CHM 101 – Chapter Eight
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Chemical Bonds, Lewis Structures & the
Octet Rule
Ionic Bonding
Covalent Bonding
Bond Polarity & Electronegativity
Drawing Lewis Structures
Resonance Structures
Exceptions to the Octet Rule
Strengths of Covalent Bonds
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Lewis Symbols
Lewis symbols display atoms using their symbol surrounded
by its valence electrons depicted as dots.
Name
Electronic
Configuration
Lewis Symbol
Potassium
Bromine
Carbon
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Ionic Compounds
• When metals react with nonmetals, electrons are
transferred from the metal to the nonmetal, forming a cation
and an anion.
• Ionic compounds such as sodium chloride {NaCl(s)} are
large arrays of cations and anions arranged so ions of
opposite charge are as close as possible.
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Ionic Compounds
The lattice energy is the energy required to break a crystal apart into
the independent ions.
• The potential energy of attraction between the ions depends directly on
the charge of the ions, and inversely on the distance between ions.
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Covalent Compounds
• In covalent compounds, atoms share electrons to achieve
a stable electronic configuration.
• While ionic compounds are typically brittle, crystalline and
possess high melting points, covalent compounds tend to be
gases, liquids or solids with low melting points.
• Bonds involve the interaction of charged species, with like
charges repelling and opposite charges attracting.
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Bond Polarity & Electronegativity
• In homonuclear diatomic molecules such as fluorine {F2(g)},
the bonding electrons are shared equally by both nuclei.
• When two different atoms are bonded together, the electrons
are often unequally shared. This is true for hydrogen fluoride
{HF(g)}.
• An unequal electron distribution results in a separation of
positive and negative charge; the bond is said to be polar.
F2(g)
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HF(g)
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Bond Polarity & Electronegativity
• To quantify the extent of bond polarity, Pauling assigned
each atom an electronegativity: The attraction for the shared
electrons the atom displays when involved in a bond.
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Bond Polarity & Electronegativity
• Electronegativity generally increases from left to right and
from bottom to top of the periodic Table. Fluorine, the most
electronegative element, is assigned a value of 4.0
• The larger the electronegativity difference between the
atoms involved in the bond, the more polar the bond. Ionic
bonds represent the extreme case
F2(g)
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HF(g)
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Drawing Lewis Structures
• Lewis structures depict molecules as collections of atoms
bonded together by shared electron pairs (depicted as lines).
• In most structures, the combination of shared and lone pairs
of electrons provides each atom (except hydrogen) with an
octet of valence electrons. These structures obey the "octet
rule"
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Lewis Structure of CCl4
1.
2.
3.
4.
Sum the valence electrons from all atoms. Add electrons to account
for negative charge, subtract to account for postitive charge.
Choose a central atom and arrange the other atoms around it.
Use the valence electrons to complete the octets of each of the surrounding
atoms except hydrogen, which only requires two electrons.
All of the valence electrons have been used, and each atom has an octet.
Complete the structure by replacing the bonding electrons with lines.
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Lewis Structure of SO3
1.
2.
3.
Sum the valence electrons from all atoms. Add electrons to account
for negative charge, subtract to account for postitive charge.
Choose a central atom and arrange the other atoms around it.
Use the valence electrons to complete the octets of each of the surrounding
atoms except hydrogen, which only requires two electrons.
4. The central atom is left with six electrons. To complete its octet, share a lone
pair from one of the surrounding atoms
5. All of the valence electrons have been used, and each atom has an octet.
Complete the structure by replacing the bonding electrons with lines.
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Lewis Structure of SO321.
2.
3.
Sum the valence electrons from all atoms. Add electrons to account
for negative charge, subtract to account for postitive charge.
Choose a central atom and arrange the other atoms around it.
Use the valence electrons to complete the octets of each of the surrounding
atoms except hydrogen, which only requires two electrons.
4. The central atom is left with six electrons. To complete its octet, add the last
lone pair to the central atom
5. All of the valence electrons have been used, and each atom has an octet.
Complete the structure by replacing the bonding electrons with lines.
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Lewis Structure of PH3
1.
2.
3.
Sum the valence electrons from all atoms. Add electrons to account
for negative charge, subtract to account for postitive charge.
Choose a central atom and arrange the other atoms around it.
Use the valence electrons to complete the octets of each of the surrounding
atoms except hydrogen, which only requires two electrons.
4. The central atom is left with six electrons. To complete its octet, add the last
lone pair to the central atom.
5. All of the valence electrons have been used, and Phosphorus has an octet
and athe hydrogen atoms have two each. Complete the structure by replacing
the bonding electrons with lines.
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Lewis Structure of CO2
1.
2.
3.
Sum the valence electrons from all atoms. Add electrons to account
for negative charge, subtract to account for postitive charge.
Choose a central atom and arrange the other atoms around it.
Use the valence electrons to complete the octets of each of the surrounding
atoms except hydrogen, which only requires two electrons.
4. The central atom is left with four electrons. To complete its octet, share lone
pairs from two of the surrounding atoms
5. All of the valence electrons have been used, and each atom has an octet.
Complete the structure by replacing the bonding electrons with lines.
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Resonance
When multiple bonds are present, the choice of the
surrounding atom that receives the extra bond can be
ambiguous. Consider SO3:
In sulfur trioxide, all bonds are equivalent in length and
strength, suggesting that the double bond is shared among all
three oxygens.
The three representations, which differ only by the placement
of the electrons, are called resonance structures, as indicated
by the double arrows.
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Resonance
Which of the following exhibit resonance?
Cl
Cl
H
C
C
Cl
H
Cl
carbon
tetrachloride
O
formaldehyde
(CH2O)
O
O -- O
O
N
N
O
C
O
O
O
nitrate anion
carbon
dioxide
Resonance requires:
1) At least one double bond
2) At least two surrounding atoms that can accommodate a
double bond.
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Exceptions to the Octet Rule
There are three cases that produce exceptions to the octet (8)
rule:
1) The central atom has less than 8 electrons. Occurs with
Be (4), B (6) and Al (6).
2) The compound has an odd number of valence electrons
3) The central atom has more than 8 electrons. Can only
occur with row three (n = 3) and higher elements.
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Bond Energies
Bond energy is the minimum energy required to
break a bond.
Bond Energy for N2
In general, the shorter the bond, the stronger the
bond.
Bond
C-F
C-Cl
C-Br
H (kJ/mol)
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Bond
C C
C C
C C
H (kJ/mol)
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Bond Energies
Bond breaking is endothermic; bond making is exothermic.
By combining the energy absorbed by breaking reactant
bonds with the energy released by forming product bonds,
the enthalpies of gas phase reactions can be estimated.
Estimate the enthalpy change for the reaction of hydrogen
gas and bromine gas to form gaseous HBr
H H + Br Br
Bond
2 H-Br
H (kJ/mol)
Break reactant bonds:
Make product bonds:
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Bond Energies
Estimate the enthalpy change for the following reaction.
H
H
H
H3C
C
C
CH 3
C
H2
H3C
+ H-Br
C
C
Br
H
CH 3
C
H2
H
Bond
H (kJ/mol)
Break reactant bonds:
Make product bonds:
CH3CHCHCH2CH3 + HBr
CH3CH2CHBrCH2CH3
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