Chapter 9: Chemical Bonds
Types of Bonds
• Ionic
– Metal and nonmetal
– Electron transfer
– Infinite lattice
• Covalent
– Nonmetal and nonmetal
– Shared electrons
– Individual molecules
• Metallic
– Metal and metal
– “electron sea”
– Infinite lattice
Ionic Bonds: Octet Rule
In forming ionic compounds, atoms tend to gain or lose
electrons in order to achieve a stable valence shell electron
configuration of 8 electrons.
Group I metals -->
Group II metals -->
Al (group III) -->
Group VII (17) -->
Group VI (16) -->
Group V (15) -->
+1 cations (Li+, Na+, etc)
+2 cations (Mg2+, Ca2+, etc)
Al3+
-1 anions (F–, Cl–, Br–, etc)
-2 anions (O2–, S2–, etc)
-3 anions (N3–, P3–)
e.g. Na 2s22p63s1 --> Na+ 2s22p6 {~Ne}
Cl 3s23p5 --> Cl– 3s23p6 {~Ar}
Remember Ch. 2 Figure 2.14 (p62)
Ionic Bonds: Lewis Symbols
simple notation for showing number of valence electrons
Cl
O
group VII (7 valence e–)
group VI (6 valence e–)
3s23p5
2s22p4
••
•O•
••
••
••Cl •
••
e.g. use Lewis symbols to illustrate the formation of a compound of
sodium and sulfur
Na2S Na+ combined with S2–
Na
+
S
+
Na
2 Na
+
2-
S
Lattice Energy
Ionic Bonding;
electrostatic attraction of positive (cation) and negative
(anion) ions
Neutral atoms
e– transfer
(IE + EA)
Ionic compound
Lattice energy
cation + anion
(lattice energy)
energy released when gaseous ions combine
to form crystalline solid (an ionic compound)
e.g. lattice energy of NaCl is 787 kJ/mol:
Na+(g) + Cl–(g) --> NaCl(s) DH = –787 kJ
Bigger ions  smaller lattice energy
Higher charge  larger lattice energy
Born-Haber Cycles
• Lattice energy can
be calculated using
a Born-Haber
cycle; a
hypothetical series
of steps describing
the formation of an
ionic compound
from the elements.
Here, DH°lattice = DH°f –
(DH°step1 + DH°step2 +
DH°step3 + DH°step4)
Covalent Bonding
1. Covalent Bond Formation
– results from sharing of one or more pairs of electrons between
two atoms
– Examples:
H
+
H
H H
H
+
F
H F
or
or
H
H
H
F
2. Octet Rule -- for covalent bonding
– In forming covalent bonds, atoms tend to share sufficient
electrons so as to achieve a stable outer shell of 8 electrons
around both atoms in the bond.
Examples
H
C
+
4 H
H C H or
H
N
+
3 H
H
+
2 H
H
H C H
H
N H
H
O
H
O H
unshared e- pairs
or "lone pairs"
Multiple Bonds
• “double” and “triple” bonds
double bond
triple bond
sharing of 2 pairs of electrons between two atoms
sharing of 3 pairs of electrons between two atoms
bond energy/bond strength
Type of Bond
Bond order
single
1
double
2
bond distance
Examples: O2
N2
CO2
{O=O double bond}
{N≡N triple bond}
{two C=O double bonds}
triple
3
Electronegativity and Bond Polarity
electronegativity
tendency of an atom in a molecule to attract
electrons to itself
Electronegativity Increases
Periodic Table
e.g. Cl is more electronegative than H, so there is partial charge
separation in the H-Cl bond:
+ H Cl
•
•
•
or
H
Cl
The H-Cl bond is described as “polar” and is said to have a “dipole”
The entire HCl molecule is also polar as a result
More complex molecules can be polar or nopolar, depending on their 3-D
shape (Later)
Lewis Electron Dot Structures
General Procedure -- stepwise process
--Write the skeletal structure (which atoms are bonded?)
--H atoms terminal
--More electronegative atoms terminal (e.g. halogens)
--Count all valence electrons (in pairs) and charges
--Place 2 electrons (1 pair) in each bond
--Complete the octets of the terminal atoms
--Use multiple bonds if needed to complete the octet of the central
atom, or
--Put any remaining electron pairs on the central atom
--Show formal charges
Apply the OCTET RULE as follows:
--H never has more than 2 electrons (I.e. one bond)
--2nd row elements (e.g. C, N, O) almost always have an octet and
never have more than 8 electrons (sometimes boron has only 6)
--3rd row and higher elements can have more than 8 electrons but
only after the octets of any 2nd row elements are completed
Formal Charge
• The “apparent” charge on an atom in a covalent bond
• = (# of valence e– in the isolated atom) - (# of bonds to the
atom) - (# of unshared electrons on the atom)
minimize formal charges whenever possible
(but the octet rule takes priority!)
NH4+
CO
+
H
H
N
C
O
H
H
Put a circle around the formal charge.
All nonzero formal charges must be shown in a Lewis structure.
Lewis Electron Dot Structures
General Procedure -- stepwise process
--Write the skeletal structure (which atoms are bonded?)
--H atoms terminal
--More electronegative atoms terminal (e.g. halogens)
--Count all valence electrons (in pairs) and charges
--Place 2 electrons (1 pair) in each bond
--Complete the octets of the terminal atoms
--Use multiple bonds if needed to complete the octet of the central
atom, or
--Put any remaining electron pairs on the central atom
--Show formal charges and resonance forms as needed
Apply the OCTET RULE as follows:
--H never has more than 2 electrons (I.e. one bond)
--2nd row elements (e.g. C, N, O) almost always have an octet and
never have more than 8 electrons (sometimes boron has only 6)
--3rd row and higher elements can have more than 8 electrons but
only after the octets of any 2nd row elements are completed
Resonance
• When multiple bonds are present, a single Lewis structure
may not adequately describe the compound or ion -- occurs
whenever there is a “choice” of where to put a multiple
bond.
e.g. the HCO2– ion is a “resonance hybrid” of two
“contributing resonance structures”
The C-O bond order is about 1.5
(average of single and double bonds)
O
H
C
O
H
O
C
O
All reasonable resonance structures must be shown in a Lewis structure.
Lewis Dot Structure Checklist
•
•
•
•
•
•
Correct total # of valence electrons
Correct connectivity
NO atoms with > octet (or duet) in group 2 or below
All atoms have octets (or duets), if possible
Lone electrons clearly shown
Nonzero formal charges included (and circled)
– Note it should be “2+” not “+2”
– If the charge is 1+ or 1- do not write the “1”!
• Important resonance structures included
• Ions have brackets and overall charge (not circled)
Exceptions to the Octet Rule
• Odd-Electron Species
– “radicals” or “free radicals”
• Incomplete Octets
– Group 3 elements often have 6 electrons total, e.g. AlCl3
– Often form bonds to complete the octets
• Expanded octets
– Only 3rd row or higher
Sample Problems
• Write Lewis Electron Dot Structures (including formal
charges and/or resonance as needed) for the following
compounds and ions.
PF3
HCN
SF5–
NO2–
SOCl2
O3
HNO3
H2CO
N3–
Sample Problems
• Write Lewis Electron Dot Structures (including formal
charges and/or resonance as needed) for the following
compounds and ions.
H
PF3
C
N
HCN
SF5
NO2
–
N
–
N
N
F
F
P
F
F
F
SOCl2
N
N
F
F
O
O
F
S
O
O
O3
HNO3
H2CO
N3–
O
O
O
O
O
O
O
Cl
S
Cl
N
O
N
N
N
N
N
N
N
N
O
H C
O
O
N
N
H
2-
O
O
O
H
2-
H
Bond Energies and Estimating DH
• Bond Energy—energy required to break 1 mole of the bond
in the gas phase, e.g.
• HCl(g)  H(g) + Cl(g) DH = 431 kJ/mol
• CH4(g)  CH3(g) + H(g) DH = 438 kJ/mol
• Endothermic!
• Estimating DHrxn using bond energies;
DHrxn = ∑(DH bonds broken) – ∑(DH bonds formed)
Sample Problem
Ethanol is a possible fuel. Use average bond energies to calculate
DHrxn for the combustion of gaseous ethanol.
Sample Problem
Ethanol is a possible fuel. Use average bond energies to calculate
DHrxn for the combustion of gaseous ethanol.
CH3CH2OH(g) + 3 O2(g)  2 CO2(g) + 3 H2O(g)
DHrxn = -1245 kJ/mol