CH 9: IONIC AND COVALENT
BONDING
Vanessa N. Prasad-Permaul
Valencia Community College
CHM 1045
1
Ionic Bonds
 Ionic Bonds: a chemical bond formed by the
electrostatic attraction between positive and
negative ions
 Bond forms when one or more electrons are
transferred from the valence shell of one atom to
the valence shell of the other.
 Cation: atom losing electrons
 Anion: atom gaining electrons
2
Ionic Bonds
Na + Cl
[Ne]3s1 + [Ne]3s23p5
Na+ + Cl-
[Ne] + [Ne]3s23p6
3
Ionic Bonds
 Lewis Electron-Dot Symbols: electrons in the
valence shell of an atom or ion are represented by
dots placed around the elemental symbol.
Na + Cl
Na+ + [ Cl ] -
 Valence electrons are those electrons with the
highest principal quantum number (n).
4
Ionic Bonds
Lewis Electron-Dot Symbols for Atoms in the
Periodic Table of Elements
5
Ionic Bonds
EXAMPLE 9.1: Use Lewis Electron-Dot symbols to
represent the transfer of electrons form magnesium
to fluorine atoms to form ions with noble gas
configuration.
Mg + Fl
F
+ Mg
+
F
MgFl2
[ F ]- + Mg2+ + [ F ]-
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Ionic Bonds
EXERCISE 9.1: Use Lewis Electron-Dot symbols to
represent the transfer of electrons form magnesium
to oxygen atoms to form ions with noble gas
configuration.
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Electron Configurations of Ions
Common monatomic ions found in compounds of the
main-group elements fall into 3 categories:
1. Cations of Groups IA to IIIA having noble gas or
pseudo-noble-gas configurations ; the ion charges
equal group numbers.
2. Cations of Groups IIIA to VA having the ns2 electrons
configurations; the ion charges equal the group
numbers minus two. Tl+, Sn2+, Pb2+, Bi3+.
3. Anions of Groups VA to VIIA having noble gas
configurations; the ion charges equal the group
number minus 8.
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Electron Configurations of Ions
EXAMPLE 9.2: Write the electron configuration and
Lewis symbol for N3-.
N = [He]2s22p3 + 3e-
[He]2s22p6
N= N
N3- = [ N ]3-
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Electron Configurations of Ions
EXERCISE 9.2: Write the electron configuration and
Lewis symbol for Ca2+ and S2-.
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Electron Configurations of Ions
EXERCISE 9.3: Write the electron configuration and
Lewis symbol for Pb and Pb2+.
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Electron Configurations of Ions
EXAMPLE 9.3: Write the electron configuration and
Lewis symbol for Fe2+ and Fe3+.
Fe = [Ar]3d64s2
=
Fe
Fe2+ = [Ar]3d6 =
[ Fe ]2+
Fe3+ = [Ar]3d5 =
[ Fe ]3+
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Electron Configurations of Ions
EXERCISE 9.4: Write the electron configuration and Lewis
symbol for Mn and Mn2+.
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Ionic Radii
Ionic Radius: A measure of the size of the spherical
region around the nucleus of an ion within the electrons
are most likely to be found.
 If the atom loses an electron, the cation will be
smaller.
 The electron-electron repulsion is initially less.
 orbitals can shrink to increase the
attraction of the electrons for the nucleus.
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Ionic Radii
EXERCISE 9.5: Which has the larger radius. S or S2-.
Explain.
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Ionic Radii
EXERCISE 9.6: Using only the Periodic Table, arrange the
following ions in order of increasing ionic radius : Sr2 ,
Mg2+, Ca2+ .
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Ionic Radii
The Comparison of Atomic and Ionic Radii
ATOMIC
RADIUS
DECREASES ACROSS A PERIOD
I
N
C
R
E
A
S
E
S
D
O
W
N
A
G
R
O
U
P
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Ionic Radii
Isoelectronic: Different species having the same number
and configuration of electrons.
Na+ < Mg2+ < Al3+
Decrease in atomic radius
 As the charge increases, the orbitals contract due to
the greater attractive forces of the nucleus  the
ionic radius decreases with increasing atomic
number.


In general , across a period the cations
decrease in radius.
As the anions are reached, there is an abrupt
increase in radius and then the radius
decreases again.
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Ionic Radii
EXAMPLE 9.4: Arrange the following ions in order of
decreasing ionic radius: F-, Mg2+, O2-.
F- = 1s22s22p6
Mg2+ = 1s22s22p6
O2- = 1s22s22p6
All are isoelectronic, so as the nuclear charge increases,
the ionic radius decreases.
O2- , F- , Mg2+
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Ionic Radii
EXERCISE 9.7: Arrange the following ions in order of
increasing ionic radius: Cl-, Ca2+, P3-.
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Covalent Bonds
 Covalent Bonds are formed by sharing at least
one pair of electrons.
 The attraction (nucleus/electrons) outweighs the
repulsions (electron/electron & nucleus/nucleus)
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Covalent Bonds
Every covalent bond has a characteristic length that
leads to maximum stability: Bond Length
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Strength of Covalent Bonds
Energy required to break a covalent bond in an
isolated gaseous molecule is called the bond
dissociation energy.
Same amount of energy released when the
bond forms.
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Electron-Dot Structures
 The electron-dot structures provide a simple,
but useful, way of representing chemical
reactions.
 Ionic:
 Covalent:
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Covalent Bonds
Coordinate covalent bond: a bond formed when both
electrons of the bond are donated by one atom
A
+
B
A B
A
+
B
A B
H+ +
NH3
H
+
H N H
H
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Covalent Bonds
OCTET RULE
 Group 1A tends to lose their ns1 valence shell




electron to adopt a noble gas electron
configuration.
Group 2A lose both ns2
Group 3A lose all three ns2 np1
Group 7A Gains one electron to attain noble
gas configuration
Group 8A inert, rarely lose or gain electrons
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Covalent Bonds
• Single Bonds:
C
• Double Bonds:
H
H
H
H
H
H
C
H
H
H
H
C
C
C
C
H
H
• Triple Bonds:
C
C
H
C
C
H
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Covalent Bonds
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Covalent Bonds
ELECTRONEGATIVITY: a measure of the ability of
an atom in a molecule to draw bonding electrons to itself.
 Bond polarity is due to electronegativity differences
between atoms.
 Pauling Electronegativity: is expressed on a scale
where F = 4.0
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Covalent Bonds
Pauling Electronegativities
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Covalent Bonds
EXAMPLE 9.5: Use electronegativity values to arrange
the following bonds in order of increasing polarity:
P H, H O, C Cl.
P
H = 0.0
H
O = 1.4
C
Cl = 0.5
P
H
, C
Cl ,
H
O
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Covalent Bonds
EXERCISE 9.8: Use electronegativity values to arrange
the following bonds in order of increasing polarity:
C O, C S, H Br.
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Drawing Lewis-Dot Structures
Rule 1: Count the total valence electrons.
Rule 2: Draw the structure using single bonds.
Rule 3: Distribute the remaining electron pairs
around the peripheral atoms.
Rule 4: Put remaining pairs on central atom.
Rule 5: Share lone pairs between bonded atoms to
create multiple bonds.
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Drawing Lewis-Dot Structures
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Lewis-Dot Structures
 NH2F Amino Fluoride: In this molecule,
nitrogen is the central atom.
 Rule 1: Number of electrons = 5 + (2 x 1) + 7 = 14 =
7 pairs
H N H
H N H
H N H
F
F
F
Rule 2
Rule 3
Rule 4
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Lewis-Dot Structures
EXAMPLE 9.6: Sulfur dichloride SCl2, is a red fuming
liquid used in the manufacture insecticides. Write the
Lewis formula for the molecule.
S = 6 Cl = 7 each for a total of 20 electrons
¨ : ¨S : Cl:
¨
:Cl
¨ ¨ ¨
¨
:Cl¨ S¨ Cl:
¨ ¨ ¨
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Lewis-Dot Structures
EXERCISE 9.9: Dichlorodifluoromethane CCl2F2, is a gas
used as a refrigerant and aerosol propellant. Write the
Lewis formula for the molecule.
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Lewis-Dot Structures
EXAMPLE 9.7: Carbonyl chloride or phosgene, COCl2, is a
highly toxic gas used as a starting material for the
preparation of polyurethane plastics. What is the electron
dot structure of this compound?
C = 4, O = 6, Cl =7 each for a total of 24 electrons
¨ : C : Cl
¨ :
: Cl
¨ :O:
¨ ¨
¨
¨:
: ¨Cl C Cl
¨
¨
:O:
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Lewis-Dot Structures
EXERCISE 9.10: Write the electron-dot structure of
carbon dioxide.
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Lewis-Dot Structures
EXAMPLE 9.8: Obtain the electron-dot formula for the
BF4- ion.
B = 3, F = 7 each (7 x 4) = 28 for a total 31 electrons. It is
an ion with one more electron so a total of 32 electrons.
: F¨ :
¨
¨
¨
:F: B: F:
¨ : F¨ : ¨
¨
: F¨ :
¨
:F
¨
B
:F:
¨
¨F :
¨
40
Lewis-Dot Structures
EXERCISE 9.11: Write the electron-dot structure of:
A. the hydronium ion, H3O+
B. The chlorite ion, ClO2-
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Resonance Structures
 When multiple structures can be drawn, the
actual structure is an average of all possibilities.
 The average is called a resonance hybrid. A
straight double-headed arrow indicates
resonance.
O
O
O
O
O
O
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Lewis-Dot Structures
EXAMPLE 9.9: Describe the electron structure of the
carbonate ion CO32-, in terms of electron-dot formulas.
C = 4, O = 3 x 6 = 18 for a total of 22 electrons, but it has
gained two electrons so there is a total of 24 electrons.
2:O:
¨:
:O
C
C
:O: :O:
¨
¨
:O
¨
2-
2-
¨:
:O
C
:O:
¨
:O:
¨
O:
¨
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Lewis-Dot Structures
EXERCISE 9.12: Describe the bonding in NO3- using
resonance formulas.
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Formal Charge
 Formal Charge: Determines the best resonance
structure.
 We determine formal charge and estimate the more
accurate representation.
Formal Charge = valence e- - # of e- in a bond - (# of lone-pair e-)
2
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Formal Charge
¨
:Cl
¨
¨
C Cl:
¨
:O:
¨
:Cl
¨
¨
C Cl
¨
:O:
¨-
+
¨
Cl
¨
¨
C Cl:
¨
:O:
¨-
Cl = 7 – (2/2) – 6 = 0
O = 6 – (4/2) – 4 = 0
C = 4 – (8/2) – 0 = 0
+
Cl = 7 – (4/2) – 4 = +1
O = 6 – (2/2) – 6 = -1
C = 4 – (8/2) – 0 = 0
Cl = 7 – (4/2) – 4 = +1
O = 6 – (2/2) – 6 = -1
C = 4 – (8/2) – 0 = 0
46
Formal Charge
EXAMPLE 9.11: Write the Lewis formula that best
describes the charge distribution in the sulfuric acid
molecule, H2SO4, according to the rules of formal charge.
¨
:O:
H
¨
O
¨
+2
S
:O:
¨
O
¨
O
¨
H
H
Ö
¨
S
O¨
¨
H
O
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Exercise 9.15: Write the Lewis formula that best
describes the phosphoric acid
molecule, H3PO4.
48
Resonance Structures
 How is the double bond formed in O3?
Move lone pair from
this oxygen?
O
O
O
O
O
or
O
Or from this
oxygen?
O
O
O
 The correct answer is that both are correct,
but neither is correct by itself.
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Example 1:
Which of the following is correct?
1. Energy is absorbed to form a bond
2. Energy is released when a bond is formed
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Example 2: Drawing Lewis-Dot Structures
Draw electron-dot structures for:
C3H8
H2O2
CO2
N2H4
CH5N
C2H4
C2H2
Cl2CO
H 3S +
HCO3–
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Example 3: Formal Charge
 Calculate the formal charge and determine
the most favorable of the following electron
dot structures:
SO2 NO3–
NCO– N2O O3
CO32–
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Example 4:
What is the overall formal charge of the
following structure?
1. -2
2. -3
3. -1
4. 0
O
O P O
O
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Example 5: Ionic Radii of Ions
 Compare ionic radii
 Fe & Fe3+
 Cl & Cl-
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