Basic Concepts of
Chemical Bonding
Chapter 8
Chapter 8
1
Lewis Symbols and the Octet Rule
There are three types of chemical bonds
(intramolecular force)
- Ionic Bond - electrostatic attraction between ions of
opposite charge (NaCl)
- Covalent Bond - sharing of electrons between two
atoms (Cl2)
- Metallic Bond - sharing of electrons between several
atoms (Ag).
Chapter 8
2
Lewis Symbols and the Octet Rule
- The electrons involved in bonding are called valence
electrons.
- Valence electrons are found in the incomplete,
outermost orbital shell of an atom.
- We can represent the electrons as dots around the
symbol for the element.
- These pictorial representations are called Lewis
Structures or Lewis Dot Structures.
Chapter 8
3
Lewis Symbols and the Octet Rule
Octet rule – Atoms tend to gain, lose or share electrons
until they are surrounded by eight valence electrons.
Chapter 8
4
Ionic Bonding
Consider the reaction between sodium and chlorine:
Na(s) + ½Cl2(g)  NaCl(s)
Chapter 8
5
Ionic Bonding
Consider the reaction between sodium and chlorine:
Na(s) + ½Cl2(g)  NaCl(s)
Chapter 8
6
Ionic Bonding
Na(s) + ½Cl2(g)  NaCl(s)
-
DH°f = -410.9 kJ
This reaction is very exothermic
Sodium loses an electron to become Na+
Chlorine gains an electron to become ClNa+ has an Ne electron configuration and Cl- has an
Ar configuration
Chapter 8
7
Ionic Bonding
Energetics of Ionic Bond Formation
Lattice Energy (DHlattice) – The energy required to
completely separate one mole of a solid ionic compound
into its gaseous ions.
Lattice energy depends on:
-the charge on the ions
-the size of the ions
Coulomb’s equation:
Chapter 8
8
Ionic Bonding
Energetics of Ionic Bond Formation
Lattice Energy (DHlattice) – The energy required to
completely separate one mole of a solid ionic compound
into its gaseous ions.
Lattice energy depends on:
-the charge on the ions
-the size of the ions
Coulomb’s equation:
Q1Q2
Ek
d
Q1, Q2 = charge on ions
k = 8.99 x 109 J-m/c2
d = distance between ions
Chapter 8
9
Ionic Bonding
Electron Configurations of Ions of the
Representative Elements
- common ions have the same electron configuration as
the nearest noble gas.
Mg : [Ne]3s2
Mg+: [Ne]3s1 not stable
Mg2+: [Ne] stable.
Cl : [Ne]3s23p5
Cl-: [Ne]3s23p6 = [Ar] =stable.
Chapter 8
10
Ionic Bonding
Transition-Metal Ions
- Common charges, +1, +2, +3
- In general, electrons are removed from orbitals in
order of increasing n (i.e. electrons are removed from
4s before the 3d).
Chapter 8
11
Ionic Bonding
Polyatomic Ions
- Covalently bound molecule with a charge
(SO42-, NO3-)
Chapter 8
12
Sizes of Ions
- Cations are smaller than their parent atoms.
- Anions are larger than their parent atoms.
- For ions of the same charge, ion size increases down a
group.
Chapter 8
13
Sizes of Ions
Isoelectric Series – a series of atoms or ions with the
same number of electrons.
O2- > F- > Na+ > Mg2+ > Al3+
As nuclear charge increases in an isoelectronic series,
the ions become smaller.
Chapter 8
14
Sizes of Ions
Chapter 8
15
Covalent Bonding
When similar atoms bond, they share pairs of electrons
to each obtain an octet.
Chapter 8
16
Covalent Bonding
When similar atoms bond, they share pairs of electrons
to each obtain an octet.
Example
Cl
+
Cl
Cl
Cl
has electrons connecting the two nuclei.
Chapter 8
17
Covalent Bonding
Multiple Bonds
- It is possible for more than one pair of electrons to be
shared between two atoms (multiple bonds)
-One shared pair of electrons - single bond (H2)
-Two shared pairs of electrons - double bond (O2)
-Three shared pairs of electrons - triple bond (N2).
Chapter 8
18
Covalent Bonding
Multiple Bonds
- It is possible for more than one pair of electrons to be
shared between two atoms (multiple bonds)
-One shared pair of electrons - single bond (H2)
-Two shared pairs of electrons - double bond (O2)
-Three shared pairs of electrons - triple bond (N2).
H H
O O
Chapter 8
N N
19
Covalent Bonding
Multiple Bonds
- It is possible for more than one pair of electrons to be
shared between two atoms (multiple bonds)
-One shared pair of electrons - single bond (H2)
-Two shared pairs of electrons - double bond (O2)
-Three shared pairs of electrons - triple bond (N2).
H H
O O
N N
- Generally, bond distances decrease as we move from
single through double to triple bonds.
Chapter 8
20
Bond Polarity and Electronegativity
- Electrons in a covalent bond may not be shared
evenly.
Electronegativity – The ability of an atom in a molecule
to attract electrons to itself.
- The periodic trend for electronegativity is up and to
the right across the periodic table.
Chapter 8
21
Bond Polarity and Electronegativity
Electronegativity
Chapter 8
22
Bond Polarity and Electronegativity
Electronegativity and Bond Polarity
- A chemical bond between elements with large
differences in eletronegativity will shift the electrons
to the atom with the higher electronegativity.
- The positive end (or pole) in a polar bond is
represented + and the negative pole -.
- This is called a polar covalent bond.
- If the electronegativity difference is small, the bond is
nonpolar; if it is large, it is a polar bond.
Chapter 8
23
Drawing Lewis Structures
1) Draw a skeleton structure of the molecule or ion
showing the arrangement of the atoms and then
connect each atom to another with a single bond.
2) Determine the total number of valence electrons in
the molecule or ion.
3) Deduct 2 electrons for each single bond used in step
1.
4) Distribute the rest of the electrons so that each atom
(except H) has 8 electrons.
-
If you are “short” electrons, form multiple bonds
If you have “extra” electrons, one of the heavy atoms may
be able to hold more than eight electrons.
Chapter 8
24
Drawing Lewis Structures
PCl3
Chapter 8
25
Drawing Lewis Structures
PCl3
Cl
P Cl
Cl
Chapter 8
26
Drawing Lewis Structures
PCl3
Cl
P Cl
Cl
Element
Number
Electrons
Total
P
1
5
5
Cl
Total Electrons
Chapter 8
27
Drawing Lewis Structures
PCl3
Cl
P Cl
Cl
Element
Number
Electrons
Total
P
1
5
5
Cl
3
7
21
Total Electrons
Chapter 8
26
28
Drawing Lewis Structures
PCl3
Cl
P Cl
Cl
Element
Number
Electrons
Total
P
1
5
5
Cl
3
7
21
Total Electrons
26
Electrons used
6
Electrons remaining
20
Chapter 8
29
Drawing Lewis Structures
PCl3
Cl
P Cl
Cl
Element
Number
Electrons
Total
P
1
5
5
Cl
3
7
21
Total Electrons
26
Electrons used
6
Electrons remaining
20
Cl
P Cl
Cl
Chapter 8
30
Drawing Lewis Structures
Formal Charge
Used to predict the correct Lewis Structure.
1) Half of the electrons in a bond are assigned to each atom
in the bond.
2) Both electrons of an unshared pair of electrons are
assigned to the atoms to which the unshared pair belong.
3) The formal charge of an atom is equal to the valence
electrons minus the number of electrons assigned to each
atom.
Formal Charge = (group number) – (assigned electrons)
4) The sum of the formal charges will equal the charge on
the molecule or polyatomic ion.
Chapter 8
31
Drawing Lewis Structures
Using Formal Charge
1) A Lewis structure in which all formal charges in a molecule
are equal to zero is preferable to one in which some formal
charges are not zero.
2) If a Lewis structure has non-zero formal charges, the one
with the fewest nonzero formal charges is preferred.
3) A Lewis structure with one large formal charge is preferable
to one with several small formal charges.
4) A Lewis structure with adjacent formal charges should have
opposite signs.
5) When choosing between several Lewis structures, the
structure with negative formal charges on the more
electronegative atom is preferable.
Chapter 8
32
Drawing Lewis Structures
Resonance Structures
- Some molecules are not well described by Lewis
Structures.
Chapter 8
33
Drawing Lewis Structures
Resonance Structures
- Some molecules are not well described by Lewis
Structures.
Example: Ozone
1
O
O
O3
1
Chapter 8
O
O
O3
34
Drawing Lewis Structures
Resonance Structures
- Experimentally, ozone has two identical bonds
whereas the Lewis Structure requires one single and
one double bond.
O
O
O
Chapter 8
35
Drawing Lewis Structures
Resonance Structures
- Resonance structures are attempts to represent a real
structure that is a mix between several extreme
possibilities.
- Each Lewis structure is call a Resonance Form
Resonance Form – Two or more Lewis structures having
the same arrangements of atoms but a different
arrangement of electrons
Chapter 8
36
Drawing Lewis Structures
Resonance Structures
- In ozone the resonance forms have one double and one
single bond.
Chapter 8
37
Drawing Lewis Structures
Resonance Structures
- In ozone the resonance forms have one double and one
single bond.
O
O
O
O
Chapter 8
O
O
38
Drawing Lewis Structures
Resonance Structures
- In ozone the resonance forms have one double and one
single bond.
O
O
O
O
O
O
- The actual structure of O3 is a combination (or
average) of the individual forms called a resonance
hybrid.
Chapter 8
39
Exceptions to the Octet Rule
There are three classes of exceptions to the octet rule:
- Molecules with an odd number of electrons
- Molecules in which one atom has less than an octet
- Molecules in which one atom has more than an octet
Chapter 8
40
Exceptions to the Octet Rule
Odd Number of Electrons
There are few molecules which fit this category
Examples. ClO2, NO, and NO2
Chapter 8
41
Exceptions to the Octet Rule
Odd Number of Electrons
There are few molecules which fit this category
Examples. ClO2, NO, and NO2
O N O
Chapter 8
42
Exceptions to the Octet Rule
Less than an Octet
- This refers to the central molecule
- Typical for compounds of Groups 1A, 2A, and 3A.
Examples: LiH, BeH2, BF3
Chapter 8
43
Exceptions to the Octet Rule
More than an Octet
- This starts for atoms in the 3rd period onwards.
- This is due to vacant d orbitals which can hold the
“extra” electrons.
- Another factor is the size of the central atom, as they
get bigger, it gets easier to place additional electrons
around the central atom.
Chapter 8
44
Strengths of Covalent Bonds
Bond Enthalpy (Energy) - The energy required to break
a covalent bond of a gaseous substance.
Cl2(g)  2Cl(g)
DH = DCl-Cl
- When more than one bond is broken the bond
enthalpy is a fraction of DH for the atomization
reaction :
CH4(g)  C(g) + 4H(g) DH = 1660 kJ
DC-H = ¼DH = ¼(1660 kJ) = 415 kJ.
- Bond enthalpies can either be positive or negative.
Chapter 8
45
Strengths of Covalent Bonds
Chapter 8
46
Strengths of Covalent Bonds
Bond Enthalpies and the Enthalpies of Reaction
- Bond enthalpies can be used to calculate DHrxn.
DHrxn = D(bonds broken) - D(bonds formed).
Chapter 8
47
Bond Enthalpies and the Enthalpies of Reaction
H
+
C C
H
H H
H
H O C C O H
H O O H
H
H H
Chapter 8
48
Bond Enthalpies and the Enthalpies of Reaction
H
+
C C
H
H H
H
H O C C O H
H O O H
H
H H
Bonds Broken
Bonds Formed
C=C
614 kJ/mol
C-C
348 kJ/mol
O-O
146 kJ/mol
C-O
358 kJ/mol
Chapter 8
49
Bond Enthalpies and the Enthalpies of Reaction
H
+
C C
H
H H
H
H O C C O H
H O O H
H
H H
Bonds Broken
Bonds Formed
C=C
614 kJ/mol
C-C
348 kJ/mol
O-O
146 kJ/mol
C-O
358 kJ/mol
DHrxn = [1mol(614kJ/mol)+1mol(146kJ/mol)]-[2mol(358kJ/mol)+1mol(348kJ/mol)]
Chapter 8
50
Bond Enthalpies and the Enthalpies of Reaction
H
+
C C
H
H H
H
H O C C O H
H O O H
H
H H
Bonds Broken
Bonds Formed
C=C
614 kJ/mol
C-C
348 kJ/mol
O-O
146 kJ/mol
C-O
358 kJ/mol
DHrxn = [1mol(614kJ/mol)+1mol(146kJ/mol)]-[2mol(358kJ/mol)+1mol(348kJ/mol)]
= -304 kJ
Chapter 8
51
Bond Enthalpies and the Enthalpies of Reaction
Cl
2
Cl
N
N
N
Cl
Chapter 8
+
Cl
3
Cl
52
Bond Enthalpies and the Enthalpies of Reaction
Cl
2
Cl
N
N
N
Cl
+
Cl
3
Cl
Bonds Broken
Bonds Formed
N-Cl
200 kJ/mol
N=N
941 kJ/mol
Cl-Cl
242 kJ/mol
Chapter 8
53
Bond Enthalpies and the Enthalpies of Reaction
Cl
2
Cl
N
N
N
Cl
+
Cl
3
Cl
Bonds Broken
Bonds Formed
N-Cl
200 kJ/mol
N=N
941 kJ/mol
Cl-Cl
242 kJ/mol
DHrxn = [2(3mol(200kJ/mol))]-[1mol(941kJ/mol)+3mol(242kJ/mol)]
Chapter 8
54
Bond Enthalpies and the Enthalpies of Reaction
Cl
2
Cl
N
N
N
Cl
+
Cl
3
Cl
Bonds Broken
Bonds Formed
N-Cl
200 kJ/mol
N=N
941 kJ/mol
Cl-Cl
242 kJ/mol
DHrxn = [2(3mol(200kJ/mol))]-[1mol(941kJ/mol)+3mol(242kJ/mol)]
= -467 kJ
Chapter 8
55
Oxidation Numbers (States)
Oxidation States in molecular compounds
1) The oxidation number of an element in its elemental form is
zero (N2, O2, P4)
2) The oxidation number of a monoatomic ion is the same as its
charge (Cl-, N3-, Na+)
3) In binary compounds, the element with the greater
electronegativity is assigned a negative oxidation number
equal to its charge in simple ionic compounds
4) The sum of the oxidation numbers equals zero for an
electrically neutral compound and equals the overall charge
for an ionic species.
Chapter 8
56
Oxidation Numbers (States)
Common oxidation numbers
Group IA, +1
Group IIA, +2
Group IIIA, +3
Fluoride in compounds(F) is always -1
Oxygen is usually -2
Peroxide (O22-) is -1
Superoxide (O2-) is -½
Hydrogen can be either +1 or -1
Chapter 8
57
Oxidation Numbers
Nomenclature
- Oxidation state are expressed in the names of
substances.
- For ionic compounds: FeF2, Fe is in the +2 oxidation
state and is called iron(II) fluoride.
Chapter 8
58
Homework
8.14, 8.30, 8.32, 8.44, 8.58, 8.62
Chapter 8
59