Chapter 9
Chemical Bonding
CHM 1045
Bushra Javed
1
Chemical Bond Concept
• Recall that an atom has core and valence
electrons.
• Core electrons are found close to the nucleus.
• Valence electrons are found in the most
distant s , p and partially filled d subshells.
• It is valence electrons that are responsible for
holding two or more atoms together in a
chemical bond.
2
Octet Rule
• The octet rule states that atoms bond in such
a way so that each atom acquires eight
electrons in its outer shell.
• There are two ways in which an atom may
achieve an octet.
(a) by transfer of electrons from one atom to
another
(b) by sharing one or more pairs of electrons
3
Types of Bonds
• Ionic bonds are formed
from the complete
transfer of electrons
between atoms to form
ionic compounds.
• Covalent bonds are
formed when two
atoms share electrons
to form molecular
compounds
4
Ionic Bonds
An ionic bond is formed by the attraction
between positively charged anions and
negatively charged anions.
• This “electrostatic attraction” is similar to the
attraction between opposite poles on two
magnets.
5
Ionic Bonding between Na and Cl
6
Ionic Bonds
• The ionic bonds formed
from the combination
of anions and cations
are very strong and
result in the formation
of a rigid, crystalline
structure. The structure
for NaCl, ordinary table
salt, is shown here.
7
Energies involved in Ionic Bonding
•
Ionic bonds have energies, which are released
when ionic bonds are formed
• In general, the energy of the ionic bond according to
Coulomb's Law is given as:
E=
qanion . qcation
r
where qanion and qcation are charges on the anion
and the cation,
• r is the distance between them
The greater the charges and the smaller the distance
between them, the stronger is this ionic bond.
8
Ionic Radii
• The radius of a cation is smaller than the
radius of its starting atom.
• The radius of an anion is larger than the radius
of its starting atom.
9
Ionic Radii
Example 1
Rank the following ions in order of decreasing ionic
radius: S2–, O2–, F–, Na+, Mg2+.
a) S2–, O2–, F–, Na+, Mg2+
b) O2–, F–, Na+, Mg2+, S2–
c) Mg2+, Na+, F–, O2–, S2–
d) Mg2+, S2–, Na+, F–, O2–
e) O2–, S2–, F–, Na+, Mg2+
10
Energies involved in Ionic Bonding
• The transfer of an electron from a sodium
atom to a chlorine atom is not in itself
energetically favorable; it requires 147 kJ/mol
of energy.
• However, 493 kJ of energy is released when
these oppositely charged ions come together.
• An additional 293 kJ of energy is released
when the ion pairs solidify.
11
Lattice Energy
• The energy necessary to separate ionic solid into
gaseous ions: NaCl (s) → Na+(g) + Cl-(g) ΔH = U
• This ‘lattice energy’ is the negative of the energy
released when gaseous ions form an ionic solid.
• For a given arrangement of ions, lattice energy
increases as the charges on the ions increase and as
their radii decrease.
12
Lattice Energy
Example 2
When the cations Na+, K+, Rb+, Cs+ are combined
with chloride ion in the gas phase to form ion
pairs, which pair formation releases the greatest
amount of energy?
a) NaCl
b) KCl
c) RbCl
d) CsCl
13
Born Haber Cycle: Energetics of Ionic Bond
Formation
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Energetics of Ionic Bond Formation
Example 3
In the Born–Haber cycle for NaCl(s), which of the
following processes corresponds to the electron affinity
of Cl?
a) Cl(g) → Cl+(g) + e–
b) NaCl(s) → Na+(g) + Cl–(g)
c) Cl2(g) → 2Cl(g)
d) Cl–(g) → Cl(g) + e–
e) Cl(g) + e– → Cl–(g)
15
Lattice Energy
Example 4
Which of the following compounds would be
expected to have the highest melting point?
a) LiF
b) LiCl
c) CsF
d) NaBr
16
Electron Dot Formulas
• An electron dot formula of an element shows
the symbol of the element surrounded by its
valence electrons.
• We use one dot for each
valence electron.
• Let’s draw EDF of phosphorus.
17
Formation of Cations
• We can use electron dot
formulas to look at the
formation of cations.
• Each of the metals in
Period 3 form cations by
losing 1, 2, or 3 electrons,
respectively. Each metal
atom becomes
isoelectronic with the
preceding noble gas,
neon.
18
Formation of Anions
• We can also use
electron dot formulas to
look at the formation of
anions.
• The nonmetals in Period
3 gain 1, 2, or 3
electrons, respectively,
to form anions. Each
nonmetal ion is
isoelectronic with the
following noble gas,
argon.
19
Electron Configuration of Ions
Example 5
The ground-state electron configuration of the Mg 2+
ion is
a) 1s22s22p6.
b) 1s22s22p63s2.
c) 1s22s22p63s23p2.
d) 1s22s22p3.
20
Electron Configuration of Ions
Example 6
What is the electron configuration for Cr2+?
a) [Ar]4s1 3d5
b) [Ar]4s2 3d2
c) [Ar]3d3
d) [Ar]4s2 3d7
e) [Ar]3d4
21
Electron Configuration of Ions
Example 7
Which of the following species is isoelectronic
with Xe?
a) Kr
b) Rb+
c) Tl3+
d) Se2–
e) Ba2+
22
Electron Configuration of Ions
Example 8
Which pair of species is isoelectronic?
a) Ne and Ar
b) Na+ and K+
c) K+ and Cl–
d) Li+ and Ne
23
Covalent Bonds
• Covalent bonds are formed when two nonmetal atoms
share electrons and the shared electrons in the covalent
bond belong to both atoms.
• In any covalent bond, the attractive energy between the
nuclei and electrons exceeds the repulsive energy arising
from nuclear-nuclear and electron-electron interactions
• For a covalent bond to be formed, e.g. between two
hydrogen atoms, the orbital overlap of the 1s orbitals
gives a molecular orbital with high electron density
between the nuclei.
24
Covalent Bonding
Example 9
Which of the following is the best explanation for a
covalent bond?
a) electrons simultaneously attracted by more than one
nucleus
b) the overlapping of two electron-filled orbitals having
different energies
c) the overlapping of unoccupied orbitals of two or
more atoms
d) a positive ion attracting negative ions
25
Covalent Bonding
26
Bond Length
• When a covalent bond is formed, the valence shells
of the two atoms overlap with each other.
• In HCl, the hydrogen 1s energy sublevel overlaps with
the chlorine 3p energy sublevel. The mixing of
sublevels draws the atoms closer together.
• The distance between the two atoms is smaller than
the sum of their atomic radii and is the bond length.
27
Bond Dissociation Energy
• Energy is released when two atoms form a covalent
bond:
• H(g) + Cl(g)  HCl(g) + heat
• Conversely, energy is needed to break a covalent
bond.
• The energy required to break a covalent bond is
referred to as the bond energy.
• To pull apart a H2 molecule into individual atoms,the
bond dissociation energy is 104 kcal/mol
28
Types of Covalent Bonds
Polar covalent bond:
• A polar covalent bond is a covalent bond in which
the bonding electrons between the atoms are shared
unequally due to the electronegativity difference.
Non polar covalent bonds:
• A non polar covalent bond is a covalent bond in
which the bonding electrons are shared equally
because the two atoms are alike.
29
Electronegativity (EN or X)
• Electronegativity - describes the ability of an
atom to attract a shared pair of electrons
• Scale devised by L. Pauling varies from 0 - 4.
• Can be related to ionization energy and
electron affinity of atoms,
X = ½(IE + EA)
30
Trend in Electronegativities
31
Electronegativity
EN decreases from the top to the bottom in a group
and increases from left to right along a period.
• ΔEN = 2 or > 2
• ΔEN = 0
• ΔEN > 0 and < 2
ionic bond
non polar covalent bond
polar covalent bond
32
Electronegativity
Example 10
Which of the following atoms is the most
electronegative?
a) B
b) Na
c) N
d) Cs
e) Al
33
Electronegativity
Example 11
The larger the difference in electronegativity between two
bonded atoms,
1. the more ionic the bond.
2. the more covalent the bond.
3. the more polar the bond.
a) 1 only
b) 2 only
c) 3 only
d) 1 and 3 only
e) 2 and 3 only
34
Valence Shell:
The outermost shell of an atom containing the valence electrons.
Core electrons:
The inaccessible electrons. Core electrons do not participate in the
chemical reactions.
Bonding electrons:
The valence electrons which are shared (in bonding orbitals) between
two nuclei.
Bonding pair:
Bond where two electrons are shared between two atoms.
Nonbonding electrons:
The valence electrons that are not involved in bonding & are localized
in individual atoms.
Hydrogen is most stable with the duet (helium electron configuration).
35
Exceptions to the Octet Rule
Not all the molecules follow the octet rule. Many elements beyond the
2nd row do not comply with this rule.
Single bond
When one pair of electrons is shared by two atoms.
Notation. Two dots or one line is used.
Multiple bonds (double bonds, triple bonds)
A covalent bond in which two pairs of electrons are shared by two
atoms is called a double bond while a covalent bond in which the three
pairs of electrons are shared between the two atoms is called a triple
bond.
36
CONSTRUCTING LEWIS STRUCTURES:
STEP 1
If the substance is ionic, treat each ion separately.
A compound is ionic if it contains a polyatomic anion and a metal or an
ammonium cation.
If there is a metal atom but no polyatomic anion, apply
electronegativity differences to identify whether or not the compound is
ionic.
STEP 2
Count the total number of valence electrons of all elements in a
substance.
Forthe anions, add electrons and for cations subtract electrons.
37
STEP 3
Assemble the bonding framework. Account for the electrons used in
the framework. Distribute the rest in Step 4.
While drawing the framework keep in mind the following points.
The order in which atoms are listed in the formula indicates the
bonding pattern.
Hydrogen atoms are always outer atoms.
Outer atoms other than hydrogen are the most electronegative.
Atoms enclosed in parentheses are bonded together and the entire
group can be bound to another atom.
38
STEP 4
• Place three nonbonding pairs of electrons on
each outer atom except H.
• An outer atom other than hydrogen is more
stable when it is associated with an octet of
electrons.
39
STEP 5
Assign the remaining valence electrons to inner atoms.
If the molecule has more than one inner atom, place nonbonding pairs
around the most electronegative atom until it has an octet.
If there are still unassigned electrons, do the same for the next most
electronegative atom. Continue in this manner until all the electrons
have been assigned.
Place any remaining electrons on an inner atom that has n>2. Atoms
with n>2 have valence dorbitals that allow them to accommodate more
than 8 electrons.
•Compounds with P, S and Cl may have between 8 & 12
electrons associated with these atoms.
40
STEP 6
Satisfy electron configurations of the inner atoms.
In case an inner atom lacks an octet, a more stable distribution can be
achieved by moving some of the electrons from outer atoms to make
double or triple bonds to the inner atoms.
For inner atoms beyond the row 2, find the formal charge. If it leaves a
+ ve charge on the atom, shift electrons to form double bonds, even if
this gives the inner atom more than eight electrons. Amore stable
structure is achieved this way.
41
STEP 7
Identify equivalent Lewis structures.
Draw resonance structures when there is more than one way
to shift electron pairs.
A formal charge calculation may be necessary if more than
one Lewis dot formulas are possible.
A structure with Formal charge gives the approximate
distribution of electrons.
Formal charge = valence electrons on free atom – ½
(number of bonding electrons in a bond) – (number of lone
pair electrons)
To gets an accurate Lewis structure, minimize formal
charges in the Lewis structure.
42
Execeptions to Octet Rule
Example 12
In which of the following molecules is the octet rule
violated?
a) PF3
b) SiF4
c) OF2
d) ClF3
e) ClF
43
Execeptions to Octet Rule
Example 13
Which of the following species represents an exception
to the octet rule?
a) CH3OH
b) CCl4
c) PH3
d) BF3
44
Execeptions to Octet Rule
Example 14
Which of the following species represents an
exception to the octet rule?
• a) CH3OH
• b) CCl4
• c) PH3
• d) BF3
45
Execeptions to Octet Rule
Example 15
Which of the following species represents an exception
to the octet rule?
a) CO2
b) SF4
c) SiO2
d) PCl3
46
Resonance in organic structures
when a single Lewis structure cannot explain
the properties of a molecule.
In such cases we draw more than one structure
to represent the molecule and take the original
structure of the molecule as the weighed
average of these structures.
47
Resonance
48
Resonance
Example 16
The concept of resonance describes molecular
structures
a) that have several different geometric arrangements.
b) that have no suitable single Lewis formula.
c) that have electrons resonating.
d) that are formed from hybridized orbitals
49
Resonance
Example 17
All the following statements about resonance are true
except
a) A single Lewis formula does not provide an adequate
representation of the bonding.
b) Resonance describes the oscillation and vibration of
electrons.
c) Resonance describes a more stable situation than
does any one contributing resonance formula.
d) Resonance describes the bonding as intermediate
between the contributing resonance formulas.
e) The contributing resonance formulas differ only in
the arrangement of the electrons.
50
Formal Charge
Example 18
What is the formal charge on the chlorine atom in the
chlorate ion, ClO3–, in the Lewis dot formula that
minimizes formal charge?
a) +2
b) +1
c) 0
d) –1
e) –2
51
Bond energies
Example 19
Based on the following data, what is the I-I bond energy?
½H2(g) + ½I2(g) → HI(g); ∆H = 26.36 kJ
Bond
Bond Energy (kJ/mol)
H-H
435
H-I
295
a)
b)
c)
d)
e)
365 kJ/mol
155 kJ/mol
–155 kJ/mol
208 kJ/mol
–208 kJ/mol
52