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Chapter 7
Chemical Bonding and
Molecular Structure
Three Types of Chemical Bonding
(1) Ionic: formed by electron transfer
(2) Covalent: formed by electron sharing
(3) Metallic: attraction between metal ions and their delocalized electrons
Ionic Bonding
Metal to Nonmetal; Metal to Polyatomic Ions
Examples: NaCl, MgCl2, NaNO3, NH4NO3
Binary Compounds:
Metals lose electrons and nonmetals gain electrons
Metal ions are called cations (positively charged)
Nonmetal ions are called anions (negatively charged)
Cations are attracted to anions and form crystal lattices called ionic compounds
Example:
An atom of Na loses one electron to form Na+
An atom of Cl gains one electron to form Cl-
Na+ is attracted to Cl- to form neutral NaCl
Remember all compounds are neutral
Ionic Bonding Question
Describe how calcium fluoride would form
from atoms of calcium and fluorine.
Lewis Electron-dot Symbols
The Lewis dot symbol of an atom depicts the number of “s” and “p” electrons in the outer energy level.
The number of dots around the symbol of an element ranges from 1-8 and reflects the valence electrons for that particle.
Generally dots are placed around the element’s symbol one at a time starting at the 9 o’clock position and continuing clockwise with no more than two dots shown at 9-12-3-6 o’clock.
Practicing Electron-Dot Symbols
Write the Lewis-dot symbols for:
Li Be B C N O F Ne
What is the Lewis-dot symbol for:
K S Al
The Octet Rule
When atoms form bonds they lose, gain, or
share electrons to attain a filled outer shell
(either two or eight electrons).
Use of Orbital Diagrams and
Electron Configurations in Bonding
Example: Lithium Fluoride
Another Example of an Ionic
Compound
Use electron-dot symbols to show how
aluminum oxide forms.
The Ionic Bonding Model
Energy considerations in ionic bonding
Overall, energy is released when ions come together in the formation of a compound
However, the process of the formation of an ionic compound involves a number of steps
Consider the simplified formation of the ionic
compound NaCl.
Let’s start with atoms of sodium and
chlorine.
A sodium atom has to lose one electron and
a chlorine atom has to gain one electron.
To remove an electron from a sodium atom it
requires energy (ionization energy)
Na(atom) + energy Na+(ion) + e-
When an atom of chlorine gains an electron it
releases energy (electron affinity)
Cl(atom) + e- Cl-(ion) + energy
When a Na+ and a Cl- ion come together into
a solid crystal a great amount of energy is
released (lattice energy).
Na+ + Cl- NaCl + energy
Lattice Energy
Lattice Energy is the energy absorbed that occurs when an ionic solid is separatedinto isolated ions in the gas phase. (Delta H is positive)
Also, the energy released when gaseous ions come together to form the crystal. (Delta H is negative)
Lattice energies for alkali metal-halogen
compounds
Summary
Note that in comparing compounds of like
charges the smaller the ion, the greater
the lattice energy.
BUT …
Lattice Energy is proportional to:
| charge M+ X charge NM-|
distance between nuclei
Which has the larger lattice energy?
KF or LiF
KF or CaF2
AlCl3 or AlBr3 or Al2S3
Note: Magnitude of charge dominates over size.
Properties of Ionic Compounds
Rigid, fixed positions of ions in solid state
Hard, brittle
Conduct electricity when melted or dissolved in water
High melting (mp) and boiling points (bp).
Note: The larger the lattice energy, the higher the mp and bp
Covalent Bonding
Nonmetal to Nonmetal Attraction: Shared
electrons
Example
H. .H H-H
A single covalent bond is formed by 2 atoms
sharing 2 electrons. Each atom has ownership
of the 2 electrons.
Covalent Bond Formation
Problem Solving
Using Lewis Dot Structures, explain how F2
forms.
F F F-F
How many bonding pairs are present?
How many lone pairs or nonbonding pairs are present?
Sharing of Bonding Pairs
Each atom in a covalent bond “counts” the
shared electrons as though they belong
entirely to that atom.
Let’s return to H2 and F2 and discuss the
meaning of this.
Another example: hydrogen fluoride
Bond Energy and Bond Length
Bond Energy: The amount of energy
required to break a bond (endothermic).
It is also the amount of energy released
in bond formation (exothermic)
Bond Length: The distance between the
nuclei of the two bonded atoms.
Examples of
Bond Strength/Bond Energy
C O C O C O
1070 kJ/mol 745 kJ/mol 358 kJ/mol
Note: Larger atoms result in longer,
thus weaker bonds.
Bond Energy and Chemical
Change
Electronegativity (EN)
Electronegativity is the ability of a bonded
atom to attract the shared electrons.
Change in Electronegativity ( EN) is the
difference in electronegativity of the two
bonded atoms.
Ionic, Polar Covalent, and Nonpolar
Covalent Bonding
Linus Pauling’s Work
Problem Solving
Consider the C-O bond, what is the EN for this
bond?
Consider the Br-Br bond, what is the EN for this
bond?
Which of the bonds above is/are polar covalent?
…nonpolar covalent?
Illustrating Bond Polarity
Conventional Methods for Illustrating the
Polarity of a bond.
C-O N-H C-H S-O H-F
Linus Pauling’s Work
Lewis Dot Structures Review
Draw Lewis Dot Structures of:
SiO2
SO32-
HNO2
Place your answer(s) on the board.
Exceptions to the Octet Rule
Examples:
Deficient Octet
BF3
Expanded Octet
H2SO4
Radicals
NO
Resonance Structures
Consider the following molecules:
CO32-
C6H6
CH3COO-
Note: Atoms do not change position! Only
pi and lone pair electrons.
Formal Charge
Determination of formal charge for each atom
within a molecule:
V.E. – (NB.E. + ½ B.E.) = formal charge
What are the formal charges of each atom in HCN?
What are the formal charges of each atom in NO3- ?
What are the formal charges of each atom in SO42- ?
Selecting the Preferred
Lewis Dot Structure
Use the concept of formal charge to select
the preferred structure of the CON- ion.
Valence Bond Theory
Central Themes:
• A covalent bond forms when orbitals of two atoms overlap and the overlap region is occupied by two electrons.
• The greater the overlap the stronger the bond.
• The stronger the bond the more stable the bond.
• Orbitals must become oriented so as to obtain the greatest overlap possible.
Types of orbital overlap:
sigma (end-to-end)
pi (side-by-side)
Note: sigma bond side-by-side overlap is
not permitted.
Let’s Consider CH4
How can carbon form four bonds?
What is the shape of methane?
What are the bond angles?
Hybrid Orbital Theory
Theory depicts a mix of the atomic orbitals (hybridize) to form the necessary number of orbitals needed for bonding.
The number of hybrid orbitals formed equals the number of atomic orbitals mixed.
The type of hybrid orbitals obtained varies with the type of atomic orbitals mixed.
Types of Hybrid Orbitals
sp
sp2
sp3
sp3d
sp3d2
sp Hybrids
Consider BeH2
Central Be needs two hybrid orbitals to
accommodate the two bonded H atoms.
To aid the visualization of sp hybrid
formation use an orbital diagram, shapes
of hybrids, and bond angles.
Bond angles for sp hybrids = 180o
Note: The bonding between the sp hybrid
orbital of beryllium and the “s” orbital of the
hydrogen atom is considered a sigma
bond overlap.
sp2 Hybrids
Consider BF3
Central boron atom needs three hybrid
orbitals to accommodate three bonded
fluorine atoms.
Visual of hybrid orbital formation theory.
Bond angles for sp2 hybrids = 120o
Note: The bond involving the sp2 hybridized
orbital of boron is an end-to-end overlap
(sigma bond) with the “p” orbital of the
fluorine atom.
sp3 Hybrids
Back to CH4
The central carbon atom needs four hybrid
orbitals.
Bond angles between sp3 hybrids = 109.5o
Other atoms using sp3 Hybrids
If the central atom has only two or three
(rather than four) bonded atoms, hybrid
orbitals may also contain the lone pairs of
electrons.
Examples: NH3 , H2O
sp3d Hybrids
Consider PCl5
The central phosphorus atom needs five hybrid
orbitals to accommodate the five bonded
chlorine atoms.
Note that since there is only one “s” orbital and
only three “p” orbitals available per energy
level, a “d” orbital must be used in the hybrid.
Bond angles between sp3d hybrids = 90o (axial) and
120o (equatorial)
sp3d2 Hybrids
Consider SF6
The six bonded atoms require six hybrids.
Bond angles between sp3d2 hybrids = 90o
Summary of Hybrid Theory
Hybrid Quiz
Predict the type of hybrid orbital you would
expect in the central atom of the following
molecules:
SiH4 BH3 AsF5 AlCl3 SF4
Multiple Bonds
Double Bond: A=B
One bond is a sigma bond
The other is a pi bond
Sigma bonds are formed by hybrid orbitals
overlapping.
Pi bonds are formed by unhybridized “p”
orbitals overlapping.
Double Bond
Triple Bond: A≡B
One bond is a sigma bond formed from
overlapping hybrid orbitals.
Two bonds are pi bonds formed from
overlapping “p” orbitals.
Triple Bond
Orbital Overlap and Molecular
Rotation
Sigma bonds allow free rotation of bonded
atoms.
Pi bond overlap restricts rotation of bonded
atoms.
Double bonds lead to cis- and trans-
isomers. Example: C2H2Cl2
Quiz
Describe the type of hybrid orbitals used by
each carbon and oxygen atom in the
following molecule:O
C
C O
HH
H H
Valence Shell Electron Pair Repulsion
Theory (VSEPR)Each group of valence electrons around a central atom is
located as far away from the other atoms as possible.
Strength of electron pair repulsions:
L.P.-L.P. > L.P.-B.P. > B.P.-B.P.
A “group” of electrons is defined as any number of electrons that occupies a space around an atom and may consist of a single, double, triple bond or lone pair of electrons.
The 3-dimensional arrangement of these groups determines the molecular arrangement (shape).
Molecular Shape
The arrangement of the atoms around a
central atom determines the shape of the
molecule or portion of the molecule.
Possible Shapes of Molecules
• Linear
• Bent
• Trigonal Planar
• Trigonal Pyramidal
• Square Planar
• Square Pyramidal
• Octahedral
• T-Shaped
• Seesaw
• Trigonal Bipyramidal
Different Shapes of Molecules
(or portions of molecules)
Two electron groups around the central atom:
– Linear shape
– Bond angle = 180o
Three electron groups
Trigonal planar arrangement
Bond angles = 120o
• Four electron groups
– Tetrahedral arrangement
– Bond angles = 109.5o
Five electron groups
Bipyramidal arrangement
Consists of three equatorial groups
And two axial groups
Six electron groups
-Octahedral arrangement
-Bond angles = 90o
Polar vs. Nonpolar Molecules
Generally speaking, a molecule will be nonpolar if …
(1) All of the bonded atoms (or groups of atoms) to the central atom are the same and equidistant from each other.
i.e. BH3 vs. BF3
(2) There are no lone pairs of electrons on the central atom(s).
(3) It is a hydrocarbon.