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Chapter 7Chemical 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

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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

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Ionic Bonding Question

Describe how calcium fluoride would form

from atoms of calcium and fluorine.

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

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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

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Fig 7.5

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

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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

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Lattice Energy

Lattice Energy is the energy absorbed that occurs when an ionic solid is separatedinto isolated ions in the gas phase. (ΔH is positive)

Also, the energy released when gaseous ions come together to form the crystal. (ΔH 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 …

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AlCl3

•Crystal structure matters for lattice energy.•Since we can’t account for every crystal structure, we will only consider attraction between one ion to another.

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

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Lattice Energy is proportional to:

| charge M+ X charge NM-|

distance between nuclei

Note: Magnitude of charge dominates over size.

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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

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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.

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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

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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

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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?

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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

GCE Daily – April 17

• Office Hours: MTWF 11:00 – 11:30 or appt

• Lab this week: #3 & #4

• Lewis Structures & Molecular Shape

• No class on Friday

• Exam II: chapters 7 & 8

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Lewis Dot Structures Review

*See pg 22-23 in lab manual

Draw Lewis Dot Structures of:

SiO2

SO32-

NO2-

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.

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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 PreferredLewis 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.

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Types of orbital overlap:

sigma (end-to-end)

pi (side-by-side)

Note: sigma bond side-by-side overlap is

not permitted.

σ bonds

π bonds

• A double bond is made of a σ bond and a π

bond

• A π bond is not a direct overlap of orbitals,

and looser sharing

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Let’s Consider CH4

How can carbon form four bonds?

What is the shape of methane?

What are the bond angles?

Tetrahedral

Hybrid Orbital Theory

Theory depicts a mix of the atomic orbitals (hybridize) to form the necessary number of orbitals needed for bonding.

# hybrid orbitals = # atomic orbitals mixed

The type of hybrid orbitals � # atoms attached + # lone pairs of e-

Types of Hybrid Orbitals

sp

sp2

sp3

sp3d

sp3d2

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Hybridization

Hybrid Orbital Theory

Theory depicts a mix of the atomic orbitals (hybridize) to form the necessary number of orbitals needed for bonding.

# hybrid orbitals = # atomic orbitals mixed

The type of hybrid orbitals � # atoms attached + # lone pairs of e-

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sp Hybrids

Consider BeCl2

Central Be needs two hybrid orbitals to accommodate the two bonded Cl 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 Linear

Note: The bonding between the sp hybrid

orbital of beryllium and the “s” orbital of the hydrogen atom is considered a sigma

bond overlap.

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Isolated C atom sp-hybridized C

2 π bonds

2 σ bonds

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 Trigonal Planar

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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.

sp2

sp3 Hybrids

Back to CH4

The central carbon atom needs four hybrid orbitals.

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Bond angles between sp3 hybrids = 109.5o

Tetrahedral

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

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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.

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Bond angles between sp3d2 hybrids = 90o

Octahedral

Summary of Hybrid Theory

v v v

& Shape

Hybrid Quiz

Predict the type of hybrid orbital you would

expect in the central atom of the following molecules:

SiH4 BH3 AsF5 AlCl3 SF4

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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.

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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

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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 (Angular)

• Trigonal Planar

• Trigonal Pyramidal

• Square Planar

• Square Pyramidal

• Octahedral

• T-Shaped

• Seesaw (Sawhorse)

• Trigonal Bipyramidal

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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

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• Four electron groups

– Tetrahedral arrangement

– Bond angles = 109.5o

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Five electron groups

Bipyramidal arrangement

Consists of three equatorial groups

And two axial groups

Six electron groups

-Octahedral arrangement

-Bond angles = 90o

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Hybridization

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.

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Another way to determine

overall polarity of a molecule…

• Do the polarity vectors cancel each other?

– Yes � Nonpolar molecule

– No � Polar molecule

• Examples:

– CO2, NCl3, CF4, NH4+

Electron Density Map: The redder the more electron density

Hybridization

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Quiz

• What are the hybridization & molecular

geometry of the following?

– CH3CCCH3, SO32-, AsF5, CHCl3

– Are they polar or non polar molecules?