Ch05 Final Fix

Copyright © 2010 Pearson Prentice Hall, Inc.

John E. McMurry • Robert C. Fay

Lecture NotesAlan D. Earhart

Southeast Community College • Lincoln, NE

General Chemistry: Atoms First

Chapter 5Covalent Bonds and Molecular Structure

Copyright © 2010 Pearson Prentice Hall, Inc. Chapter 5/2

Molecules and the Covalent Bond

Covalent Bond: A bond that results from the sharing of electrons between atoms.

Molecule: The unit of matter held together by covalent bonds.

Chapter 5/3



Strengths of Covalent Bonds

Chapter 5/6

Chapter 5/7

A Comparison of Ionic and Covalent Bonds


Polar Covalent Bonds: Electronegativity

Electronegativity: The ability of an atom in a molecule to attract the shared electrons in a covalent bond.

Chapter 5/9


Chapter 5/10


Chapter 5/11


Chapter 5/12



Naming Molecular Compounds

Because nonmetals often combine with one another in different proportions to form different compounds, numerical prefixes are usually included in the names of binary molecular compounds.


Naming Molecular Compounds

N2O4

The second element listed is more anionlike and takes the name of the element with an “ide” modification to the ending.

The first element listed is more cationlike and takes the name of the element.

The prefix is added to the front of each to indicate the number of each atom.

dinitrogen tetraoxide


Electron-Dot Structures

Electron-Dot Structures (Lewis Structures): A representation of an atom’s valence electrons by using dots and indicates by the placement of dots how the valence electrons are distributed in the molecule.

Chapter 5/16


Chapter 5/17


Chapter 5/18


Chapter 5/19

Electron-Dot Structures of Polyatomic Molecules

Step 1: Valence Electrons• Count the total number of valence electrons for

all atoms in the molecule.• Add one additional electron for each negative

charge in an anion or subtract one for each positive charge in a cation.

Chapter 5/20


Step 2: Connect Atoms• Draw lines to represent bonds between atoms.• For hydrogen and second row atoms, use the

number of bonds listed below.• For third row and greater atoms, they may have

more bonds than predicted by the octet rule.• The least electronegative atom is usually the

central atom.



Step 4: Assign Electrons to the Central Atom• If unassigned electrons remain after step 3,

place them on the central atom.

Step 3: Assign Electrons to the Terminal Atoms• Subtract the number of electrons used for

bonding in the previous step from the total number determined in step 1.

• Complete each terminal atom’s octet (except for hydrogen).



Step 5: Multiple Bonds• If no unassigned electrons remain after step 3

but the central atom does not yet have an octet, use one or more lone pairs of electrons from a neighboring atom to form a multiple bond (either a double or a triple).



2(1) + 6 = 8 valence electrons

Step 4:

Step 1:

Step 2:

bonding pair of electrons

lone pair of electrons

HO

H

Draw an electron-dot structure for H2O.

HO

H



Draw an electron-dot structure for CCl4.

4 + 4(7) = 32 valence electrons

Step 3:

Step 1:

Step 2: ClC

Cl

Cl

Cl

ClC

Cl

Cl

Cl



Draw an electron-dot structure for H3O1+.

3(1) + 6 - 1 = 8 valence electrons

Step 4:

Step 1:

Step 2: HO

H

HHO

H

H

1+



Draw an electron-dot structure for CH2O.

4 + 2(1) + 6 = 12 valence electrons

Step 3:

Step 1:

Step 2: HC

O

H

HC

O

H

Step 5: HC

O

H

HC

O

H



Draw an electron-dot structure for SF6.

6 + 4(7) = 34 valence electronsStep 1:

Step 2: Step 3:

F

F

S

FF

F F

F

F

S

FF

F F



Draw an electron-dot structure for ICl3.

7 + 3(7) = 28 valence electronsStep 1:

Step 2: Step 4:

Cl

I

ClCl

Cl

I

ClCl

Step 3:

Cl

I

ClCl


Electron-Dot Structures and Resonance

Draw an electron-dot structure for O3.

Step 1:

Step 2:

3(6) = 18 valence electrons

Step 4:

Step 5:

OOO

OOOStep 3:

OOO

OOO


Electron-Dot Structures and Resonance

Step 4: OOO

Or, move a lone pair from this oxygen?

Move a lone pair from this oxygen?

OOO OOO

Resonance


Formal Charges

FormalCharge

# ofvalence e-

in free atom-

21

-# of

nonbondinge-

# ofbonding

e-=

Calculate the formal charge on each atom in O3.

OOO

6 - (2) - 6 = -1126 - (4) - 4 = 0

12 6 - (6) - 2 = +1

12


Molecular Shapes: The VSEPR Model

VSEPR: Valence-Shell Electron-Pair Repulsion model

Electrons in bonds and in lone pairs can be thought of as “charge clouds” that repel one another and stay as far apart as possible, this causing molecules to assume specific shapes.

Working from an electron-dot structure, count the number of “charge clouds,” and then determine the molecular shape.

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Two Charge Clouds

Chapter 5/34


Three Charge Clouds

Chapter 5/35


Four Charge Clouds

Chapter 5/36


Four Charge Clouds

Chapter 5/37


Five Charge Clouds

Chapter 5/38


Five Charge Clouds

Chapter 5/39


Five Charge Clouds

Chapter 5/40


Five Charge Clouds


Five Charge Clouds

Chapter 5/42


Six Charge Clouds

Chapter 5/43


Six Charge Clouds

Chapter 5/44


Six Charge Clouds

Chapter 5/45


Six Charge Clouds

Chapter 5/46

Chapter 5/47


Valence Bond Theory

sigma () bonds

Valence Bond Theory: A quantum mechanical model which shows how electron pairs are shared in a covalent bond.


Valence Bond Theory

Valence Bond Theory: A quantum mechanical model which shows how electron pairs are shared in a covalent bond.

• Covalent bonds are formed by overlap of atomic orbitals, each of which contains one electron of opposite spin.

• Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms.

• The greater the amount of overlap, the stronger the bond.


Hybridization and sp3 Hybrid Orbitals

How can the bonding in CH4 be explained?

4 valence electrons2 unpaired electrons




4 valence electrons4 unpaired electrons



4 nonequivalent orbitals




4 nonequivalent orbitals


4 equivalent orbitals



Other Kinds of Hybrid Orbitals

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Chapter 5/59


Chapter 5/60



Molecular Orbital Theory: The Hydrogen Molecule

Atomic Orbital: A wave function whose square gives the probability of finding an electron within a given region of space in an atom.

Molecular Orbital: A wave function whose square gives the probability of finding an electron within a given region of space in a molecule.



* antibonding orbitalhigher in energy

bonding orbitallower in energy

Bond Order =(# Bonding e- - # Antibonding e-)

2



= 12

2 - 0Bond Order =



= 02

2 - 2Bond Order: =

21

22 - 1


Molecular Orbital Theory: Other Diatomic Molecules

Oxygen, O2, is predicted to be diamagnetic by electron-dot structures and valence bond theory.

However, it is known to be paramagnetic.

O2 OO

Diamagnetic: All electrons are spin-paired. It is weekly repelled by magnetic fields.

Paramagnetic: There is at least one unpaired electron. It is weakly attracted by magnetic fields.


Chapter 5/68


Chapter 5/69


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Ch05 Final Fix

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