1. Chapter 6 Table of Contents Chemical Bonding
Section 1 Introduction to Chemical Bonding
Section 2 Covalent Bonding and Molecular Compounds
Section 3 Ionic Bonding and Ionic Compounds
Section 4 Metallic Bonding
Section 5 Molecular Geometry

2. Chapter 6 Objectives Define chemical bond.
Section 1 Introduction to Chemical Bonding
Chapter 6
Objectives
Define chemical bond.
Explain why most atoms form chemical bonds.
Describe ionic and covalent bonding.
Explain why most chemical bonding is neither purely ionic nor purely covalent.
Classify bonding type according to electronegativity differences.

3. Chemical Bond
An attractive force that holds atoms or ions together

4. Types of Chemical Bonds
Ionic
-electrical attraction between positive and negative ions
-transfer of electrons
-usually a metal and a nonmetal
Cation- positive ion
Anion- negative ion

5. Types of Chemical Bonds, cont’d
Covalent
-sharing of electrons between two atoms to make a bond
-can be: nonpolar covalent (equal sharing)
polar covalent (unequal sharing)

6. Types of Chemical Bonds, cont’d
Determining the Type of Bond
-bonds are not 100% ionic or covalent
-use the table of electronegativities (pg.161)
3.3 – 1.8 ionic bonds
polar covalent
nonpolar covalent

7. Chapter 6 Objectives Define molecule and molecular formula.
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Objectives
Define molecule and molecular formula.
Explain the relationships among potential energy, distance between approaching atoms, bond length, and bond energy.
State the octet rule.

8. Chapter 6 Molecular Compounds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Molecular Compounds
A molecule is a neutral group of atoms that are held together by covalent bonds.
A chemical compound whose simplest units are molecules is called a molecular compound.

9. Visual Concepts
Chapter 6
Molecule

10. Chapter 6 Molecular Compounds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Molecular Compounds
The composition of a compound is given by its chemical formula.
A chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.
A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound.

11. Molecular Compounds One type of molecular compound is called a
Diatomic molecule
examples- O2, this is composed of two atoms of oxygen
others:

12. Formation of a Covalent Bond
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Formation of a Covalent Bond
Atoms have lower potential energy when they are bonded to other atoms
The figure below shows potential energy changes during the formation of a hydrogen-hydrogen bond.
Check fig. explanation in the textbook.

13. Characteristics of the Covalent Bond
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Characteristics of the Covalent Bond
The distance between two bonded atoms at their most stable energy state is the bond length.
In forming bond, atoms release energy. The same amount of energy must be added to separate the bonded atoms. This is called Bond energy .

14. Bond Energies & Bond Lengths for Single Bonds- Inverse relationship
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Bond Energies & Bond Lengths for Single Bonds- Inverse relationship

15. Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
The Octet Rule
Noble gas atoms are unreactive because their electron configurations are especially stable.
noble-gas atoms’ outer s and p orbitals are completely filled by a total of eight electrons.
Other atoms can fill their outermost s and p orbitals by sharing electrons through covalent bonding.
Such bond formation follows the octet rule: Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest energy level.

16. The Octet Rule, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
The Octet Rule, continued
Exceptions to the Octet Rule
Exceptions to the octet rule –atoms have more or less than 8 electron in outer energy level
Hydrogen - only two electrons.
Boron – only six electrons.
Main-group elements in Periods 3-7 can form bonds with expanded valence, involving more than eight electrons.

17. Chapter 6 Lewis Structures
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures
Electron-dot notation can also be used to represent molecules.
The pair of dots between the two symbols represents the shared electron pair
For a molecule of fluorine, F2, the electron-dot notations of two fluorine atoms are combined.

18. Chapter 6 Lewis Structures
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures
The pair of dots between the two symbols represents the shared pair of a covalent bond.
In addition, each fluorine atom is surrounded by three pairs of electrons that are not shared in bonds.
An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.

19. Chapter 6 Lewis Structures
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures
The pair of dots representing a shared pair of electrons in a covalent bond is often replaced by a long dash.
example:
A structural formula indicates the kind, number, and arrangement, and bonds but not the unshared pairs of the atoms in a molecule.
example: F–F H–Cl

20. Chapter 6 Objectives, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Objectives, continued
Use the basic steps to write Lewis structures.
Determine Lewis structures for molecules containing single bonds, multiple bonds, or both.
Explain why scientists use resonance structures to represent some molecules.

21. Rules for Drawing Lewis Structures
Determine the total valence electrons for all atoms.
Draw a skeleton. Central atom is C or least electronegative. Never H.
Give each atom an octet. (some exceptions)
Double or triple bond if you don’t have enough electrons. (COPSN)
Extra electrons? Put on the central atom.

22. Lewis Structures, continued
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Lewis Structures, continued
Examples:
Draw the Lewis structure of
a. iodomethane, CH3I.
b. Ammonia, NH3
c. Silane, SiH4
d. Phosphorous trifluoride, PF3

23. Multiple Covalent Bonds
Section 2 Covalent Bonding and Molecular Compounds
Chapter 6
Multiple Covalent Bonds
A double covalent bond, or simply a double bond, is a covalent bond in which two pairs of electrons are shared between two atoms.
A triple covalent bond, or simply a triple bond, is a covalent bond in which three pairs of electrons are shared between two atoms.
In general, double bonds have greater bond energies and are shorter than single bonds.
Triple bonds are even stronger and shorter than double bonds.

24. Lewis Structures with Multiple Bonds
Examples:
Formaldehyde, CH2O
Carbon dioxide, CO2
Hydrogen cyanide, HCN
Ethyne, C2H2

25. Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Objectives
Compare a chemical formula for a molecular compounds with one for an ionic compound.
Discuss the arrangements of ions in crystals.
Define lattice energy and explain its significance.
List and compare the distinctive properties of ionic and molecular compounds.
Write the Lewis structure for a polyatomic ion given the identity of the atoms combined and other appropriate information.

26. Chapter 6 Ionic Compounds
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Ionic Compounds
Most of the rocks and minerals that make up Earth’s crust consist of positive and negative ions held together by ionic bonding.
example: table salt, NaCl,
An ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal.

27. Ionic Vs. Covalent Bonding
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Ionic Vs. Covalent Bonding

28. Formation of Ionic Compounds
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Formation of Ionic Compounds

29. Formation of Ionic Compounds, continued
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
Formation of Ionic Compounds, continued
In an ionic crystal, ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice.
The combined attractive and repulsive forces within a crystal lattice determine:
the distances between ions.
the pattern of the ions’ arrangement in the crystal

30. NaCl and CsCl Crystal Lattices
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
NaCl and CsCl Crystal Lattices

31. A Comparison of Ionic and Molecular Compounds
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
A Comparison of Ionic and Molecular Compounds
The force that holds ions together in an ionic compound is a very strong electrostatic attraction.
In contrast, the forces of attraction between molecules of a covalent compound are much weaker.
This difference in the strength of attraction between the basic units of molecular and ionic compounds gives rise to different properties between the two types of compounds.

32. A Comparison of Ionic and Molecular Compounds, continued
Section 3 Ionic Bonding and Ionic Compounds
Chapter 6
A Comparison of Ionic and Molecular Compounds, continued
Molecular compounds have relatively weak forces between individual molecules.
They melt at low temperatures.
The strong attraction between ions in an ionic compound gives ionic compounds some characteristic properties, listed below.
very high melting points
hard but brittle
not electrical conductors in the solid state, because the ions cannot move