1. CHAPTER 8 COVALENT BONDS
8.1 Molecular Compounds
8.2 The Nature of Covalent Bonding
8.3 Bonding Theories
8.4 Polar Bonds and Molecules

2. 8.1 Molecular Compounds
Key Concepts:
How are melting points and boiling points of molecular compounds different from ionic compounds?
What information does a molecular formula provide?

3. Molecules and Molecular Compounds
Covalent bond – occurs when two or more atoms share valence electrons.
Molecule – is a neutral group of atoms joined together by covalent bonds.
Diatomic molecule – is a molecule consisting of two atoms.
Compound – a substance that contains two or more elements chemically combined in a fixed proportion
Molecular compounds – a compound composed of molecules
a. Have low melting points
b. Have low boiling points
c. Most are gas or liquid at room temperature
d. Composed of two or more non-metals
6. Using page 214 illustrate some differences between ionic and covalent compounds.

4. Molecular Formula:
Molecular Formula – the chemical formula of a molecular compound
a. Describes how many of each atom a molecule contains
b. Subscripts are used after the element’s symbol to indicate the number of atoms of each element in the molecule.
c. Reflects the actual number of atoms in each molecule and are not necessarily the lowest whole-number ratios.
d. Can describe molecules consisting of one element.
e. Does not tell you about the molecule’s structure
2. Using page 215 state the molecular formula for Ammonia and describe the types of diagrams and models used to represent Ammonia.

5. 8.2 The Nature of Covalent Bonding
Key Concepts
How does electron sharing occur in forming covalent bonds?
How do electron dot structures represent shared electrons?
How do atoms form double or triple covalent bonds?
How are coordinate covalent bonds different other covalent bonds?
How is the strength of a covalent bond related to its bond dissociation energy?
How are oxygen atoms bonded in ozone?
What are some exceptions to the octet rule?

6. The Octet Rule in Covalent Bonding
In forming covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases.
That is to say the valence electrons arrange themselves so that each atom sees an octet.
Hydrogen has a noble gas configuration with 2 electrons
Groups four to seven are likely to form covalent bonds

7. Single Covalent Bonds
1. Single Covalent Bond – Two atoms held together by sharing a pair of electrons.
2. Hydrogen is an example.
3. An electron dot structure can be used to show the shared pair of electrons of the covalent bond.
4. Using page 218 use electron dots to combine two Fluorine atoms then show the electron configuration for each atom.
5. Structural Formula – represents the covalent bonds by using dashes, each dash represents one electron pair.
6. Unshared Pair – are electrons not shared between atoms – also called lone pair, nonbonding pair.
7. Draw the electron dot structure for ammonia (NH3) show the unpaired bonds and the shared pairs properly.

8. 8. Now draw the structure for methane
9. Draw the electron configuration for Carbon then using p220 of the text explain why Carbon usually forms four bonds.

9. Double and Triple Covalent Bonds
Atoms sometime bond by sharing more than one pair of electrons.
Double Covalent Bond – Shares two pair of electrons
Triple Covalent Bond – Shares three pairs of electrons
Try showing bonding Carbon Dioxide

10. Coordinate Covalent Bonds
Is a covalent bond in which one atom contributes both bonding electrons
In a structural formula, you can show coordinate covalent bonds as arrows that point from the atom donating the pair of electrons.
Once formed, a coordinate covalent bond is like any other covalent bond.
Most polyatomic cations and anions contain both covalent and coordinate covalent bonds.
Compounds containing polyatomic ions include both ionic and covalent bonding.
Polyatomic ions have charge in order to satisfy the octet rule for each atom present in the group.
Show the coordinate covalent bond of Carbon Monoxide.
Show the formation of the Ammonium ion.
Show the formation of Sulfate.

11. 10. Using page 224 of your text, show the chemical and structural formula for the following Molecular Compounds.
Nitrous Oxide
Sulfur Trioxide
Hydrogen Fluoride
Nitric Oxide
Hydrogen Peroxide
Nitrogen Dioxide
Hydrogen Cyanide
Hydrogen Chloride
Sulfur Dioxide
11. The electron dot structure for a neutral molecule contains the same number of electrons as the total number of valence electrons in the combining atoms.
12. The negative charge of a polyatomic ion shows the number of electrons in addition to the valence electrons.
13. Because a negatively charged polyatomic ion is part of an ionic compound, the positive charge of the cation of the compound balances these additional electrons.

12. Bond Dissociation Energies
The energy required to break the bond between two covalently bonded atoms.
Usually expressed as the energy needed to break one mole of bonds.
A large bond dissociation energy corresponds to a strong covalent bond.
High dissociation energies tend to create very stable compounds that tend to be chemically unreactive.
Units are measured in kJ/mo1
A mol is a chemical quantity of an element or compound in which there are 6.02x1023 atoms or molecules present.
Link

13. Resonance
A structure that occurs when it is possible to draw two or more valid electron dot structures that have the same number of electron pairs for a molecule or ion.
Although no back-and-forth changes occur, double –headed arrows are used to connect resonance structures.
Show the structural formation of ozone.

14. Exceptions to the Octet Rule
The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number. There are also molecules in which an atom has fewer, or more, than a complete octet of valence electrons.
Draw two resonance structures for Nitrogen Dioxide
Other odd number electron molecules are Chlorine Dioxide and Nitric Oxide.
Several molecules with an even number of electrons, such as some compounds of Boron, also fail to follow the octet rule.
Draw the structure for Boron Trifluoride and show the significance of it reacting with ammonia.
A few atoms, Phosphorus and Sulfur, can have ten or twelve electrons instead of eight
Draw the structure for Phosphorus Pentachloride and Sulfur Hexafluoride

15. Exceptions to the Octet Rule
There are three classes of exceptions to the octet rule
1) Molecules with an odd number of electrons;
2) Molecules in which one atom has less than an octet;
3) Molecules in which one atom has more than an octet.
Odd Number of Electrons
Few examples. Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons.
Less than an Octet Less
Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A.
Most typical example is BF3.
More electrons than an Octet
This is the largest class of exceptions. Atoms from the 3rd period onwards can accommodate
more than an octet. Beyond the third period, the d orbitals are low enough in energy to participate in bonding and accept the extra electron density.

16. 8.3 Bonding Theories
How are atomic and molecular orbitals related?
How does VESPR theory help predict the shapes of molecules?
In what ways is orbital hybridization useful in describing molecules?

17. Molecular Orbitals
Molecular orbitals are created when two atoms combine by the overlap of each atoms atomic orbital creating an orbital that applies to the entire molecule.
Each atomic orbital is full when it contains two electrons.
Bonding Orbitals – in covalent bonds two electrons are also required to fill a molecular orbital.
Sigma Bonds – are created when two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei
Two examples of sigma bonds
Are H2 and F2

18. 5. Pi Bonds – Are created by the side by side overlap of p orbitals
the bonding electrons are most likely to be found in sausage-shaped
regions above and below the bond axis.
Atomic orbitals of pi bonding
overlap less than in sigma bonding
therefore, pi bonds tend
to be weaker than sigma bonds.

19. VSEPR – Valence Shell Electron-Pair Repulsion Theory
The repulsion theory between electron pairs causes molecular
shapes to adjust so that the valence –electron pairs stay as far apart
as possible creating three dimensional structures
Therefore, VSEPR diagrams are characterized by the number of
lone pair electrons (unshared electron pairs) and the angles between
the shared pairs of electrons

20. The AXE system
American* general chemistry textbooks adopt the excellent AXmEn system, where A is the central atom, m the number of ligands X, and n the number of nonbonded lone-pairs of electrons, E, about the central atom.
methane, CH4, is AX4 ammonia, H3N:, is AX3E1 water, H2O, is AX2E2
Note that different AXmEn designations can give rise to the same overall geometry or shape:
AX2E1 and AX2E2 both give rise to bent or angular geometries AX2 and AX2E3 both give rise to linear geometries

21. The AXE system gives rise to a pattern, from which the various atomic geometric shapes can be determined/assigned:

22. A Couple of More Advanced Examples:

23. Hybrid Orbitals
-Provides information about both molecular bonding and molecular shape unlike VSEPR theory that just deals with molecular shape.