CHAPTER 8 COVALENT BONDS8.1 Molecular Compounds8.2 The Nature of Covalent Bonding8.3 Bonding Theories8.4 Polar Bonds and Molecules

8.1 Molecular Compounds

Key Concepts:1. How are melting points and boiling points of

molecular compounds different from ionic compounds?

2. What information does a molecular formula provide?

Molecules and Molecular Compounds1. Covalent bond – occurs when two or more atoms share

valence electrons.2. Molecule – is a neutral group of atoms joined together

by covalent bonds.3. Diatomic molecule – is a molecule consisting of two

atoms.4. Compound – a substance that contains two or more

elements chemically combined in a fixed proportion5. 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-metals6. Using page 214 illustrate some differences between

ionic and covalent compounds.

Molecular Formula:1. 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 structure2. Using page 215 state the molecular formula for Ammonia

and describe the types of diagrams and models used to represent Ammonia.

8.2 The Nature of Covalent Bonding

Key Concepts1. How does electron sharing occur in forming covalent bonds?2. How do electron dot structures represent shared electrons?3. How do atoms form double or triple covalent bonds?4. How are coordinate covalent bonds different other covalent

bonds?5. How is the strength of a covalent bond related to its bond

dissociation energy?6. How are oxygen atoms bonded in ozone?7. What are some exceptions to the octet rule?

The Octet Rule in Covalent Bonding1. In forming covalent bonds, electron sharing

usually occurs so that atoms attain the electron configurations of noble gases.

2. That is to say the valence electrons arrange themselves so that each atom sees an octet.

3. Hydrogen has a noble gas configuration with 2 electrons

4. Groups four to seven are likely to form covalent bonds

Single Covalent Bonds1. 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. 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.

Double and Triple Covalent Bonds1. Atoms sometime bond by sharing more than one

pair of electrons.2. Double Covalent Bond – Shares two pair of

electrons3. Triple Covalent Bond – Shares three pairs of

electrons4. Try showing bonding Carbon Dioxide

Coordinate Covalent Bonds1. Is a covalent bond in which one atom contributes both

bonding electrons2. In a structural formula, you can show coordinate covalent

bonds as arrows that point from the atom donating the pair of electrons.

3. Once formed, a coordinate covalent bond is like any other covalent bond.

4. Most polyatomic cations and anions contain both covalent and coordinate covalent bonds.

5. Compounds containing polyatomic ions include both ionic and covalent bonding.

6. Polyatomic ions have charge in order to satisfy the octet rule for each atom present in the group.

7. Show the coordinate covalent bond of Carbon Monoxide.8. Show the formation of the Ammonium ion. 9. Show the formation of Sulfate.

10. Using page 224 of your text, show the chemical and structural formula for the following Molecular Compounds.

a. Nitrous Oxideb. Sulfur Trioxidec. Hydrogen Fluorided. Nitric Oxidee. Hydrogen Peroxidef. Nitrogen Dioxideg. Hydrogen Cyanideh. Hydrogen Chloridei. Sulfur Dioxide11. 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.

Bond Dissociation Energies1. The energy required to break the bond between two

covalently bonded atoms.2. Usually expressed as the energy needed to break one

mole of bonds.3. A large bond dissociation energy corresponds to a

strong covalent bond.4. High dissociation energies tend to create very stable

compounds that tend to be chemically unreactive.5. Units are measured in kJ/mo16. A mol is a chemical quantity of an element or

compound in which there are 6.02x1023 atoms or molecules present.

Link

Resonance1. 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.

2. Although no back-and-forth changes occur, double –headed arrows are used to connect resonance structures.

3. Show the structural formation of ozone.

Exceptions to the Octet Rule1. 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.

2. Draw two resonance structures for Nitrogen Dioxide3. Other odd number electron molecules are Chlorine Dioxide

and Nitric Oxide.4. Several molecules with an even number of electrons, such

as some compounds of Boron, also fail to follow the octet rule.

5. Draw the structure for Boron Trifluoride and show the significance of it reacting with ammonia.

6. A few atoms, Phosphorus and Sulfur, can have ten or twelve electrons instead of eight

7. Draw the structure for Phosphorus Pentachloride and Sulfur Hexafluoride

Exceptions to the Octet RuleThere 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.

8.3 Bonding Theories1. How are atomic and molecular orbitals related?2. How does VESPR theory help predict the shapes of molecules?3. In what ways is orbital hybridization useful in describing molecules?

Molecular Orbitals

1. 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.

2. Each atomic orbital is full when it contains two electrons.3. Bonding Orbitals – in covalent bonds two electrons are also

required to fill a molecular orbital.4. 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 bondsAre H2 and F2

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.

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 apartas possible creating three dimensional structures

Therefore, VSEPR diagrams are characterized by the number of lone pair electrons (unshared electron pairs) and the angles betweenthe shared pairs of electrons

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 AX4ammonia, H3N:, is AX3E1water, 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 geometriesAX2 and AX2E3 both give rise to linear geometries

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

A Couple of More Advanced Examples:

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