1. Covalent Bonding
Chapter 9

2. The Covalent bond
Remember that atoms bond to gain stability, usually meaning an octet of electrons.
Sometimes atoms who need to gain electrons will reach their octet by sharing electrons.
The bond formed by sharing electrons is called a covalent bond.
When two or more atoms bond covalently, a molecule is formed.
In covalent bonds the atoms move close enough together so that the repulsive force of like charged particles is balanced by the attractive forces between oppositely charged particles.

3. Examples of molecules Examples of molecules include diatomic elements:H2 N2 O2 F2
Cl2
Br2 I2

4. Lewis structures
A Lewis structure is a representation used when electrons are shared.
These structures show how the electrons are arranged in the molecules.
In the Lewis structure shared electron pairs, often called bonding pairs, are represented either by either a pair of dots or a line.
For example, the hydrogen molecule can be represented as:
H:H or H—H .

5. Bond type
In the formation of covalent bonds there are several types of bonds that can form.
Single bonds are always sigma bonds, in which the electron pair is shared in an area centered between two atoms.
These bonds can form by the overlap of two s orbitals, an s and a p orbital, or two p orbitals.

6. Multiple Bonds
Sometimes atoms fain stability by sharing more than one pair of electrons forming a multiple covalent bond.
For example, in a double covalent bond, two pairs of electrons are shared.
In a triple covalent bond, three pairs of electrons are shared.
A multiple covalent bond always consists of a sigma bond and at least one pi bond, a bond in which parallel orbitals overlap.
A pi bond occupies space above and below the line that represents where two atoms are joined.

7. Bond length and strength
The distance between the nuclei of two bonded atoms is called bond length.
Energy is absorbed when a bond breaks and released when a bond forms.
The amount of energy required to break a covalent bond is called the bond dissociation energy and is always positive.
The stronger the bond, the greater the bond dissociation energy, and the more difficult to break the bond.
Double bonds are stronger than single bonds; triple bonds are stronger than double bonds.
Double bonds are shorter than single bonds and triple bonds are shorter than double bonds.

8. Total energy change
The total energy change in a chemical reaction is determined from the energy of the bonds that have to be broken and formed.
In an endothermic reaction, more energy is required to break existing bonds in the reactants than is released when the new product bonds are formed.
In an exothermic reaction, more energy is released in the formation of the new products than is required to break bonds in the reactants.

9. Naming molecules
Section 9.2

10. Naming binary molecular compounds
Binary molecular compounds only contains two different elements.
In naming such compounds, given their formulas, use the following rules:
Name the first element in the formula using its name unchanged.
Name the second element, using the root of its name and adding the suffix –ide.
Use the following prefixes, given the number of each type of atom present:

11. Covalent Prefixes
Number
Prefix 1
Mono- 2
Di- 3
Tri- 4
Tetra- 5
Penta- 6
Hexa- 7
Hepta- 8
Octa- 9
Nona- 10
Deca-
****Mono- is not used as a prefix for the name of the first element.

12. Problems Name the following binary molecular compounds. SO2 P4O10 N2O3
SiF6

13. Naming binary acids
A binary acid contains hydrogen and one other element.
To name the acid, use the prefix hydro- to name the hydrogen part of the mole. Then to the root of the second element add the suffix –ic. Finally, add the word acid.

14. Oxyacids An oxyacid is a polyatomic ion that contains oxygen.
An oxyacid is an acid that contains hydrogen and an oxyanion.
To name an oxyacid, first write a form of the root of the name of the oxyanion. If the oxyanion name ends in –ate, name the acid by adding the suffix –ic to the root. If the oxyanion name ends in –ite, name the acid by adding the suffix –ous to the root. Then add the word acid.

15. Problems
Name the following acids.
H2Se
HBrO3
H2CO3 HI
HClO4
H2SO3

16. Formulas from Names
Given the name of a molecular compound, you can write its formula by analyzing the name in terms of the naming rules. Use the prefixes that indicate number to write the proper subscripts.

17. Problems Write the formulas for the following molecular compounds.
Disulfur dichloride
Dinitrogen tetroxide
Hydrosulfuric acid
Sulfuric acid

18. Molecular structures
Section 9.3

19. Molecular models
There are several types of molecular structures that illustrate the positions of the atoms in a molecular compound.
Those structures include:
Molecular formula
Space filling model
Lewis structure
Ball and stick model
Structural formula

20. Structural formulas
A structural formula is a molecular model that uses letter symbols and bonds to show relative positions of the atoms.
It can be predicted using the Lewis structure.

21. Lewis structure The steps for drawing Lewis structures are as follows:
Predict the location of certain atoms. Hydrogen is always a terminal atom. The atom with the least attraction for electrons is the central atom.
Find the total # of e- available for bonding (valence e-). If the structure is to represent a positive or negative polyatomic ion, the ion charge must be subtracted or added, respectively.
Divide the total # of available e- by 2 to obtain the # of bonding pairs.
Place one bonding pair between the central atom and each terminal atom.

22. Lewis rules (cont’d)
To find the total # of lone pairs and pairs available for multiple bonding, subtract the # of bonding pairs used in step 4 from the # of bonding pairs determined in step 3. Place lone pairs around the terminal atoms to satisfy the octet rule. Assign remaining pairs to the central atom.
If the central atom is not surrounded by four e- pairs, convert one or two lone pairs to the terminal atoms to a double or triple bond to the central atom.

23. problem Draw the Lewis structures for each of the following. CF4 CO
SiS2
NH4+

24. Resonance structures
Resonance occurs when more than one valid Lewis structure can be written for a molecule or an ion.
For example, three resonance structures can be written for the NO3- ion because the double bond can be placed between the central N atom and any of the three O atoms.

25. problem
Draw the three resonance structures for the carbonate ion (CO32-). (Hint: Each structure contains one double bond.)

26. Exceptions to the octet rule
Sometimes there are exceptions to the octet rule.
This may be because there are an odd # of total e-, as in ClO2.
Sometimes the central atom may have more or fewer than eight e-.
In the latter, an atom with a lone pair may attach to it by means of a coordinate covalent bond, a bond in which both shared e- are donated by only one of the atoms.

27. Problem
State why each of the following is an exception to the octet rule.
NO2
BCl3
PF5

28. Molecular shape
Section 9.4

29. Vsepr model Valence Shell Electron Pair Repulsion
The VSEPR model is used to determine molecular shape.
It assumes arrangements that minimize repulsion of electron pairs around the central atom.
Learn the molecular shapes in the table on page 260.

30. hybridization
During bonding, atomic orbitals can undergo hybridization, or mixing to form new identical hybrid orbitals.
An s orbital and three p orbitals can hybridize to form four identical sp3 hybrid orbitals.
These hybridized orbitals will form varying molecular shapes based on the number of lone pairs present.

31. Problem
Use VSPER model and the concept of hybridization to describe the hybrid orbitals and the shape of each of the following molecules.
H2S
BeF2
CBr4
NF3
BI3