1. Covalent Bonding
Chapter 9

2. The Covalent Bond Section 9.1
Why do atoms bond?
To achieve full outer electron shells
Octet Rule- atoms gain, lose, or share electrons to achieve the electron configuration of noble gases
Gain and Lose IONIC BONDING
Share  COVALENT BONDING

3. What is a covalent bond?
A chemical bond that results in the SHARING of valence electrons
Occurs between 2 or more nonmetals

4. A molecule is formed when two or more atoms bond covalently
Examples: sugars, DNA, proteins, fats, carbohydrates, cotton, synthetic fibers

5. Formation of a covalent bond
REMEMBER: Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2),Chlorine (Cl2) Bromine (Br2) and Iodine (I2) occur in nature as diatomic molecules

6. Covalent Bonding
An attractive force occurs between the protons of one atom and the electrons of the other atom
When a single pair of electrons is shared, such as in the hydrogen molecule, a single covalent bond forms

7. Lewis Structures
Use electron-dot diagrams to show how electrons are arranged in molecules

8. Group 7A (Halogens) have 7 VE
One more VE is necessary
A single covalent bond will form
Group 6A have 6 VE
Two more VE are necessary
Two covalent bonds will form
Group 5A have 5 VE
Three more VE are necessary
Three covalent bonds will form
Group 4A have 4 VE
Four more VE are necessary
Four covalent bonds will form

9. Sigma Bonds- single covalent bonds- when electron pairs are centered between two atoms
Multiple Bonds
In many molecules, atoms attain noble gas configuration by sharing more than one pair of electrons between two atoms
Carbon, Nitrogen, Oxygen, and Sulfur most often form multiple bonds

10. Strength of Covalent Bonds
The strength of covalent bonds depends on how much distance separates both nuclei
The distance between the two bonding nuclei at the position of maximum attraction is called bond length
Determined by the size of the atoms and how many electron pairs are shared
Bond length decreases as the number of bonds increases (triple bond has a shorter bond length than a single bond)

11. Energy Changes
An energy change accompanies the forming or breaking of a bond between atoms in a molecule.
Energy is released when a bond forms
Energy must be added to break the bonds of a molecule
The amount of energy required to break a specific covalent bond is called bond dissociation energy

12. Bond dissociation energy indicates the strength of a chemical bond because a direct relationship exists between bond energy and bond length
In chemical reactions, bonds in reactant molecules are broken and new bonds are formed as product molecules form

13. Endothermic reactions occur when a greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds from in the product molecules
Exothermic reactions occur when more energy is released forming new bonds than is required to break bonds in the initial reactants

14. Checkpoint
What is a covalent bond? How does it differ from an ionic bond?
What type of elements form covalent bonds?
Draw the Lewis Structures for each of these molecules:
PH3
H2S
CCl4

15. Naming Molecules Section 9.2
Naming Binary Molecular Compounds
The first element in the formula is always named first, using the entire element name.
The second element in the formula is named using the root of the element and adding the suffix – ide.
Prefixes are used to indicate the number of atoms of each type that are present in the compound.

16. Prefixes in Covalent Compounds
Number of Atoms
Prefix 1
Mono- 6
Hexa- 2
Di- 7
Hepta- 3
Tri- 8
Octa- 4
Tetra- 9
Nona- 5
Penta- 10
Deca-

17. Practice Naming
CCl4
As2O3 CO
SO2
NF3

18. Naming Acids Binary Acids
Use the prefix hydro- to name the hydrogen part of the compound
The rest of the name consists of a form of the root of the second element plus the suffix –ic, followed by the word acid.
Examples:
HCl  Hydrochloric Acid
HCN Hydrocyanic acid (even though there are more than 2 elements present, if no oxygen is present- the acid is named as a binary)

19. Oxyacids
Acids that contain an oxyanion (polyatomic ion that contains oxygen)
The name of the oxyacids consists of a form of the root of the anion, a suffix, and the word acid
If the anion suffix is –ate, it is replaced with the suffix –ic
If the anion suffix is –ite, it is replaced with the suffix –ous.

20. Practice Problems HI
HClO3
HClO2
H2SO4
H2S

21. Formulas and Names of Covalent Compounds
Checkpoint
Write the molecular formula for each of the following compounds
Disulfur trioxide
Iodic acid
Dinitrogen monoxide
Hydrofluoric acid
Phosphorus pentachloride
What is the difference between a binary acid and oxyacid?
Complete the following table 
Formulas and Names of Covalent Compounds
Formula
Name
PCl5
Hydrobromic acid
H3PO4
Oxygen difluoride
SO2

22. Molecular Structures Section 9.3
Structural Formula- uses letter symbols and bonds to show relative positions of atoms
Lewis Structure Procedure
1. Predict the location of certain atoms
Hydrogen is always a terminal, or end, atom. Because it can share only one pair of electrons, hydrogen can be connected to only one other atom
The atom with the least attraction of shared electrons in the molecule is the central atom. This element usually is the one closer to the left on the periodic table. The central atom is located in the center of the molecule, and all other atoms become terminal atoms.

23. 2. Find the total number of electrons available for bonding
2. Find the total number of electrons available for bonding. This total is the number of valence electrons in the atoms in the molecule.
3. Determine the number of bonding pairs by dividing the number of electrons available for bonding by two.
4. Place one bonding pair (single bond) between the central atom and each of the terminal atoms.

24. 5. Subtract the number of pairs you used in step 4 from the number of bonding pairs you determined in step 3. The remaining electron pairs include lone pairs as well as pairs used in double and triple bonds. Place lone pairs around each terminal atom bonded to the central atom to satisfy the octet rule. Any remaining pairs are assigned to the central atom.

25. 6. If the central atom is not surrounded by 4 electron pairs, it does not have an octet. You must convert one or two of the lone pairs on the terminal atoms to a double bond or a triple bond between the terminal atom and the central atom. These pairs are still associated with the terminal atom as well as with the central atom. Remember that, in general, carbon, nitrogen, oxygen, and sulfur can form double or triple bonds with the same element or with another element.

26. Space-filling Structure Ball-and-stick molecular model
Lewis Structure
Ball-and-stick molecular model

27. Examples
NH3
PO4-3
CO2
BH3

28. Resonance Structures
Resonance is a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion
Nitrate Ion Resonance Structures:

29. Exceptions to the Octet Rule
A small group of molecules has an odd number of valence electrons and cannot form an octet around each atom. (i.e. NO2 , ClO2, and NO)
Some compounds form with fewer than eight electrons present around an atom. (Example: Boron)
When one atom donates a pair of electrons to be shared with an atom or ion that needs two electrons to become stable, a coordinate covalent bond forms

30. 3. Some central atoms contain more than eight valence electrons (expanded octet)
Examples: PCl5, SF6, and XeF4

31. Molecular Shape Section 9.4
The shape of the molecule determines many of its physical and chemical properties.
Molecular shape is determined by the overlap of orbitals that share electrons
Valence Shell Electron Pair Repulsion model or VSEPR model
Based on an arrangement that minimizes the repulsion of shared and unshared pairs of electrons around the central atom

32. VSEPR Model
The angle formed by any two terminal atoms and the central atom is a bond angle.
Shared electron pairs repel one another
Lone pairs of electrons occupy a slightly larger orbital than shared electrons
Shared bonding orbitals are pushed slightly together by lone pairs

33. Look at the VSEPR cheat sheet 

34. Hybridization
A hybrid results from combining two of the same type of object, and it has characteristics of both.
Hybridization- a process in which atomic orbitals are mixed to form new, identical hybrid orbitals.

35. Practice Problems
Determine the molecular geometry and bond angle for the following:
BF3
NH4+
OCl2
CF4

36. Electronegativity and Polarity Section 9.5
Electron affinity is a measure of the tendency of an atom to accept an electron

37. The character and type of a chemical bond can be predicted using the electronegativity difference of the elements that are bonded
Polar Covalent- Unequal sharing
Nonpolar Covalent- Equal sharing
Generally- Ionic bond form when the electronegativity difference is greater than 1.70

38. Polar Covalent Bonds
Polar covalent bonds form because not all atoms that share electrons attract them equally
The shared pair of electrons is pulled toward one of the atoms
Partial charges occur at the ends of the bond
Partially negative
Partially positive 
The resulting polar bond is referred to as a dipole (two poles)

39. Molecular Polarity
Molecules are either polar or nonpolar, depending on the location and nature of the covalent bonds they contain.
A polar molecule has a partial negative charge on one side, while the other side of the molecule has a partial positive charge
Note: symmetric molecules are usual nonpolar and molecules that are asymmetric are usually polar

40. Polar Molecule or not?
Compare H2O and CCl4

41. Solubility of polar molecules
The ability of a substance to dissolve in another substance is known as the physical property solubility
The bond type and the shape of the molecules present determine solubility
“Likes dissolve likes”
Polar compounds are usually soluble in polar substances
Nonpolar molecules only dissolve in nonpolar substances

42. Properties of Covalent Compounds
Lower melting and boiling points (indicating weak bond strength)
Many are liquids or gases at room temperature
Do not conduct electricity
Many do not dissolve in water (polar)

43. Practice Problems
Decide whether each of the following molecules is polar or nonpolar
SCl2
H2S
CF4
CS2

44. Vocabulary on Test Structural formula molecule VSEPR model
Coordinate covalent bond
hybridization
oxyacid
electronegativity
polar covalent
covalent bond
resonance
endothermic
exothermic
terminal atom
Sigma bond