1. TOPIC: Covalent Bonding
Essential Questions: How is Covalent bonding similar or different from ionic bonding

2. Why Do Atoms Bond?
The stability of an atom, ion or compound is related to its energy
lower energy states are more stable.
Metals and nonmetals gain stability by transferring electrons (gaining or losing) to form ions that have stable noble-gas electron configurations.
Ionic Bonding
Another way atoms can gain stability is by sharing valence electrons with other atoms, which also results in noble-gas electron configurations.
Covalent Bonding

3. The Covalent Bond
Atoms will share electrons in order to form a stable octet.
Covalent bond : the chemical bond that results from the sharing of valence electrons. Formed between nonmetals only.
also called a molecular bond

4. The Molecule formed when two or more atoms bond covalently
The smallest piece in a covalent compound

5. Models
Molecules

6. Single Covalent Bonds
In a single covalent bond a single pair of electrons is shared
A single line represents a single covalent bond
A single pair of electrons

7. Bonding Pair vs. Lone Pair
Bonding pair: a pair of electrons shared by two atoms
Lone pair: an unshared pair of electrons on an atom

8. Formation of Water

9. Group 17 elements will form one covalent bond.

10. Group 16 elements will form two covalent bonds.

11. Group 15 elements will form three covalent bonds.

12. Group 14 elements will form four covalent bonds.

14. Sigma Bonds Single covalent bonds are also called sigma bonds:
the electron pair is centered between two atoms.

16. Multiple Covalent Bonds
When more than one pair of electrons is shared, a multiple covalent bond is formed
Multiple bonds are made up of sigma bonds and pi bonds: formed when parallel orbitals share electrons.

17. Double Covalent Bond Two pairs of electrons are shared
Contains one sigma and one pi bond.

18. Triple Covalent Bond Three pairs of electrons are shared
Has one sigma and two pi bonds.

19. Strength of Covalent Bonds
The strength of covalent bonds is determined by the bond length:
distance between the bond nuclei
Shorter bond=stronger bond
Bond length is determined by:
The size of the atoms involved—larger atoms have longer bond lengths
How many pairs of electrons are shared—the more pairs of electrons shared, the shorter the bond length is.

20. Bond Dissociation Energy
the amount of energy required to break a bond
Indicates the strength of a covalent bond
When a bond forms, energy is released;
When a bond breaks, energy must be added
Each covalent bond has a specific value for its bond dissociation energy.

21. Bond Energy and Bond Length
A direct relationship exists between bond energy and bond length
Shorter Bond
Stronger Bond
Higher Bond Dissociation Energy
Longer Bond
Weaker Bond
Lower Bond Dissociation Energy

22. Energy Changes
An endothermic reaction is one where a greater amount of energy is required to break a bond in reactants than is released when the new bonds form in the products.
An exothermic reaction is one where more energy is released than is required to break the bonds in the initial reactants.

23. Naming Molecules Molecular Formula
Shows what atoms and how many are in a molecule
Examples:
Nonmetal-Nonmetal Combinations

24. Naming (Binary Compounds)
The first element is always named first using the entire element name
The second element is named using its root and adding the suffix -ide
Prefixes are used to indicate the number of atoms of each element that are present in the compound
mono=1 hexa= 6
Di= 2 hepta= 7
Tri=3 octa = 8
Tetra=4 nona = 9
Penta = 5 deca = 10

25. Naming Molecular Compounds
Greek prefixes are used to denote the number of atoms of each element present.

26. Naming Molecular Compounds
The prefix mono- is generally omitted for the first element.
For ease of pronunciation, we usually eliminate the last letter of a prefix that ends in “o” or “a” when naming an oxide.
Example: N2O5 is dinitrogen pentoxide not dinitrogen pentaoxide

27. Common Names
Many compounds were discovered and given common names long before the present naming system was developed (water, ammonia, hydrazine, nitric oxide).

28. Binary Acids
An acid that contains hydrogen and one other element Ex. HCl
ion ends –ide.
Name the acid with hydro-root of the anion-ic
HCl (hydrogen and chloride ) becomes hydrochloric.
HCl in a water solution is called hydrochloric acid.

29. Naming Acids Acids contain hydrogen as the first element
Binary Acids: H bonded to one other element
An ion that ends –ide
Name the acid with hydro-root-ic
Example: HCl
Hydrogen ion and chloride ion
Hydrochloric acid

30. Oxyacids An acid that contains both a hydrogen atom and an oxyanion.
Example: HNO3
Identify the oxyanion present.
name ends with the suffix –ate, replace it with the suffix –ic.
If the name ends with suffix –ite, replace it with suffix –ous,
NO3 is the nitrate ion so the acid is nitric acid.

31. Acid Naming Summary Anion name Acid name --ide Hydro-root-ic --ite
Root---ous
--ate
Root-ic

33. Naming Molecular Compounds
Greek prefixes are used to denote the number of atoms of each element present.

34. Naming Molecular Compounds
The prefix mono- is generally omitted for the first element.
For ease of pronunciation, we usually eliminate the last letter of a prefix that ends in “o” or “a” when naming an oxide.
Example: N2O5 is dinitrogen pentoxide not dinitrogen pentaoxide

35. Worked Example 5.7
Name the following binary molecular compounds: (a) NF3 and (b) N2O4.
Strategy Each compound will be named using the systematic nomenclature including, where necessary, appropriate Greek prefixes.
Solution (a) nitrogen trifluoride
(b) dinitrogen tetroxide
Think About It Make sure that the prefixes match the subscripts in the molecular formulas and that the word oxide is not preceded immediately by an “a” or an “o”.

36. Structural Formulas
A structural formula uses letter symbols and bonds to show relative positions of atoms.

37. Lewis Structures Used to predict the structural formula
Show arrangement of the atoms and un-bonded electrons

38. Five steps to draw Lewis structures:
Count the total number of valence electrons in all atoms involved.
Decide how the elements are arranged in the structure and draw it out.
Hydrogen is always an end atom.
Central atom is usually written first in compound
Central atom has least attraction for the electrons
Usually closer to left on periodic table
Subtract the # of electrons used in the bonds.

39. Satisfy the octets of the terminal atoms.
Place any remaining electrons around the central atom to satisfy its octet. If the central atom cannot be satisfied, make a multiple bond using a lone pair from the terminal atoms.
Check your work 

40. Examples

41. Count the total number of valence electrons in all atoms involved.
Drawing Lewis structures for polyatomic ions is very similar to drawing Lewis structures for covalent compounds EXCEPT in finding the number of electrons available for bonding
Count the total number of valence electrons in all atoms involved.
If the polyatomic ion is negatively charged, ADD the charge to the number of valence electrons.
If the ion is positively charged, SUBRACT the charge from the number of valence electrons.
Follow the rest of the steps to drawing Lewis structures.

42. Examples

43. Resonance Structures
When a molecule or polyatomic ion has both a double bond and a single bond, it is possible to have more than one correct Lewis structure:

44. Resonance
a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion.
The structures are called resonance structures.
A molecule that undergoes resonance behaves as if it has only one structure.

45. Exceptions to the Octet Rule
Three Ways Molecules Might Violate the Octet Rule

46. Odd Number of Valence Electrons
Some molecules have an odd number of valence electrons and cannot form an octet around each atom
Example: NO2

47. Sub Octet
Some compounds form with fewer than 8 electrons present around an atom.
Boron
BF3

48. Coordinate covalent bond
when one atom donates an entire pair of electrons to be shared with atoms or ions that need two more electrons.
Boron compounds often do this

49. Expanded Octet
Some elements can have more than eight electrons in their valence shell
Because of d-level electrons
PCl5

51. How? The d orbital starts to hold electrons.
This occurs in atoms in Period 3 or higher.  
When you draw Lewis structures for these compounds, extra lone pairs are added to the central atom OR the central atom will form more than four bonds.

52. Things to Remember
Any exceptions to the Octet Rule are on the central atom

53. Molecular Shape How a molecule “looks” Determines properties
The shape of a molecule determines whether or not two molecules can get close enough to react
We describe shape using the VSEPR model

54. VSEPR
This model is based on the fact that electrons pairs will stay as far away from each other as possible
Valence
Shell
Electron
Pair
Repulsion

55. How to apply VSEPR Draw the Lewis Structure for a Molecule
Count the pairs of bonded electrons
Count the pairs of unbonded electrons
Match the information with the VSEPR chart to classify the shape of the molecule

58. Atoms will assume certain bond angles: the angle formed by any two terminal atoms and the central atom
Lone pairs take up more space than bonded pairs do.

60. Molecular Polarity
Molecules are either polar or nonpolar depending on the bonds in the molecule.
We must look at the shape (geometry) of a molecule to determine polarity.
Symmetric molecules are nonpolar.
Asymmetric Molecules are Polar

61. Polar or NonPolar Determine if a molecule is polar or nonpolar by
Looking at a model of the molecule
Looking at a Lewis Structure of the molecule

62. Solubility of Polar Molecules
Bond type and shape of the molecule determine solubility
Polar substances and ionic substances will dissolve in polar solvents
Nonpolar substances will only dissolve in nonpolar substances

63. Intermolecular Forces
Nonpolar Molecules: Van der Waals intermolecular Forces
Very Weak forces between molecules
Polar Molecules: have dipole-dipole intermolecular bonding.
Stronger intermolecular Forces
Polar Molecules with Hydrogen Bonding: hydrogen bonded to nitrogen, oxygen or fluorine, it will have hydrogen bonding between molecules. A very strong dipole-dipole interaction
Very strong intermolecular Forces
High boiling points, high melting points