1. Covalent Bonding

2. 8.1 Molecules & Molecular Compounds
Molecule: a neutral group of atoms joined by covalent bonds
Diatomic Molecule: two atoms joined by a covalent bond
Examples: H2, Cl2, O2, NO, CO
Diatomic elements: Dr. Brinclhof
Molecular Compounds: Compounds composed of molecules (covalent bonds)

3. Comparison of Molecular & Ionic Compounds
Bonding
Covalent
Melting point
Lower
Higher
Electrolyte
Weak or non
Strong
Physical room temp
(s), (l), (g)
(s)

4. Molecular Formulas Show number & type of atoms in a molecule CH4, H2S
HNO3
C6H6
C3H7OH
(NH4)3PO4

5. Structural Formulas
Show the arrangement of atoms in a molecule

6. 8.2 Nature of Covalent Bonding
Octet rule is a guide
Electrons are shared to form a covalent bond

7. Formation of a Single Covalent Bond
Formed when two atoms share one pair of electrons

8. Why do some elements form diatomic molecules?

9. Single Covalent Bonds
The hydrogen and oxygen atoms attain noble-gas configurations by sharing electrons.

10. Ammonia, NH3

11. Drawing Electron Dot (Lewis) Structures
Lewis structure is a type of structural formula that depicts all the valence electrons in the molecule or ion
See Tutorial
Determine the total # ve
Connect atoms in such a way that all have a noble gas configuration (octet rule)
Carbon is often a central atom
Check

12. Draw Lewis Structures for these Molecular Compounds
HCl hydrogen chloride
Cl2 chlorine
I2 iodine
H2O2 hydrogen peroxide
PCl3 phosphorous trichloride
CH4 methane

13. Single, Double and Triple Covalent Bonds
Sometimes atoms share more than one pair of ve’s
A bond that involves on shared pair of e-s is a single covalent bond
Two shared pairs of electrons is a double covalent bond.
Three shared pairs of electrons is a triple covalent bond.

15. Acetylene A gas used in cutting steel Molecular formula is C2H2
Draw the Lewis structure for acetylene
Connect the atoms
Calculate ve’s
Form single covalent bonds between atoms
Complete octets until remainder of ve’s are used
Form double or triple bonds if needed to complete octets.

16. Polyatomic Ions Same process except…
Add or subtract e-s to account for the charge of the ion, for example
[NH4]+
[SO4]2-

17. Coordinate Covalent Bonds
Bonds in which one of the shared pair comes completely from one of the bonding atoms
Carbon Monoxide

18. Bond Energies Energy required to break a chemical bond
Energy released when a bond is formed
Is a measure of the strength of the bond
Large bond energies = strong bonds
Type of bond
Bond Energy (kJ/mol)
C─C
347
C=C
657
C≡C
908

19. Resonance Structures
A resonance structure is 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.
Actual bonding is a hybrid of all the possible resonance structures

20. Ozone Is an allotropic form of oxygen Molecular formula is O3
Is a pollutant (smog)
Protects earth by absorbing UV radiation
Draw the resonant Lewis structures for ozone

21. Nitrogen Dioxide Formed by lightning strikes Molecular formula NO2
Also a pollutant in automobile exhaust
Draw the resonance structures for NO2
Why is this an exception to the octet rule?

22. Exceptions to Octet Rule
When there is an odd number of ve,
NO2
Less than an octet:
Boron BF3
More than an octet:
Phosphorous PCl5
Sulfur SF6
Unfilled d-shells accept additional electrons, creating an “expanded” octet

23. 8.3 Bonding Theories Molecular orbitals
When covalent bonds form, atomic orbitals merge to form molecular orbitals

24. Sigma and Pi Bonds
Sigma bonds result atomic orbitals merge along the axis between nuclei (internuclear axis)
Pi bonds result when atomic orbitals merge to around the internuclear axis

25. Sigma Bonds
σ bonds are present in single covalent bonds.

26. Pi Bonds
π bonds are present in double and triple covalent bonds

27. Sigma and Pi Bonds C2H2

28. VSEPR Theory Valence Shell Electron Pair Repulsion Theory
The big idea:
Because covalent bonds and non-bonding pairs of electrons are areas of negative charge, they repel one another
Covalent bonds and non-bonding electrons are called “electron domains”

29. VSEPR Predicts the shape of small molecules
According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible.
How to predict the shape of the following molecules:
Draw the Lewis structure
Count the electron domains
Determine the geometry of the molecule (the way the atoms are arranged

30. Methane, CH4
Tetrahedron, bond angles of 109.5°

31. Ammonia, NH3 Trigonal pyramid, 107° Why is this not trigonal planar?
Why is the H-N-H bond angle not °?

32. Water, H2O Draw the Lewis structure Determine the total domains
Determine the bonding domains
Determine the shape of the molecule
Why is water a bend molecule and not a linear one?

33. Hybrid Orbitals
When covalent bonds form, atomic orbitals mix together to form hybrid orbitals
Atomic orbitals involved in bonding often contain a single unpaired electron
When the orbitals hybridize, a pair of electrons is shared
These hybrid orbitals are equal in number to the atomic orbitals which made them

34. Covalent Bond formation in CH4
In order for carbon’s 4 ve to be used in bonding, one 2s2 electron is promoted to 2p.
This results in 4 unpaired ve, which can then bond with unpaired e’s of other atoms.
In order to accomplish this, the atomic orbitals of C containing these ve hybridize.
One s and three p orbitals hybridize to form four equivalent orbitals, called sp3 orbitals

35. Covalent bonding in CH4
The s (one) and p (three) orbitals in the valence shell of C hybridize (merge) to form four equivalent sp3 orbitals.
They are called sp3 orbitals because they are formed from one s orbital and three p orbitals

36. Formation of Hybrid Orbitals

37. Hybrid Orbitals Hybridization Involving Single Bonds
In methane, each of the four sp3 hybrid orbitals of carbon overlaps with a 1s orbital of hydrogen.

38. Hybrid Orbitals Hybridization Involving Double Bonds
In an ethene molecule, two sp2 hybrid orbitals from each carbon overlap with a 1s orbital of hydrogen to form a sigma bond. The other sp2 orbitals overlap to form a carbon–carbon sigma bond. The p atomic orbitals overlap to form a pi bond. Inferring What region of space does the pi bond occupy relative to the carbon atoms?

39. Hybrid Orbitals Hybridization Involving Triple Bonds
In an ethyne molecule, one sp hybrid orbital from each carbon overlaps with a 1s orbital of hydrogen to form a sigma bond. The other sp hybrid orbital of each carbon overlaps to form a carbon–carbon sigma bond. The two p atomic orbitals from each carbon also overlap. Interpreting Diagrams How many pi bonds are formed in an ethyne molecule?

40. How to Determine Hybridization about an Atom
The principle: the number of hybrid orbitals must equal the number of atomic orbitals hybridized
Count the number of covalent bonds about an atom
This must equal the number of hybridized orbitals
Beginning with s, continue to add orbitals until the total equals the number of covalent bonds about the atom

41. Hybridization Chart
# bonds
Hybridization 2 sp 3
sp2 4
sp3 5 ?? 6

42. Predicting Hybridization
What hybridzation would be found about carbon in the following molecules?
HC≡CH sp
H2C=CH2
sp2
H3C-CH3
sp3

43. 8.4 Polar Bonds and Molecules
Electrons in a covalent bond are attracted to the nuclei of both atoms. Why?

44. Unequal Sharing of Bonding Electrons
When covalently bonded to another atom, some atoms attract electrons more strongly than others
These atoms have greater “electronegativity”
When bonded atoms differ in electronegativity, they do not share the bonding electrons equally

45. Bonding Electrons in HCl
Bonding e’s spend more time near Cl than H
What does this imply about Cl?
What does this imply about the distribution of electrical charge in HCl?

46. Polar Covalent Bonds
When bonded atoms are sufficiently different in electronegativity, the bond develops negative (-) and positive (+) ends
Why? Because the bonding e’s spend more time around the more electronegative element
This unequal distribution of (-) charge is called a dipole
The bond is called a polar covalent bond

47. Bond Character
Describes the type of charge distribution in a chemical bond
Based upon differences in electronegativity

48. Differences in Electronegativity and Bond Character

49. Polar Molecules
Molecules containing polar bonds may have an net dipole
The molecule may have a (+) and (-) side
Depends upon two factors
Presence of polar bonds
Geometry (shape) of molecule

50. Intermolecular Forces
Types of intermolecular forces account for differences between ionic and molecular substances.

51. Polar Molecules

52. Intermolecular Forces of Attraction
Not chemical bonds
Much weaker than covalent or ionic bonds
Van der Waals Forces
dipole-dipole interactions
London dispersion forces
Hydrogen Bonds
very important

53. Hydrogen Bonds Hydrogen bonds
Attraction between a hydrogen covalently bonded to a very electronegative atom to an unshared electron pair of another electronegative atom
often involve different molecules
Hydrogen bonding accounts for the unusual properties of water.

54. Hydrogen Bonding in Water