1. Chemical Bonding Batrachotoxin
A steroid alkaloid derived from skin secretions of the Phyllobates and Dendrobates genera of South American poison-arrow frogs. It is one of the most potent venoms known.
Notes 6

2. Bonds Forces that hold groups of atoms together and make them function
as a unit.
Ionic bonds – transfer of electrons
Covalent bonds – sharing of electrons

3. Electronegativity
The ability of an atom in a molecule to attract shared electrons to itself.
Linus Pauling

4. Table of Electronegativities

5. Covalent Bonds Polar-Covalent bonds Nonpolar-Covalent bonds
Electrons are unequally shared
Electronegativity difference between .3 and 1.7
Nonpolar-Covalent bonds
Electrons are equally shared
Electronegativity difference of 0 to 0.3

6. Polarity
A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

7. Bonding Forces Electron – electron repulsive forces Proton – proton
Electron – proton
attractive forces

8. Bond Length Diagram

9. Bond Energy It is the energy required to break a bond.
It gives us information about the strength of a bonding interaction.

10. Lewis Dot Diagrams
Electron dot diagrams, which are also called Lewis dot diagrams, are very useful tools in Chemistry. 
They will give you the ability to determine the type(s) of covalent bonds that an element may make in certain situations. 
They can also be used to predict the type of ion that an atom might make when it forms an ion. 
Each dot diagram consists of; an elemental symbol, which represents the kernel of the atom, and a group of 1-8 dots which shows the configuration of the outer-most electron shell of the atom, also called the valence shell.

11. An example of the proper Lewis dot diagram for the element oxygen.
The "O" in the example above represents the kernel of the atom, that is the nucleus and all of the electrons, except those in the valance (outer) shell. 
Each of the four "sides" of the symbol represents an orbital in the outermost energy level of the atom.

12. To make a Lewis dot diagram
You need to know how many electrons are in the valence shell. 
You fill in one valence electron on each side of the elemental symbol, and then double up as many sides as you need to in order to include each one.  
Remember that each side can only hold up to two dots!
You can place the first two dots on any side, but the rest of the dots should be placed in either a clockwise or counter clockwise manner, with no side receiving two dots until each side gets one.

13. More….
   By looking at the electron dot diagram for oxygen we can see that oxygen has two unpaired electrons, so it has two electrons available for standard covalent bonds.

14. Electron Dot Notation

15. The Octet Rule
Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.
Diatomic Fluorine

16. Hydrogen Chloride by the Octet Rule

17. Formation of Water by the Octet Rule

18. Comments About the Octet Rule
2nd row elements C, N, O, F observe the octet rule.
2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.
3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.
When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

19. Lewis Structures
Shows how valence electrons are arranged among atoms in a molecule.
Reflects central idea that stability of a compound relates to noble gas electron configuration.

20. Completing a Lewis Structure -CH3Cl
Make carbon the central atom
Add up available valence electrons:
C = 4, H = (3)(1), Cl = 7 Total = 14
Join peripheral atoms
to the central atom
with electron pairs. H .. ..
Complete octets on
atoms other than
hydrogen with remaining
electrons H .. C .. Cl .. .. .. H

21. Multiple Covalent Bonds: Double bonds
Two pairs of shared electrons

22. Multiple Covalent Bonds: Triple bonds
Three pairs of shared electrons

23. Resonance
Occurs when more than one valid Lewis structure can be written for a particular molecule.
These are resonance structures.
The actual structure is an average of
the resonance structures.

24. Resonance in Ozone
Neither structure is correct.

25. Resonance in Polyatomic Ions
Resonance in a carbonate ion:
Resonance in an acetate ion:

26. Covalent Network Compounds
Some covalently bonded substances DO NOT form discrete molecules.
Diamond, a network of covalently bonded carbon atoms
Graphite, a network of covalently bonded carbon atoms

27. Models
Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.
Models can be physical as with this DNA model
Models can be mathematical
Models can be theoretical or philosophical

28. Fundamental Properties of Models
A model does not equal reality.
Models are oversimplifications, and are therefore often wrong.
Models become more complicated as they age.
We must understand the underlying assumptions in a model so that we don’t misuse it.

29. Bonding with Lewis diagrams
These unpaired electrons might make two single covalent bonds, as is the case in water (H2O).  Or, they might make one double covalent bond, as the case of magnesium oxide (MgO).
  When Lewis dot diagrams are used for compounds, "x's" are often used to substitute for the dots of one or more elements in order to show which electrons came from which element. 

30. More with bonding..
Lewis dot diagrams for both oxygen and hydrogen as free elements, and then at water as a compound.

31. Ionic Bonds Electrons are transferred
Electronegativity differences are
generally greater than 1.7
The formation of ionic bonds is
always exothermic!

32. Sodium Chloride Crystal Lattice
Ionic compounds form solids at ordinary temperatures.
Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.

33. Table of ionic radii

34. VSEPR Model
(Valence Shell Electron Pair Repulsion)
The structure around a given atom is determined principally by minimizing electron pair repulsions.

35. Predicting a VSEPR Structure
Draw Lewis structure.
Put pairs as far apart as possible.
Determine positions of atoms from the way electron pairs are shared.
Determine the name of molecular structure from positions of the atoms.

36. Table – VSEPR Structures

37. VSEPR and the water molecule

38. VSEPR and the ammonia molecule

39. VSEPR and a molecule of I3
Which structure is the correct one?

40. VSEPR and Xenon tetrafluoride
Which one will it be???

41. VSEPR and Phosphorus hexachloride

42. Table of dipole moments

43. The Blending of Orbitals
Hybridization
The Blending of Orbitals

44. Lets look at a molecule of methane, CH4.
We have studied electron configuration notation and
the sharing of electrons in the formation of covalent
bonds.
Lets look at a molecule of methane, CH4.
Methane is a simple natural gas. Its molecule has a
carbon atom at the center with four hydrogen atoms covalently bonded around it.

45. Carbon ground state configuration
What is the expected orbital notation of carbon
in its ground state?
Can you see a problem with this?
(Hint: How many unpaired electrons does this
carbon atom have available for bonding?)

46. Carbon’s Bonding Problem
You should conclude that carbon only has TWO
electrons available for bonding. That is not not enough!
How does carbon overcome this problem so that
it may form four bonds?

47. Carbon’s Empty Orbital
The first thought that chemists had was that carbon promotes one of its 2s electrons…
…to the empty 2p orbital.

48. However, they quickly recognized a problem with such
an arrangement…
Three of the carbon-hydrogen bonds would involve
an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.
A Problem Arises

49. This would mean that three of the bonds in a methane
molecule would be identical, because they would involve
electron pairs of equal energy.
But what about the fourth bond…?
Unequal bond energy

50. The fourth bond is between a 2s electron from the
carbon and the lone 1s hydrogen electron.
Such a bond would have slightly less energy than the other bonds in a methane molecule.
Unequal bond energy #2

51. This bond would be slightly different in character than
the other three bonds in methane.
This difference would be measurable to a chemist
by determining the bond length and bond energy.
But is this what they observe?
Unequal bond energy #3

52. The simple answer is, “No”.
Measurements show that
all four bonds in methane
are equal. Thus, we need
a new explanation for the
bonding in methane.
Chemists have proposed an explanation – they call it
Hybridization.
Hybridization is the combining of two or more orbitals
of nearly equal energy within the same atom into
orbitals of equal energy.
Enter Hybridization

53. In the case of methane, they call the hybridization
sp3, meaning that an s orbital is combined with three
p orbitals to create four equal hybrid orbitals.
These new orbitals have slightly MORE energy than
the 2s orbital…
… and slightly LESS energy than the 2p orbitals.
sp3 Hybrid Orbitals

54. sp3 Hybrid Orbitals
Here is another way to look at the sp3 hybridization
and energy profile…

55. sp Hybrid Orbitals While sp3 is the hybridization observed in methane,
there are other types of hybridization that atoms
undergo.
These include sp hybridization, in which one s
orbital combines with a single p orbital.
Notice that this produces two hybrid orbitals, while
leaving two normal p orbitals

56. sp2 Hybrid Orbitals
Another hybrid is the sp2, which combines two orbitals from a p sublevel with one orbital from an s sublevel.
Notice that one p orbital remains unchanged.

57. Relative magnitudes of forces
The types of bonding forces vary in their strength as measured by average bond energy.
Strongest
Weakest
Covalent bonds (400 kcal)
Hydrogen bonding (12-16 kcal )
Dipole-dipole interactions (2-0.5 kcal)
London forces (less than 1 kcal)

58. Hydrogen Bonding Bonding between
hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen
Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests.

59. Hydrogen Bonding in Water

60. Hydrogen Bonding between Ammonia and Water

61. Dipole-Dipole Attractions
Attraction between oppositely charged regions of neighboring molecules.

62. The water dipole

63. The ammonia dipole

64. London Dispersion Forces
The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors.
London forces increase with the size of the molecules.
Fritz London

65. London Forces in Hydrocarbons