1. Unit 3: Chemical Bonding and Molecular Structure
Cartoon courtesy of NearingZero.net

2. Bonds
Forces between nuclei and valence electrons of different atoms that binds the atoms together– bonds form because it LOWERS the potential energy
Ionic bonds – transfer of electrons; smallest unit is called a “formula unit”
Covalent bonds – sharing of electrons; smallest unit is a “molecule”

3. Properties of Ionic Compounds
Structure:
Crystalline solids
Melting point:
Generally high
Boiling Point:
Electrical Conductivity:
Excellent conductors, molten and aqueous
Solubility in water:
Electrons:
Elements:
Generally soluble
transferred
Metal + nonmetal

4. Properties of Covalent Compounds
Structure:
Gases and some liquids
Melting point:
Generally low
Boiling Point:
Electrical Conductivity:
Poor conductors, except aqueous acids
Solubility in water:
Elements:
Two Types:
Polar covalent are soluble
Nonmetal + nonmetal
Polar and nonpolar

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

6. Table of Electronegativities

7. 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
IONIC Compounds have EN differences >1.7

8. Determining Bond Type:
Bond EN Diff Type? More EN?
S to H
Ca to Cl
Mg to O
F to F
Which of the above bonds has the LEAST ionic character?
Which of the above bonds has the MOST ionic character?

9. The Octet Rule – Ionic Compounds
Ionic compounds tend to form so that each atom, by gaining or losing electrons, has an octet of electrons in its highest occupied energy level.

10. Ionic Bonding: The Formation of Sodium Chloride
Sodium has 1 valence electron
Chlorine has 7 valence electrons
An electron transferred gives
each an octet
Na 1s22s22p63s1
Cl 1s22s22p63s23p5

11. Ionic Bonding: The Formation of Sodium Chloride
This transfer forms ions, each with an octet:
Na+ 1s22s22p6
Cl- 1s22s22p63s23p6

12. Ionic Bonding: The Formation of Sodium Chloride
The resulting ions come together due to electrostatic attraction (opposites attract):
Na+
Cl-
The net charge on the compound must equal zero

13. Examples of Ionic compounds
All salts, which are composed of metals bonded to nonmetals, are ionic compounds and form ionic crystals.
Examples:
MgCl2
Na2O KI
CaO
BaS
LiF

14. Monatomic Cations
Name H+
Hydrogen
Li+
Lithium
Na+
Sodium K+
Potassium
Mg2+
Magnesium
Ca2+
Calcium
Ba2+
Barium
Al3+
Aluminum

15. Monatomic Anions
Name F-
Fluoride
Cl-
Chloride
Br-
Bromide I-
Iodide
O2-
Oxide
S2-
Sulfide
N3-
Nitride
P3-
Phosphide

16. 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.

17. Representation of Components in an Ionic Solid
Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

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

19. Metallic Bonding
The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons
Electrons are “delocalized”
Occurs between like metals in the D block Example: Cu to Cu

20. Properties of Metals
Metals are good conductors of heat and electricity
Metals are malleable
Metals are ductile
Metals have high tensile strength
Metals have luster

21. Bond Energy Bonding Energy gives the strength of covalent compounds…
Bonding Energy is the energy required to break a bond
For ionic bonds, we measure Lattice Energy – the amount of energy it takes to form the bond

22. The Octet Rule – Covalent Compounds
Covalent compounds tend to form so that each atom, by sharing electrons, has an octet of electrons in its highest occupied energy level.
Diatomic Fluorine

23. Hydrogen Chloride by the Octet Rule

24. Electron Dot Notation

25. 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.

26. 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

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

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

29. Bond Length and Bond Energy
Length (pm)
Energy (kJ/mol)
C - C
154
346
C=C
134
612
CC
120
835
C - N
147
305
C=N
132
615
CN
116
887
C - O
143
358
C=O
799
CO
113
1072
N - N
145
180
N=N
125
418
NN
110
942

30. 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.

31. Resonance in Ozone
Neither structure is correct.

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

33. 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.

34. Table – VSEPR Structures

35. 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.

36. Intermolecular Forces
Forces between molecules
3 Types:
Hydrogen bonding
Dipole-dipole forces
London Disperson Forces

37. 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.

38. Hydrogen Bonding in Water

39. Hydrogen Bonding between Ammonia and Water

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

41. The water dipole

42. 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

43. 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)

44. The Blending of Orbitals
Hybridization
The Blending of Orbitals

45. 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.

46. 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?)

47. 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?

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

49. 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

50. 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

51. 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

52. 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

53. 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

54. 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

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

56. 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

57. 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.