1. 6-1 Introduction to Bonding
Chemical Bond
Mutual attraction between the nuclei and valence electrons of different atoms that binds them together.
Types of Bonds
Ionic Bonds
Force of attraction between oppositely charged ions.
Covalent Bond
Force of attraction for electrons, that results in a pair of electrons being shared by two atoms.

2. Ionic or Covalent?
The Difference in electronegativity determines the bond type, page 161, according to the following scale:
Example Na-Cl
Na = 0.9 Cl = 3.0 Difference = 3.0 – 0.9 = 2.1
Ionic
Non Polar
Ionic
Polar Covalent
3.3
100%
1.7
50%
0.3 5% 0%

3. Polar Vs. Nonpolar Covalent Bonds
Non-Polar – electrons are equally shared
Balanced Distribution of Electrical Charge
Polar – electrons are not equally shared
Unbalanced Distribution of Electrical Charge
 +
 -

4. Table of Electronegativities

5. Identifying Bond Types
Determine if the following are going to form; Ionic, Polar Covalent or Non-polar Covalent Bonds:
C-N
Ca-F
Br-Br
H-Br

6. Summary of Bond Types

7. Chapter 6.2- Covalent Bonding
Molecule (Molecular Substance)
Group of atoms that are held together by a covalent bond.
Molecular Formula
Describes the composition of atoms in a single molecule of a compound.
Number and type of each atom in the molecule.
Diatomic molecule
A molecule containing only 2 atoms.
Seven Diatomic Elements
H2, N2, O2, F2, Cl2, Br2, I2

8. Formation of Covalent Bonds
Bond results from the attraction forces between the nuclei and electrons of the two atoms
Attraction Forces  - Potential Energy 
Repulsion Forces  - Potential Energy 
If the attraction forces overcome the repulsion forces the bond will form. (AF > RF)

9. Bond Length – the avg. distance between two atoms
AF  - BL 
Bond Energy – energy needed to break the bond
AF  - BE 

10. Electron Configurations
1s 2s p
Bonding Electron Pair
N ___ ___ ___ ___ ___
1s 2s p
Triple Bond – 3 shared pairs

11. Lewis structures
Formula that specifies which atoms are bonded to each other in a molecule.
Places atoms in a pattern in which they can share electrons to satisfy the octet rule.
Every shared pair of electrons are depicted as a dash ( — ).
Shared Pair, are 2 electrons that are shared between 2 atoms. (bonding pairs)
All unshared electrons must be paired.
Unshared Pair, are 2 electrons that are NOT shared between 2 atoms. (non-bonding pairs)

12. Single Covalent Bonds
Draw the Lewis Structure for the molecule that has 1 N and 3 H.
Identify which atom has the most free electrons.
Usually the lowest electronegative element.
Place this atom in the center of the structure.

13. Systematic over Trial-Error
Begin with the lowest electronegative element in the center.
Count the total number of valence electrons (N) needed to account for the atoms (based on the column of the atom in the periodic table) and charge (add one electrons for each negative charge; subtract one electron for each positive charge).
Draw the framework with single bonds. Some knowledge of the way the atoms are connected may be required.
Using lone pairs, complete octets around the non-central atoms.
Count the number of electrons depicted (two for each bond and two for each lone pair). If this number is less than N, then add electrons to the central atom until the total number of electrons depicted is N.
If the octet rule is not satisfied for the central atom and lone-pair electrons are nearby, use those electrons to make double or triple bonds to the central atom.

14. Practice Single Covalent Bonds
Write the Lewis Structure for the following formulas:
Cl2
OF2
C2H6
C3H7Cl

15. Multiple Covalent Bonds
Some structures need more than a single pair of electrons to reach the octet.
Double covalent bonds
Covalent bond between 2 atoms in which there are a total of 4 electrons being shared. (2pair)
Triple covalent bonds
Covalent bond between 2 atoms in which there are a total of 6 electrons being shared. (3pair)
CO2 and HCN

16. Polyatomic Ion Structures
Polyatomic ions, are covalently bonded atoms that form a charge due to the gain or loss of electrons.
When drawing the structures of polyatomic ions:
(–) ions must have extra electrons in the structure that is equivalent to its charge.
(+) ions must lose the number of electrons in the structure that is equivalent to its charge.

17. Example NH4 +1, the ammonium ion.
Starts out like the ammonia molecule NH3
But it has an extra H and a +1 charge, meaning the structure has 1 less electron.

18. Exceptions to the Octet
Incomplete Octet
Some atoms that have less than eight electrons.
Hydrogen only needs two and Boron needs six electrons.
Expanded Octet
Atoms that can have more than eight electrons.
Common expanded octets:
Cl, Br, I, S, and P
IMPORTANT!!!
Only one atom may exceed the standard eight valence.
Must be the central atom only!!

19. Coordinate Covalent Bond
Like a single covalent bond, but a single atom is sharing 2 electrons with another atom, in which it doesn’t increase the # of electrons in the atom sharing the 2 electrons.
This is determined by one atom having more electrons drawn than it originally contained.
Indicated by using an arrow rather than a dashed line. () The arrow must point away from the atom that has the pair of electrons.

20. Resonance
Bonding in molecules or ions that cannot be correctly represented by a single Lewis Structure. Can be determined by Formal Charges.
FC = V-N-B/2
Example: Ozone or CO2
O = C → O or O = C = O or O←C → O

21. 6.3 Ionic Bonding Ionic Compounds Formula Unit
Composed of positive and negative ions of equal but opposite charge.
Formula Unit
Simplest collection of ion in which the ionic compound can be formed.

22. 3D Array of forces is created (Lattice)
Strength of Array – varies w/ sizes, charges, & number of ions.
Arrangement of ions
gives the crystal its strength
To compare the strengths
Chemists compare energies released when the ionic compound forms.
Lattice Energy – energy released, 1 mole of an ionic compound is formed. (from gaseous ions)
CaF2

24. A comparison of Ionic and Molecular Compounds
Ionic Bonds are stronger than Covalent compounds
Due to the array of forces holding the solid together.
Ionic compounds
Higher Melting Point, Boiling Point
More Brittle
Do not vaporize easily
Conductors of electricity when
dissolved in H2O
Great repulsion when
layers shift

25. Lewis Dot Diagrams
Using electron dot notation, sketch the exchange of valence electrons in atoms as they form ionic bonds.

26. Class work
Show the electron transfer for the following elements and write the formula unit for the binary ionic compound:
1) Zn + I
2) Al + S
3) K + P
4) Mg + N
5) Ba + Se
6) Fe O

27. 6.4 Metallic Bonding
Bonding resulting from the attraction of positive ions and mobile electrons
Metal Atoms
Small Number of Valence Electrons
2 in the s-Block & 0 in the p-Block
Low Ionization Energy & Electronegativity
Easily give up electrons
At best – weakly covalent
Electrons are Delocalized
Electrons do not belong to any one ion

29. Metallic Properties Electron-Sea Model Explains Properties of Metals
Lustrous
Good Conductors of Heat & Electricity
Malleability
Bent or Shaped w/ Hammer
Ductility
Able to be drawn into a wire

30. 6.5 Molecular Geometry
Properties of Molecular Compounds depend upon 2 main things
Bonding
Molecular Geometry – arrangement of atoms in space
Used to determine Molecular Polarity
VSEPR Theory is used to predict molecular Geometry

31. Where would you expect the 2 Fluorine atoms to be located around Beryllium in BeF2?
2 F’s ?

32. Where would you expect the 2 Fluorine atoms to be located around Beryllium in BeF2?
2 F’s
Repulsion e- e- ?

33. Where would you expect the 2 Fluorine atoms to be located around Beryllium in BeF2?
2 F’s
Repulsion e- e- ?

34. Where would you expect the 2 Fluorine atoms to be located around Beryllium in BeF2?
2 F’s
Repulsion e- e-
The Greatest Distance Apart.

35. VSEPR Theory (Ves-per Ther ee)
Valence-Shell Electron-Pair Repulsion Theory
Electron pairs surrounding a central atom repel each other as far apart from each other as possible
Bonded Pairs – electrons in bonds
Unshared Pairs – unattached electrons
Only those around the central atom. A B
Unshared Pair
Bonded Atom

36. VSEPR & Molecular Geometry
Type of Molecule
Bonded Atoms
Unshared Pairs
Molecular Shape
AB2 2
Linear
AB2E 1
Bent
AB3 3
Triangular Planar-120o
AB4 4
Tetrahedral o
AB3E
Triangular Pyramidal
AB2E2
AB5 5
Triangular Bipyramidal-90o
AB6 6
Octahedral

37. Practice
Draw the Lewis structure then predict the molecular geometry of the structure:
SO2
SF6
NI3
CF4
SiO2
IF4 1-
BCl3

39. Molecular Polarity Polar Bond – uneven sharing of electrons
Oxygen has a greater electronegativity than carbon C O C
 +
 -

40. Molecular Polarity
Non-Polar Molecule O C
 +
 -

41. Molecular Polarity N H
Overall Polarity
• •

42. General Rule for determining Polarity.
Non-Polar
Zero Unshared Pairs
If all the bonded atoms are of the same type.
Polar
1 or More Unshared Pairs
Unless linear shape

43. Identify the polarity of the following:
F F
F S F F
F C F
  N
I I I F
F C F Cl
    O H
O = Si = O

44. Hybridization Carbon’s Electron Configuration .     1s 2s 2p
.    
1s 2s p
Each single electron is a bonding site.
Carbon only shows 2 bonding sites.
How is it possible for Carbon to bond to 4 Hydrogen atoms in Methane(CH4)?

45. Hybridization Carbon’s New Electron Configuration     . 1s 2s 2p
    .
1s 2s p
Moves one of the 2s electrons into the empty 2p orbital.
Creating four bonding sites. 
So we would expect 3 bonds to be identical and one to be different!

46. Hybridization
So we would expect 3 bonds to be identical and one to be different. C H s p
However, we find them all to be identical!

47. Hybridization
Hybridization – mixing 2 or more different orbitals, forming new identical orbitals.
For Carbon
1 s orbital + 3 p orbitals = 4 sp3 orbitals C H
sp3 s

48. 10.4

49. 10.4

50. Possible Combinations
1 s +1 p = 2 sp orbitals
1 s +2 p’s = 3 sp2 orbitals
1s +3 p’s = 4sp3 orbitals
1 s +3 p’s +1 d = 5 sp3d orbitals
1 s +3 p’s +2 d’s = 6 sp3d2 orbitals

51. Simplify To determine hybrid orbitals for a molecule.
1st count the number of bonded atoms & unshared pairs
2nd use the following chart, where the number is the count from step 1.
3rd use a superscript for a letter that appears more than once. Example H2O = 4  sppp  sp3
Example BF3= 3  spp  sp2
123456
spppdd

52. Practice
Determine the hybridization and Polarity of the central atom for the following molecules:
1) CO2
2) SO2
3) SF6
4) SiO2

53. Molecular Bonding Molecular Orbitals 2 Types of Molecular Bonds
Resulting orbital formed by the overlap of 2 atomic orbitals.
2 Types of Molecular Bonds
Sigma Bond
Molecular orbital that is formed along the bonding axis.
Pi Bond
Molecular orbital that is formed above or below the bonding axis.

54. Sigma Bond

55. Sigma with Pi bond

56. Sigma with 2 Pi bond

57. Sigma bond (s) – electron density between the 2 atoms
Pi bond (p) – electron density above and below plane of nuclei of the bonding atoms
Sigma bond (s) – electron density between the 2 atoms
10.5

58. 10.5

59. Sigma (s) and Pi Bonds (p)
1 sigma bond
Single bond
Double bond
1 sigma bond and 1 pi bond
Triple bond
1 sigma bond and 2 pi bonds
How many s and p bonds are in the acetic acid
(vinegar) molecule CH3COOH? C H O
10.5

60. Intermolecular Forces
Forces of attraction between molecules.
Vary in strength
Weaker than Intramolecular Forces (Bonds)
Strength of IMF’s is related to the substance’s boiling points
IMF  BP 
3 Types of IMF’s
Dipole-Dipole
Hydrogen Bonding
London Dispersion

61. Dipole-Dipole
Attraction between polar molecules + -

62. Hydrogen Bonding Particularly Strong Dipole-Dipole
Attraction between Hydrogen atoms & an unshared electron pair on a strongly electronegative atom in another molecule
F, O, N
Holds DNA Together

63. London Dispersion Forces
Attractions between molecules that have an induced dipole or a momentary dipole
Due to the constant motion of molecules Induced dipole

64. Bond Type Substance Boiling Point (°C) Noble Gas He -269 Non-Polar H2
Boiling Points & Bond Types
Bond Type
Substance
Boiling Point (°C)
Noble Gas He
-269
Non-Polar H2
-253
Cl2
-34
Br2 59
Polar
NH3
-33
H2S
-61
H2O
100
Ionic
NaCl
1413
Metallic W
5660