1. Unit 7 – Bonding & Molecular Geometry

2. Definitions Chemical Bonds Force that holds atoms together
It’s all about the electrons (e-)
Electrons available for bonding are called valence electrons!

3. Types of Chemical Bonds
Ionic Bond
Bond between metal and nonmetal due to “electrostatic interactions”
Attraction between positively and negatively charged ions (cations and anions)
Electrons are completely transferred from metal to nonmetal

4. Ionic bonds Result from a Transfer of Valence Electrons+ -

5. Types of Chemical Bonds
Covalent Bond
Bonds in which e- are shared
Most common type

6. Shared Electrons Complete ShellsF F

7. Definitions Octet rule (Rule of 8)
Most atoms want 8 e- in the outer shell this is most stable
H2 and He want a “duet” (2 e-)
Electron configuration for duet = ns2
Electron configuration for octet = ns2 np6

8. Lewis Structures (Atoms)
A Lewis dot diagram depicts an atom as its symbol and its valence electrons.
Ex: Carbon . . . C .
Carbon has four electrons in its valence shell (carbon is in group 14), so we place four dots representing those four valence electrons around the symbol for carbon.

9. Drawing Lewis Dot Diagrams
Electrons are placed one at a time in a clockwise manner around the symbol in the north, east, south and west positions, only doubling up if there are five or more valence electrons.
Same group # = Same Lewis Dot structure
Ex. F, Cl, Br, I, At
Example: Chlorine (7 valence electrons b/c it is in group 17) . . . . . Cl . .

10. Paired and Unpaired Electrons
As we can see from the chlorine example, there are six electrons that are paired up and one that is unpaired.
When it comes to bonding, atoms tend to pair up unpaired electrons.
A bond that forms when one atom gives an unpaired electron to another atom is called an ionic bond.
A bond that forms when atoms share unpaired electrons between each other is called a covalent bond.

11. Maximum # of valence electrons = 8
Note: In the final structure, the placement of the dots around the element is not crucial:
Maximum # of valence electrons = 8

12. Bonding in Ionic Compounds
The ionic bond forms from attraction of cations for anions.

13. Structure of Ionic Compounds
Ionic compounds have
formula units—these show ratio of ions in the crystal lattice.

14. Writing Lewis Structures for Ions
Uses either 0 or 8 dots, brackets and a superscript charge designate to ionic charge
Ex.) Li+, Be+2, B+3, C+4, N-3, O-2, F-1

15. Writing Lewis Structures (Ionic Compounds)
Lewis Dot Diagrams of Ionic Compounds
Ex. 1) MgO
Ex. 2) Li2O

16. Lewis Representations of Ionic Structures
MgO
Li2O

17. Drawing Lewis Structures (Covalent Compounds)
Covalent compound - a substance made up of atoms which are held together by covalent bonds (in which e- are shared)
They are also called molecules.

18. Covalent Compounds and Lewis Structures
Lewis structures for covalent compounds show the bonds and how the atoms will connect.
Shared e- = bonding e-
Non-shared e- = lone pair e- (a.k.a. non-bonding e-)
Ex. H2O

19. Drawing Electron Dot Diagrams for Molecules
Chemists usually denote a shared pair of electrons as a straight line.
F F
Sometimes the nonbonding pair of electrons are left off of the electron dot diagram for a molecule

20. Examples H
CH4 H C H H H N H
NH3 H

21. Types of Covalent Bonds
Single Bond
2 e- are shared in a bond (1 from each atom)
Double Bond
2 pairs of e- are shared (4 e- total, 2 from each atom)
Triple Bond
3 pairs of e- are shared (6 e- total, 3 from each atom)

22. Rules for Drawing Lewis Dot Diagrams
Add up the total number of valence e- for each atom in the molecule.
Each (-) sign counts as 1 e-, each (+) sign subtracts one e-
Write the symbol for the central atom then use one pair of e- to form bonds between the central atom and the remaining atoms.
Count the number of e- remaining and distribute according to octet rule (or the “duet” rule for hydrogen)
If there are not enough pairs, make sure the most electronegative elements are satisfied. Then, start shifting pairs into double and triple bonds to satisfy the octet rule.
If there are extra e-, stick them on the central atom.

23. Hints: H is NEVER a central atom!
Halogens (Group 17) are usually not central atoms.
If you only have 1 of a certain element, it is usually the central atom.

24. Checking Your Work! But Remember....
The Structure MUST Have: the right number of atoms for each element, the right number of electrons, the right overall charge, and 8 electrons around each atom (ideally).

25. Examples: F2
H2O
OCl-
PO43-

26. Examples: O2
CH4 HF
NH3

27. Examples:
NH4+
SO32- N2
CH3OH

28. Exceptions to the Octet Rule
Reduced Octets – electron deficient molecules
(Be and B)
Be: 2 valence e-, doesn’t form
octet (BeH2: Be has 4 e-)
B: 3 valence e-, doesn’t form octet
(BF3: B has 6 e-)
BF3

29. Exceptions to the Octet Rule
Expanded Octets
(Examples: P, S, Cl, As, Se, Br, Kr, Xe)
How to recognize:
The central atom in PERIOD 3 or greater is surrounded by > 4 atoms.
You draw the Lewis diagram and the results don’t make sense – the central atom has > 8 e-

30. Expanded Octets (P, S, Cl, As, Se, Br, Kr, Xe)
Examples:
PF XeF4

31. Resonance Structures are created by moving electrons, NOT atoms.
Definition:
When a single Lewis structure does not adequately represent a substance, the true structure is intermediate between two or more structures which are called resonance structures.
Resonance Structures are created by moving electrons, NOT atoms.

32. Resonance Structure Example, SO2
Central atom = S
This leads to the following structures:
These equivalent structures are called
RESONANCE STRUCTURES. The true structure is a HYBRID of the two.
Arrow means “in resonance with”

33. Resonance Structures, NO3-
Draw the Lewis diagram for NO3- with all possible resonance structures.

34. Free Radicals
When there is an odd # of total electrons, there will be a single, unpaired electron in the structure!
Example: NO
Radicals are extremely reactive; they want to have paired electrons to be more stable!!

35. Molecular Geometry Molecular Geometry describes the
3-D arrangement of atoms in a molecule.
We will use VSEPR theory to predict these 3-D shapes!

36. There is a fundamental geometry that corresponds to the total number of electron pairs around the central atom: 2, 3, 4, 5 and 6
linear
trigonal
planar
tetrahedral
trigonal
bipyramidal
octahedral

37. VSEPR: Shapes of Molecules
VSEPR Theory (definition)
“Valence Shell Electron Pair Repulsion”
Based on idea that e- pairs want to be as far apart as possible
The molecule adopts the shape that minimizes the electron pair repulsions.
Based on molecular shape of Lewis structure

38. VSEPR: Shapes of Molecules
We define the electron pair geometry by the positions in 3D space of ALL electron pairs (bonding and non-bonding).
The molecular geometry only considers the positions of the bonded electrons.

39. To determine the electron pair geometry:
1. Draw the Lewis structure.
2. Count the number of bonded (X) atoms and non-bonded or lone pairs (E) around the central atom.
3. Based on the total of X + E, assign the electron pair geometry.
4. Multiple bonds count as one bonded pair!

40. Basic Electron Pair Geometries
Shapes Sum of Bonded Atoms & Lone e-
Linear
Trigonal planar
Tetrahedral
Trigonal bipyramidal
Octahedral

41. Molecular Geometry Notation
4/28/2017
Molecular Geometry Notation
A: Central Atom
X: Bonded Atom
E: Non-bonding electron pair
(Lone pair e- on central atom)
BP’s: Bonding Pairs
LP’s: Lone Pairs
Dr. Mihelcic Honors Chemistry

42. Molecular Geometry: Two Bonded Items
Electron Pair Geometry = linear
Total Bonds to C.A.
# bonded atoms (X)
# lone pairs (E)
AXE Notation
MolecularGeometry
Bond
Angles
Example 2
AX2
Linear
180°
BeCl2

43. Molecular Geometry: Three Bonded Items
Electron Pair Geometry =
trigonal planar
Total Bonds to C.A.
# bonded atoms (X)
# lone pairs (E)
AXE Notation
MolecularGeometry
Bond
Angles
Example 3
AX3
Trigonal Planar
120°
BCl3 2 1
AX2E
Bent
118°
NO2-

44. Molecular Geometry: Four Bonded Items
e- pair geometry = tetrahedral
Total Bonds to C.A.
Bonded atoms (X)
Lone pairs (E)
AXE Notation
MolecularGeometry
Bond
Angles
Example 4
AX4
Tetrahedral
109.5°
CCl4 3 1
AX3E
Trigonal pyramid
107°
NH3 2
AX2E2
Bent
104.5°
H2O

45. 5 AX5 AsCl5 4 1 AX4E SeCl4 3 2 AX3E2 BrCl3 AX2E3 Linear XeF2
Total Bonds to C.A.
Bonded atoms (X)
Lone pairs (E)
AXE Notation
MolecularGeometry
Bond
Angles
Example 5
AX5
Trigonal bipyramid
90, 180, 120°
AsCl5 4 1
AX4E
See-saw
<90, 180, 120°
SeCl4 3 2
AX3E2
T-shaped
<90, 180°
BrCl3
AX2E3
Linear
180°
XeF2

46. Molecular Geometries for Five Electron Pairs
AX AX4E AX3E AX2E3
Molecular Geometries for Five Electron Pairs
All based on trigonal bipyramid!

47. Molecular Geometry:6 Bonded Items
Total Bonds to C.A.
Bonded atoms (X)
Lone pairs (E)
AXE Notation
Molecular Geometry
Bond
Angles
Example 6
AX6
Octahedral
90, 180°
TeBr6 5 1
AX5E
Square Pyramid
<90, 180°
BrF5 4 2
AX4E2
Square Planar
<90°
XeF4 3
AX3E3
T-shaped
---
AX2E4
Linear
180°

48. Molecular Geometry: Six Bonded Items
AX6 AX5E AX4E2

49. Electronic Flashcards
Flash cards on molecular geometry and hybridization

50. Molecules with More than One Central Atom
Determine geometry for each central atom separately!
Example: In acetic acid, CH3COOH, there are three central atoms:

51. Molecules with only two atoms are always linear!
Examples: HCl N2

52. What shape (molecular geometry) would the following compounds have according to VSEPR theory?
CFCl3

53. What shape (molecular geometry) would the following compounds have according to VSEPR theory?
H2S
PBr5

54. What shape (molecular geometry) would the following compounds have according to VSEPR theory?
SeCl22-
C2H2
(hint: classify each C separately)

55. Bond Polarity Polar Bond Non-polar Bond Predicting Bond Polarity
Covalent bond in which the electrons are unequally shared
Ex.) O--H bond
Non-polar Bond
Covalent bond in which the electrons are equally shared
Ex.) F--F bond in F2
Predicting Bond Polarity
Use Electronegativity!! (see next slide)

56. Predicting Bond Polarity
Calculate the difference between the Pauling electronegativity values for the 2 elements
Type of Bond
IONIC
(COVALENT)
POLAR
NON-POLAR
Types of Atoms
1 metal & 1 nonmetal
(ex. NaCl)
(generally)
2 nonmetals
Ex. NH3, H2O
Ex. CCl4, O2
Electronegativity Difference
≥ 1.7
> 0.4 but < 1.7
≤ 0.4
0 – 0.4  Non-polar covalent
0.4 – 1.7  Polar covalent (more e/n element has greater pull)
1.7 and up  Ionic (e- are transferred between atoms)

57. Bond Polarity What type of bonds are these? Cl—Na O—F
E/N 

58. Polar molecules are NOT the same as polar bonds!
Polar Molecules and Dipole Moments
Polar molecules are NOT the same as polar bonds!
Can’t use ∆ E.N. to calculate if something is a polar molecule!

59. Polar Molecules (a.k.a dipole)
Molecule with separate centers of (+) and (-) charge
Molecules are polar if the pull in any one direction is not balanced out by an equal & opposite pull in the opposite direction
Positive and negative regions shown by “delta”(δ+ or δ-)
Ex. CH3Cl

60. Dipole Moment
Polar molecules have a DIPOLE MOMENT; will align with an electric field.

61. Determining Molecular Polarity
Nonpolar Polar

62. Determining the Polarity of a Molecule
See polarity flowchart!!
Shape is crucial (you must determine the VSEPR shape first)
All non-polar bonds = nonpolar molecule
Polar bonds  see if they cancel each other out
If they all cancel = nonpolar molecule
If they are unbalanced = polar molecule

63. Rules for Determining Polarity
4/28/2017
Simple Molecules:
Rules for Determining Polarity
A molecule is polar if:
It has only 2 atoms in it and both are different.
It has 3 or more atoms and has lone pairs on the central atom (i.e., it is classified as an AXE)
**Exception: where the lone pairs are symmetrical to the axis of bonded atoms in AX4E2 or AX2E3.**
3. It has 3 or more atoms in an AXn classification and all of the X’s are not the same atom.
Dr. Mihelcic Honors Chemistry

64. Determining Molecular Polarity
Nonpolar Polar
(pulls cancel) (pulls don’t cancel)

65. Example: Determine the polarity of the following molecules:
BF3 BF2Cl CO2
H2O H3O XeF4
PF5

66. Hybridization of Atomic Orbitals
The solutions of the Schrodinger equation led to these atomic orbitals:
1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f, etc.
However, overlap of these orbitals does not give a satisfactory explanation. In order to explain bonding, these orbitals are combined to form new sets of orbitals – this method is called hybridization.

67. Need for Hybridization (Carbon example)
Should be able to form only 2 bonds (since 2 unpaired electrons)
Combines s & p orbitals to form 4 sp3 hybrid orbitals (which now leaves 4 unpaired electrons that can make 4 stable bonds)

68. Hybridization of Atomic Orbitals
sp 2 sp hybrid orbitals from mixing of a s and a p orbital
sp 2 3 sp2 hybrid orbitals from mixing of a s and 2 p orbitals
sp3 4 sp3 hybrid orbitals from mixing of a s and 3 p orbitals
sp3d 5 sp3d hybrid orbitals from mixing of a s and 3 p and a d orbital
sp3d 2 6 sp3d2 hybrid orbitals from mixing of a s and 3 p and 2 d orbitals

69. Hybrid Orbitals Two atomic orbitals produce two hybrid orbitals
One s + one p  two sp

70. Hybridization of Atomic Orbitals
sp 2 sp hybrid orbitals from mixing of a s and a p orbital
sp 2 3 sp2 hybrid orbitals from mixing of a s and 2 p orbitals
sp3 4 sp3 hybrid orbitals from mixing of a s and 3 p orbitals
sp3d 5 sp3d hybrid orbitals from mixing of a s and 3 p and a d orbital
sp3d sp3d2 hybrid orbitals from mixing of a s and 3 p and 2 d orbitals
Notice that there are five hybrid orbital types: they match up with the five electron pair geometries!!!

71. Hybridization of Atomic Orbitals
sp 2 sp hybrid orbitals from mixing of a s and a p orbital
sp 2 3 sp2 hybrid orbitals from mixing of a s and 2 p orbitals
sp3 4 sp3 hybrid orbitals from mixing of a s and 3 p orbitals
sp3d 5 sp3d hybrid orbitals from mixing of a s and 3 p and a d orbital
sp3d sp3d2 hybrid orbitals from mixing of a s and 3 p and 2 d orbitals
Superscripts on s, p, d added together =
Sum of X + E in the designation for the electron pair geometry (sum of bonded atoms and lone pair e-)

72. How to determine:
ALWAYS results from tetrahedral electron pair geometry
ONLY hybridization explains why C forms four equal bonds!
Link for animations:

73. Sample Problem OF2 NH4+ CO2 COCl2 XeF4
Predict the molecular geometry and hybridization of the central atom in the following compounds:
OF2 NH4+ CO2
COCl2 XeF4

74. Sigma and Pi Bonds Sigma (σ) bonds Pi bonds (π)
Electron density is located between the nuclei
Single bond
Pi bonds (π)
Electron density is located above and below or in front of and in back of the nuclei
One pair of a double bond is called pi (π)
Two pairs of a triple bond are called pi (π)

75. Sigma & Pi Bonds: Ethylene and Acetylene

76. Sigma/Pi Bonds Link for animations: