1. Chemical Bonding & Molecular Geometry
Unit 12
Chemical Bonding & Molecular Geometry

2. Definitions Chemical Bonds Force that holds atoms together
It’s all about the electrons (e-)
Electrons are attracted to positively charged nucleus of other atom

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 transferred from metal to nonmetal

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

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

6. Shared Electrons Complete ShellsF F

7. Types of Chemical Bonds
Metallic Bonds
Atoms are bonded to one another (not to other elements)
Positive ions in a “sea” of negative charge (e-)

8. Metallic Bonding Video

9. Definitions Octet rule (Rule of 8)
8 e- in the outer shell very stable
H2 and He want a “duet”
Electron configuration for duet = ns2
Electron configuration for octet = ns2 np6

10. Examples of Bonding Types
Ionic Bonding:
NaCl, K2S, Ca(NO3)2
Covalent Bonding
H2 , Cl2
Metallic Bonding
Cu, Ag

11. Lewis Dot Diagrams
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.

12. 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 . .

13. 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.

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

15. Writing Lewis Dots Structures (Ionic Compounds)
Lewis Dot Diagrams of Ionic Compounds
Ex. 1) NaCl
Ex. 2) MgF2

16. Lewis Representations of Ionic Structures
MgO
Li2O

17. Lewis Dot Diagrams for Covalent Compounds
A substance made up of atoms which are held together by covalent bonds is a covalent compound.
They are also called molecules.

18. Covalent Compounds and Lewis Dot Diagrams
Diagrams show bonds in a covalent compound and tells us how the atoms will combine
Shared e- = bonding e-
Non-shared e- = lone pair e- (a.k.a. non-bonding e-)
Ex. F2

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. (see hints)
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!
Carbon will always be 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. Covalent Compounds and Lewis Dot Diagrams
HF NH3
CCl NH4+
CO NO3- *****

26. Resonance Structures 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 (in double or triple bonds), NOT atoms.

27. 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”

28. Resonance Structure Example, NO3-
Draw the Lewis diagram for NO3- with all possible resonance structures.

29. 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!

30. 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
Gives molecule its shape

31. 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

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

33. To determine the molecular 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 number of X and E, assign the molecular geometry.
4. Multiple bonds count as one bonded pair!

34. Molecular Geometry Notation
10/2/2017
Molecular Geometry Notation
A: Central Atom
X: Bonded Atom
E: Non-bonding electron pair
(Lone pair e- on central atom)
Dr. Mihelcic Honors Chemistry

35. 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

36. 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-

37. 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

38. 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
In Accelerated, you are not responsible for knowing anything about geometries involving 5 or 6 bonded items (Trigonal bipyramidal & octahedral)

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

40. 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°
In Accelerated, you are not responsible for knowing anything about geometries involving 5 or 6 bonded items (Trigonal bipyramidal & octahedral)

41. Molecular Geometry: Six Bonded Items
AX6 AX5E AX4E2

42. Electronic Flashcards
Flash cards on molecular geometry and hybridization

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

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

45. To determine the molecular 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!

46. VSEPR Examples:
What shape would the following compounds have according to VSEPR theory?
CH4
CO2

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

48. (hint: classify each C separately)
What shape (molecular geometry) would the following compounds have according to VSEPR theory?
H2S
C2H2
(hint: classify each C separately)

49. Bond and Molecule Polarity
Polar Bond
Covalent bond in which the electrons are unequally shared
Ex. H2O
Non-polar Bond
Covalent bond in which the electrons are equally shared
Ex. F2 or CH4
Predicting Bond Polarity
Use Electronegativity!! (see next slide)

50. 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)

51. Polar Molecules Polar Molecules (dipole)
Molecule with separate centers of (+) and (-) charge
In other words, molecules are polar if the pull in any one direction is not balanced out by an equal & opposite pull in the opposite direction

52. Polar Bonds and Polar Molecules
Drawing Polar Molecules
Positive and Negative regions shown by “delta”(δ)
Ex. CH3Cl

53. Determining the Polarity of a Molecule
Shape is crucial (determine the VSEPR shape 1st)
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

54. Determining Molecular Polarity
Nonpolar Polar

55. Examples: Polar or non-polar?
Determine if the following molecules are polar or nonpolar.
H2S F2
H2O

56. Special Types of Bonding
Hydrogen Bonding
Force in which a hydrogen atom covalently bonded to a highly electronegative element (F, O, or N) is simultaneously attracted to a neighboring nonmetal atom

57. Hydrogen Bonding Elements that undergo H-bonding
Hydrogen bonding is FON!
(Fluorine, Oxygen, and Nitrogen)
Effects on Physical Properties
H2O is most notable example of H-bonds
Ice forms rigid, open structures
Increases volume upon freezing (floats)
Molecules w/ higher molar mass have lower BP than H2O