1. Covalent Bonding and Lewis Structures

2. My Favorite “NO” #5
The diagram represents a mixture of S atoms and O2 molecules in a closed container.
Which diagram shows the results after the mixture reacts as completely as possible according to the equation:2S + 3O2 2SO3

3. My Favorite “NO” #5
2S + 3O2- 2SO3
My Favorite “NO” # D

4. Day 1
Bonds and Bond Types

5. Page 3: Watch the following videos and take notes!
Ionic Bond video click here.
Metallic Bond video: click here.
Covalent Bond video: click here.

6. Homework Pg 4

7. My Favorite “NO” #6
. The circle on the left shows a magnified view of a very small portion of liquid water in a closed container.
What would the magnified view show after the water evaporates?

8. My Favorite “NO” #6
My Favorite “NO” # E

9. Day 2
Lewis Dot Structures and Bond Polarity

10. Ionic Bonding Generally occurs between metals and nonmetals.
Can also occur with polyatomic ions.

11. Ionic Bonding
Involves the transfer of electrons, followed by electrostatic attraction.

12. Covalent Bonding Generally occurs between nonmetals.
Involves sharing of electrons, rather than transfer.

13. Octet Rule
Atoms will acquire, through sharing or transfer, the electron configuration of a noble gas. This happens in order for the atoms to gain stability.
Most noble gases
have 8 valence
electrons
He are the exception

14. Octet Rule Other Exceptions** These elements do not need a full octet:
H, He, Li, Be, B
Needs only 6 electrons to be stable
Needs only two electrons to be stable

15. Valence Electrons
The electrons in the highest occupied energy level. There are two ways to determine the number of valence electrons.
A. Using the electron configuration: 1s22s22p63s23p2, how many valence electrons are in this element?
2 + 2 = 4

16. Valence Electrons
B. Look at the group number to determine valence electrons

17. Dot Models
The number of dots is equal to the number of valence electrons. 1 2 3 4 5 6 7 8 P
Example: Phosphorus

18. Examples of Ionic compounds: Remember…..
NaCl MgBr2

19. Bonding in Covalent Molecules
Each atom in the molecule is connected by bonds.
Bonds are shared pairs of electrons, and they are represented by a dash.
Pairs that are not shared are called unshared electrons or lone pairs.

20. A single bond is created by one shared pair of electrons
A double bond is created by 2 shared pairs of electrons
A triple bond is created by 3 shared pairs of electrons.

21. Steps for Drawing Lewis Structures for Covalent compounds:
C has 4
4H has 4
C needs 8
H needs 2x4 = 8
Has 8
Need16
= 8
8/2 = 4 bonds

22. Covalent bonds dashes 1st first Hydrogen
Use the second element in the formula instead.
Subtract the pairs used in step 2
unshared elections
1 pair
octet of 8 valence electrons
aren’t enough pairs

23. Practice: Has Needs 2H = 2 2H= 4 O = 6 O= 8 8 12 12- 8 = 4/2 = 2
12- 8 = 4/2 = 2
2 bonds needed
Has Needs
H = 1 H=2
C = 4 C=8
P = 5 P=8
18-10 = 8/2 = 4
Has Needs
C = C = 8
2O =12 2O= 16
= 8/2 = 4
4 bonds needed
Has Needs
2N =10 2N=16
= 6/2 = 3

24. Remember: Exceptions to the octet rule: Atoms with less than an octet.
EX: BF3

25. Steps for Drawing Lewis Dot Structures:
Look at your chemical formula and arrange your atoms as they most likely be positioned using the following rules:
The first atom in your formula will be your central atom unless it is hydrogen. Hydrogen is never central – use the next element in the formula if hydrogen is listed first.
Example: CH4
H H C

26. H 2. Add your valence electrons to each atom.
3. Draw a skeleton of the molecule, connecting atoms in the molecule with single bonds (1 shared pair)
Make sure each atom is happy even if they have to share. Each atom wants to have eight electrons around it, two to each side of the box. Each bond, or stick, counts as two electrons for each atom. H C

27. Drawing CH4

28. Extra Practice:
H2O HCP
CO2 N2

29. Homework: Page 7
Optional…
check with
your teacher

30. My Favorite “NO” #7
True or False? When a match burns, some matter is destroyed.
A. True
B. False
Answer
misconception
False

31. Day 3 Notes Electron and Molecular Geometry
Once we have drawn the Lewis Structure, now we can describe how the atom looks in 3D.

32. Electron Geometry
The geometry based on the number of electron groups on the central atom.
It does not matter if the groups are a single bond, a multiple bond, or lone pairs.

33. Know the following Geometries
Electron Geometry
Linear
Trigonal Planar
Tetrahedral
Number of electron regions on the Central Atom: 2 3 4
Picture

34. Molecular Shape
In order to predict molecular shape, we assume the valence electrons repel each other. Therefore, the molecule adopts whichever 3D geometry minimizes this repulsion.
We call this process Valence Shell Electron Pair Repulsion (VSEPR) theory.

35. Know the following Molecular Geometries
Molecular Geometry
Linear
Trigonal Planar
Tetrahedral
Pyramidal
Bent
Number of Lone pairs on the Central Atom 1
Number of atoms bonded to the central atom 2 3 4
Bond Angle
180
120
109.5
107.5
<104.5
Picture

36. Day 3 Homework: page 9 1. 2. 3.

37. My Favorite “NO” #8
What is the reason for your answer to the question above? #7
a. This chemical reaction destroys matter.
b. Matter is consumed by the flame.
c. The mass of ash is less than the match it came from.
d. The atoms are not destroyed, they are only rearranged.
e. The match weighs less after burning.

38. My Favorite “NO” #8
What is the reason for your answer to the question above? #7
The other answers are all good misconceptions
d. The atoms are not destroyed, they are only rearranged.

39. Day 4 Notes: Bond Polarity
In polar bonds, shared pairs of electrons are pulled between the nuclei of atoms sharing them. Sometimes electrons are pulled equally and sometimes they are not. This has to do with electronegativity.
Recall: electronegativity is the ability to attract electrons.

40. Nonpolar Covalent Bonds
The electrons are shared equally.
EX: All diatomic elements are nonpolar.

41. Polar Covalent Bonds
The electrons are shared unequally due to electronegativity
The more electronegative atom will have a stronger attraction for the bonded electrons and will have a slightly negative charge.
The less electronegative atom will have a slightly positive charge.

42. Example - HCl
Look up the electronegativity values for hydrogen and chlorine.

44. Example: HCl Which is more electronegative? H: 2.1 Cl: 3.0
Chlorine has a slightly negative charge, while hydrogen has a slightly positive charge.

45. There are two ways to communicate the polarity of HCl:
 +
 -
The lowercase Greek letter delta shows that the atoms involved acquire only partial charges.
H – Cl
The arrow points to the more electronegative atom.

47. Example: Water
Is hydrogen or oxygen more electronegative?

48. Oxygen!!

49. O H H  -  +  + Example: Water
Is hydrogen or oxygen more electronegative? Oxygen!
 - O
H H
The O-H bonds are polar.
 +
 +

50. Electronegativity Difference
The difference in electronegativities indicates the type of bond the atoms will form.
Electronegativity Difference
Type of Bond
Example
Nonpolar
H-H (0.0)
Polar
H-Cl (0.9)
Very polar
H-F (1.9)
>2.0
Ionic
Na+Cl+ (2.1)
What type of bond will form between:
N and H?
F and F?
Ca and O?
Br and Cl?

51. 3.0 – 2.1 = .9

52. Electronegativity Difference
The difference in electronegativities indicates the type of bond the atoms will form.
Electronegativity Difference
Type of Bond
Example
Nonpolar
H-H (0.0)
Polar
H-Cl (0.9)
Very polar
H-F (1.9)
>2.0
Ionic
Na+Cl+ (2.1)
What type of bond will form between:
N and H?
F and F?
Ca and O?
Br and Cl?
Polar!

53. 4.0 – 4.0 = 0.0

54. Electronegativity Difference
The difference in electronegativities indicates the type of bond the atoms will form.
Electronegativity Difference
Type of Bond
Example
Nonpolar
H-H (0.0)
Polar
H-Cl (0.9)
Very polar
H-F (1.9)
>2.0
Ionic
Na+Cl+ (2.1)
What type of bond will form between:
N and H?
F and F?
Ca and O?
Br and Cl?
Polar!
Nonpolar!

55. 3.5 – 1.0 = 2.5

56. Electronegativity Difference
The difference in electronegativities indicates the type of bond the atoms will form.
Electronegativity Difference
Type of Bond
Example
Nonpolar
H-H (0.0)
Polar
H-Cl (0.9)
Very polar
H-F (1.9)
>2.0
Ionic
Na+Cl+ (2.1)
What type of bond will form between:
N and H?
F and F?
Ca and O?
Br and Cl?
Polar!
Nonpolar!
Ionic!

57. 3.0 – 2.8 = .2

58. Electronegativity Difference
The difference in electronegativities indicates the type of bond the atoms will form.
Electronegativity Difference
Type of Bond
Example
Nonpolar
H-H (0.0)
Polar
H-Cl (0.9)
Very polar
H-F (1.9)
>2.0
Ionic
Na+Cl+ (2.1)
What type of bond will form between:
N and H?
F and F?
Ca and O? Ionic!
Br and Cl?
Polar!
Nonpolar!

59. Summary

60. Day 4 Homework
Page 11
Polar and Nonpolar
..Ha Ha

61. Day #5 Review: Bond Polarity
In covalent bonds, shared pairs of electrons are pulled between the nuclei of atoms sharing them.
Sometimes electrons are pulled equally and sometimes they are not.
There are two types of covalent bonds:
Non-polar
Polar

62. Molecule Polarity (PAGE 12)
If a molecule has all nonpolar bonds then the molecule is nonpolar.
If a molecule has a polar bond then the whole molecule is usually polar, but not always.

63. To determine if a molecule is polar or nonpolar, look at the central atom.
If the central atom has no lone pairs and has all of the same type of atoms attached to it, then the molecule is nonpolar.
If the central atom has no lone pairs but different atoms attached to it, the molecule is polar.
If the central atom has lone pairs, the molecule is polar.
In other words: If the molecular geometry is symmetrical, the bond polarities cancel, and the molecule is nonpolar.

64. Remember your SYMMETRICAL shapes!
These molecules are symmetrical!!!

65. Example: carbon dioxide
CO2 has two polar bonds, but since the structure is linear (and symmetrical) the bonds cancel.
 +
 -
 -
O C O
Nonpolar molecule!

67. H H C H H  +  -  +  + Nonpolar molecule!  + Example: methane
CH4 has four polar bonds, but since the structure is tetrahedral (symmetrical) the bonds cancel.
 + H
 -
 +
H C H
 +
Nonpolar molecule! H
 +

68. O H H  -  +  + Example: Water Polar molecule
Water has a bent shape, due to unshared electrons (so NOT symmetrical)- The polarities do NOT cancel.
 -
Polar molecule O
H H
 +
 +

69. In a polar molecule, one end has a positive charge and the other has a negative charge. A molecule that has poles is called a dipolar molecule or a dipole.

70. Intermolecular Forces
Click to view video

71. Intermolecular Forces
Intermolecular Forces are the attractive forces between neighboring particles.
Dispersion Forces – occurs between
non-polar molecules, temporary partial charges

72. Dipole – dipole interaction – occurs when polar molecules are attracted to one another

73. Hydrogen bonding – occurs when a hydrogen atom involved in an extremely polar bond is strongly attracted to an adjacent molecule. This is the strongest intermolecular force.
Extremely polar bonds: H—N, H—O, H—Cl , and H—F
Hydrogen bonding

74. Day 5 Homework: Page 13 These forces are why: Comb bends water
Water beads up
Gecko sticks
to glass