1. Covalent Bonding Chapter 8 Chapter 8.1 Vocab Covalent Bond Molecule
Lewis Structure
Sigma Bond
Pi Bond
Endothermic
Exothermic
Covalent Bonding
Chapter 8

2. Ch. 8 Vocabulary Covalent bond Exothermic Reaction Molecule
Structural Formula (Ch. 8.3)
Lewis Structure
Polar Covalent Bond
Sigma bond
Pi bond
Bond Dissociation Energy
Endothermic Reaction

3. OBJECTIVES Draw Lewis Structures of Covalent compounds
Name covalent compounds
Write formulas for covalent compounds
Describe characteristics of covalent molecules
Review electronegativity and compare they types of bonds of different molecules using the electronegativity scale

4. Covalent Bonding Remember the Octet Rule:
Similar arrangement of valence electrons
Electron arrangements determines chemical properties
Presents a model of chemical stability
RULE: Atoms become stable by having 8 electrons in their outer energy level (2 for smaller atoms)
When they have gotten 8 electrons they have achieved NOBLE GAS CONFIGURATION (NGC)
One way to get NGC is by transferring electrons as in an ionic bond.

5. Covalent Bonding The other way to achieve Noble Gas Configuration:
Sharing Electrons
Take the case of Water (H2O)
Hydrogen can’t transfer its electron, otherwise it would be just a proton and not a noble gas configuration.
It also can’t gain one. Hydrogen and oxygen can’t both gain electrons.

6. Sharing of Electrons
In the case of NaCl, chlorine has a much stronger affinity for electrons and sodium holds its valence electron very weakly.
In the case of H2O, both hydrogen and oxygen have similar affinities.
In other words, the attraction for electrons is not strong enough.
So what do they do?

7. Molecular Elements Molecules can vary greatly in size
Can be just two atoms (CO) to thousands or millions of atoms (DNA)
:CO:
Two or more atoms of the same element can form a covalent bond – this is called a molecular element.

8. Molecular Elements H-H NN O=O F-F Cl-Cl Br-Br Diatomic Molecules –
I -I
They are:
Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, and Iodine
Their Formulas are:
H2, N2, O2, F2, Cl2, Br2, and I2
All are gases except Br2 (liquid) and I2 (Solid)
Show electron dot configuration for H2, Cl2 and F2 for single bond
And O2 for double bond and N2 for triple bond.
There is a blank slide next slide to do this.
Diatomic Molecules –
Seven non metal elements are found naturally as molecular elements of two identical atoms

9. Molecular Elements
Allotropes
Examples:

10. Molecular Elements White Black Allotropes of Red Phosphorus
White is most common. Turns yellow when exposed to sunlight. It ignites at 93F and needs to be stored under water.
White can be converted to red phosphorus by heating (230 to 300 C)in the absence of air.
White can be converted to black phosphorus by heating at 200 C in the presence of air.
Other Examples: Diamond and graphite (both carbon)

11. Electron Dot Structures
(Lewis Symbols)
Remember (Chapter 5) electron dot structures (Lewis symbols).
The number of electrons available for bonding are indicated by unpaired dots.
Place the electrons one four sides of a square around the element symbol, e.g. ·C·
Lewis Structures – quick method – Pair up the single electrons in each atom to form a bond.
See Chemistry: Matter and Change page 212.
The quick-and-dirty method noted in the 4th bullet works best for binary compounds (O2, N2, CH4, NH3, etc). For more complicated molecules we’ll have a formal set of rules to use.

12. Electron Dot Structures
(Lewis Symbols)
Review of electron dot structures

13. Single Covalent Bonds
Chlorine
forms a
covalent
bond
with
itself
Cl2

14. How
will
two
chlorine
atoms
react? Cl Cl

15. Cl Cl
octet

16. Cl Cl The octet is achieved by each atom sharing the
electron pair in the middle

17. Cl Cl It is called a SINGLE BOND also called a sigma (σ) bond
Lone pair Cl Cl
It is called a SINGLE BOND
also called a sigma (σ) bond
Lone Pairs are non-bonding electrons or unshared electrons
The Cl2 molecule has how many lone pairs? 6

18. Cl Cl Cl2 Single bonds are abbreviated with a dash
This is the chlorine molecule,
Cl2
Contrast this with ionic bonding - next slide

19. NaCl - + Na Cl Cl2 Cl Cl This is the formation of an ionic bond.
electron transfer
and the formation of ions
Cl2
This is the formation of a covalent bond. Cl Cl
sharing of a pair of electrons
and the formation of molecules

20. Covalent Bonding
Electron cloud a “glue: to hold the two atoms together.
Maximizes the electron density between the two atoms.

21. Molecular Lewis Structures – Other Single Covalent Bonds
Show the Lewis structures of these on the board.
Also do on board Practice Problems 1-5 on P. 244.
Bonding pairs ?
Lone pairs?

22. Multiple Covalent Bonds
Oxygen is also one of the diatomic molecules

23. O
Each atom has two unpaired electrons

24. O Oxygen atoms are highly electronegative.
So both atoms want to gain two electrons.

25. Both electron pairs are shared.

26. O O
6 valence electrons
plus 2 shared electrons
= full octet

27. (a sigma (σ) and a pi (π) bond)
two bonding pairs,
making a double bond
(a sigma (σ) and a pi (π) bond)

28. O O = For convenience, the double bond can be shown as two dashes.
How many bonding pairs in O2?
How many lone pairs in O2?

29. This is the oxygen molecule,=
This is the oxygen molecule, O2

30. σ and π Bonds Single bond (Cl2) has one σ bond
Bond directly between the atoms
Double bond (O2 and C2H4) have two bonds –one σ and one π
σ bond goes directly between the atoms
π bond goes above the axis

31. σ and π Bonds in a Molecule
Do other N2, and some other halogens.
Ask how many sigma and how many pi bonds are in this molecule.
After this slide give out Covalent Compounds WS 2 – Structures and Bonds

32. Strength of Covalent Bonds
The energy required to break a covalent bond is called the bond dissociation energy.
As number of bonds increase, length gets shorter and bond energy increases.
So, H-H bonds are fairly easy to break and makes hydrogen gas reactive.
Nitrogen has a triple bond, which takes a lot of energy to break and is pretty much inert.
Use rubber band analogy

33. Strength of Covalent Bonds
Exothermic reactions release energy. They occur when the bonds of the products are more stable than the bonds of the reactants. In other words, energy is released forming the new bonds in the products.
Example: CH4 + O2  H2O + CO2 + energy
Endothermic reactions absorb energy. The occur when bonds of the reactants are more stable than the bonds of the products. In other words, energy has to be entered to make the products.
Example: 2H2O(l) + energy  2H2(g) + O2(g)
This might be a good place to do the yeast/H2O2/dish soap experiment to show exothermic reactions.
Can also do NH4NO3 or something like that for endo.

34. Chapter 8.2 - Formulas and Names
Rules for Naming Binary Inorganic Compounds
Write out the name of the first nonmetal (Left most or bottom most first)
Follow it by the name of the second nonmetal and end in –ide
Add a prefix to the name of each element to denote how many are present.
Example: SO2 is sulfur dioxide because sulfur is below oxygen in the same group.
Start of 8.2
8.2 Vocab
oxyacid
Organic compounds – Contain C-H bonds
Inorganic Compounds – do not contain C-H bonds
The same prefixes we used to denote how many waters were present in the last section.

35. Prefixes for Molecular Compounds
Number of Atoms
Prefix 1
Mono- 2
Di- 3
Tri- 4
Tetra- 5
Penta- 6
Hexa- 7
Hepta- 8
Octa- 9
Nona- 10
Deca-
Note: Same as prefixes used for hydrates.

36. Formulas and Names Rules for Naming Binary Inorganic Compounds
Omit mono- if there is only one atom of the first element
If o-o or a-o vowels appear next to each other, the first of the pair is omitted for easier pronunciation.
Example: NO is Nitrogen Monoxide rather than mononitrogen monooxide

37. Practice Formulas and NamesNO
NO2
N2O
N2O5
Give out a copy of that ‘dihydrogen monoxide’ danger sheet that I got in the science workshop while they are working on these problems.

38. More Practice Formula or Name Name or Formula CCl4 CO
Diarsenic Trioxide
Sulfur hexafluoride
P2O5

39. Common Names Some molecules are so common they have common names.
What is the common name of dihydrogen monoxide?
Another common compound is NH3. What is this commonly called?
Give out a copy of that ‘dihydrogen monoxide’ danger sheet that I got in the science workshop.

40. Common Names of Compounds
Acids
HCl – Hydrochloric Acid
H2SO4 – Sulfuric Acid
H3PO4 – Phosphoric Acid
HNO3 – Nitric Acid
HC2H3O2 – Acetic Acid (vinegar)

41. Common Names of Compounds
Bases
NaOH – Sodium Hydroxide
KOH – Potassium Hydroxide
NH3 - Ammonia

42. Organic Compounds Organic Compounds Made up of carbon
Carbon is able to bond with other carbon atoms to form long chains, rings, sheets, and larger networks
Carbon forms 4 covalent bonds
The simplest with 4 hydrogens to form methane (CH4)

43. Organic Compounds
Hydrocarbons – compounds containing just hydrogen and carbon
The first was methane (CH4)
The names of others change as more carbons are added.

44. Organic Compounds Formula Name #of Carbons CH4 Methane 1 C2H6 Ethane 2
C3H8 Propane 3
C4H10 Butane 4
C5H12 Pentane 5
C6H14 Hexane 6
C7H16 Heptane 7
C8H18 Octane 8
C9H20 Nonane 9
C10H22 Decane 10
After this slide give out Covalent Bonding WS 3 – Bond Strength and Nomenclature

45. Ch. 8.3 - Molecular Structures
Structural Formulas use letter symbols and bond symbols to show relative positions of atoms.
You have already done some like H-H, O=O, N≡N, HCl, and CCl4
What happens when there are more than two types of atoms and/or lone pairs are involved?
Start of 8.3

46. Apply to CO2 Central atom: C Total number valence electrons: 16
(4 from C and 6 each from O)
Draw skeleton structure: O-C-O
Place electrons on Os – form octet
All atoms have octet?
Draw final structure
Note the first step is to place at least 2 electrons between atoms—we know a bond must exist there.
After showing skeleton structure, do the rest on the board.
Work out Ammonia, methane, oxygen, ethanol out on the board

47. Try SO2 (neutral molecule)
Central atom
Total number valence electrons
Draw skeleton structure
Place electrons on Os – form octet
Place remaining electrons on S
All atoms have octet?
No? need to form multiple bonds
Draw final structure 18
O – S - O
O = S -O and satisfy octet rule.

48. Try ClO4- (anion) Central atom Total valence electrons
Skeleton structure

49. ClO4- (anion) Place remaining electrons on Os
Place remaining electrons on Cl.
All atoms have an octet?
Draw ClO4- on board. Cl surrounded by single bonds to 4 oxygens, each with 6 lone pair electrons around them.

50. Ch Molecular Shapes
The shape of a molecule determines many of its physical and chemical properties.
Molecular geometry (shape) can be determined with the Valence Shell Electron Pair Repulsion model, or VSEPR model which minimizes the repulsion of shared and unshared atoms around the central atom.
I’m going to cover shapes very basically. Really just go through linear, trigonal planar, tetrahedral and its offshoots (pyramidal and bent), trigonal bipyramid (central atom with 5 terminal) and octahedral (central with 6 atoms).
I’m not going to discuss offshoots of trigonal bipyramidal nor octahedral (i.e. square pyramidal and square planar).
Also, I’m not going to discuss hybridization.

51. Molecular Shapes
Electron pairs repel each other and cause molecules to be in fixed positions relative to each other.
Unshared electron pairs also determine the shape of a molecule.
Electron pairs are located in a molecule as far apart as they can be.

54. Practice Other examples from PhET.
Draw the Lewis Structures and note the shapes of the following:
NF3
CS2
BH3
NH4+
SO42-
These are the practice problems on p. 253.
This site also shows the shapes, both ball-and-stick and space filling.

55. Exceptions to the Octet Rule
There are three classes of exceptions to the octet rule:
Molecules with an odd number of electrons;
Molecules in which one atom has less than an octet;
Molecules in which one atom has more than an octet.
Odd Number of Electrons
Few examples. Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons.

56. Exceptions to the Octet Rule
Less than an Octet
Relatively rare.
Molecules with less than an octet are typical for compounds of Groups 1, 2, and 13 (3A).
Most typical example is BF3.

57. Exceptions to the Octet Rule
More than an Octet
This is the largest class of exceptions.
Atoms from the 3rd period on down can accommodate more than an octet.
From the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.
Example: SF6
Note: Raise screen and draw SF6. Sulfur surrounded by 6 flourines in an octahedron. Have to account for 48 electrons.
Other examples: PCl5 (Trigonal bipyramid) – 40 electrons total
ClF3 (Cl with three bonds to F’s and 2 sets of lone pairs) – 28 e-

58. Ch. 8.5 – Electronegativity and Polarity
Atoms form bonds to increase their stability.
They acquire an octet of electrons and get noble gas configuration.
The types of bonds we looked at so for are...
Ionic and Covalent

59. Electronegativity and Polarity
Bonding involves a sharing of electrons.
Some are shared equally, slightly or none at all.
Ionic bonds – sharing is so unequal that it is considered a transfer of electrons
Covalent – Pure covalent bonds are shared equally, such as with O2 or Cl2
Most bonds fall in between.

60. Electronegativity
Electronegativity – a measure of the ability of an atom in a bond to attract electrons.
You can make a comparison between two atoms.
The difference in electronegativity (ΔEN, pronounced ‘delta E N’) tells you the type of bond.
Note to them that the delta (Δ) means “difference”.

61. Electronegativity
The higher the number, the more electronegative the atom is, i.e. the more it wants the electrons.
Electronegativity increases across a period. Why?
Protons increase in the nucleus, the same energy level being filled and the pull of the nucleus increases the attraction for the valence electrons. (Higher Z* across a period)

62. Electronegativity LiF NaCl CsF KBr Which one is “most” ionic?
Use your Electronegativity table to find ΔEN for the following ionic compounds.
LiF
NaCl
CsF
KBr
Which one is “most” ionic?
Which one is “least” ionic?
ΔEN = 3.0
ΔEN = 2.1
ΔEN = 3.3
ΔEN = 2.0
-Also have them define the bond by the Delta E number.
-The greater the difference, the more ionic the character of the bond.
So most ionic is CsF
Least ionic is KBr
The dividing line in the figure to the right are somewhat arbitrary. Different texts might have different numbers. Sometimes the differences vary between compound and compound with the same difference. This may be due to the size of the ions, reactivity, and other factors.

63. Electronegativity
What happens when the electronegativity numbers are similar?
Then the difference between them is very small.
The bond is then more described as sharing between the atoms and the bond is covalent.

64. Electronegativity What happens when the electronegativity is the same?
ΔEN = 0 and the bond is described as pure covalent bond.
Ex: The F-F bond. EN for F = 4.0 and 4.0 for the other F. So, ΔEN = 0.
All electrons are shared equally for the other diatomic molecules.
All other diatomics, of course, being Cl2, Br2, I2, O2, N2, and H2.
Mostly covalent on the scale is from 0.0 to 0.5.

65. Electronegativity
What happens when the ΔEN is greater than 0 and ≤0.5?
The bonds are still considered covalent. The compounds they make tend to be gases or low-boiling liquids at room temperature.
Example: Find ΔEN for C-H
C EN =2.5; H EN =2.1 ; ΔEN = 0.4
The molecules CH4, C2H6, C3H8 are gases at RT.

66. Electronegativity
So, ionic bonds are on our scale are between ΔEN≥ 2.0 and ΔEN=3.3.
(Mostly) Covalent bonds are between ΔEN=0.0 and ΔEN≤0.5.
(Pure Covalent bonds are ΔEN=0.0)
What happens between 0.5 and 2.0?

67. Electronegativity
Between ΔEN > 0.5 and ΔEN < 2.0, there is a partial transfer of the shared electrons to the more electronegative element.
The bond that forms is called polar covalent.
A polar covalent bond has some degree of ionic character.
The numbers are arbitrarily set and to be used as a guideline or rule of thumb. Sometimes it depends on the molecule itself and its properties before you’d classify it as one or the other. Other texts may set the dividing lines at different ΔEN’s.

68. Electronegativity Cl So what happens in a polar covalent bond?
The electron spends time around each element, but more time around the more electronegative element.
ΔEN= EN (Cl) – EN(H) = 3.0 – 2.1 = 0.9
The arrow points in the direction of the negative end of the bond.
Also note that the (δ – small delta) means “slightly” or “somewhat”, as in slightly positive or slightly negative. Cl δ- δ+ H

69. Bond Polarity Cl has a greater share in bonding electrons than does H.
HCl is POLAR because it has a positive end and a negative end. (difference in electronegativity)
Cl has a greater share in bonding electrons than does H.
Cl has slight negative charge (-d) and H has slight positive charge (+ d)

70. Practice – Calculate the ∆EN and classify the type of bond
Pair of Atoms
∆EN and Bond Type
Ca-S
Ba-O
C-Br
Ca-F
H-Br

71. Electronegativity What is the ΔEN for O-H bonds in water? O EN =
H EN =
Δ EN =
How does this affect the properties of water?
Maybe show Water – What is it? Video in Chemistry/Videos folder.
Put Fig. 9.1 in right of slide.
Quick Demo (p. 268 of TWE)
Order Reichardt’s Dye from Sigma-Aldrich (~$55/250 mg). This shows different solutions’ polarities by color.

72. Electronegativity
ΔEN=1.4
ΔEN=0.0
ΔEN=0.3
ΔEN=0.4