1. Chemical Bonds Ionic & Covalent Bonds Lewis Structures
Nonpolar and Polar Covalent Bonds
Electronegativity
Resonance Structures
Exceptions to the Octet Rule
Hybridization
Molecular Shapes

2. Noble gas configuration
The noble gases are noted for
their chemical stability and
existence as monatomic
molecules.
Except for helium,
they share a common
electron configuration
that is very stable.
This configuration has 8 valence-shell electrons.
valence e- He Ne Ar Kr Xe Rn

3. The octet rule
Atoms are most stable if they have a filled or empty outer layer of electrons.
Except for H and He, a filled layer contains 8 electrons - an octet.
Atoms will
gain or lose (ionic compounds)
share (covalent compounds)
electrons to make a filled
or empty outer layer.

4. Ionic bonds
Ionic bonds consist of the electrostatic attraction between positively charged and negatively charged ions.
Ionic bonds are commonly formed between reactive metals and nonmetals.
Cl-
Na+

5. Ions and the octet rule Na Na+ + e- group IA (1) e- + Cl Cl-
Simple ions are atoms that have gained or lost electrons to satisfy the octet rule.
They will typically form based on what requires the smallest gain or loss of electrons to complete an octet.
Na Na+ + e-
group IA (1)
e- + Cl Cl-
group VIIA (17)

6. Ionic compounds Don’t exist as individual molecules
Tend to form crystals
Ions touch many others
Formula represents the average ion ratio
Empirical formula is
the lowest whole
number ratio.
NaCl
1 to 1 ratio
sodium chloride

7. Metals Ion charges Group IA (1)
All elements have a 1+ charge also known as the oxidation state. Ion has same name as metal. No numbers are required in the name.
Group IIA (2)
All have a single oxidation state,2+ charge. Ion has same name as metal. No numbers are required in the name.
Group IIIA (13)
All have a 3+ oxidation state. (Tl also has a 1+ oxidation state.) Ion has same name as metal. (You must use numbers with thallium but not the other Group IIIA (13) metals.)

8. Metals with multiple charges
Ion name is metals with multiple oxidation states is the metal name followed by
Roman Numeral for charge in parentheses.
Group IV (14)
All metals and semimetals have oxidation states of 2+ and 4+.
Group V (15)
All metals and semimetals have oxidation states of 3+ and 5+.
Group VI (16)
Metals and semimetals have oxidation states of 4+ and 6+ except Po (+2 only).

9. Why does this happen? There d shell electrons exist
When a metal has electrons of d shell electrons, it is possible to lose the s and p electrons or d electrons or shift electrons to get empty outer shell.
This is why there are two or more possible oxidation states for many elements.

10. Transition metals
Listing the ones that only have a single common ionic charge is easy.
All Group III B - 3+
Ni, Zn, Cd - 2+ Ag
The ion name is the same name as the metal. No numbers are needed for these few.
The other elements are able to form two or more cations and require roman numerals in parentheses.

11. Examples of transition metals form two or more cations?
When a metal has d electrons, they can play a role in forming an ion.
Examples.
Fe loses two 4s electrons to form Fe2+. It loses two 4s and one 3d electron to form Fe3+.
Cu loses on 4s electron and moves the other 4s to the 3 d to form Cu+. It loses both 4s to form Cu2+.
It would be difficult for you to predict all the possible charges and stay within the scope of this class.

12. Lewis structures
This is a simple system to help keep track of electrons around atoms, ions and molecules - invented by G.N. Lewis.
If you know the number of electrons in the valence-shell of an atom, writing Lewis structures is easy.
Lewis structures are used primarily for s- and p-block elements.

13. X Lewis symbols Basic rules Draw the atomic symbol.
Treat each side as a box that can hold up to two electrons.
Count the electrons in the valence shell.
Start filling box - don’t make pairs unless you need to. X

14. O Lewis symbols Oxygen has 6 electrons in its valence - VIA.
Start putting them in
the boxes.

15. Lewis symbols

16. Lewis Symbols O
This is the Lewis
symbol for oxygen.

17. Li Be B C N O F Ne Lewis symbols
Lewis symbols of second period elements
Li Be B C
N O F Ne

18. Lewis dot formula and the formation of NaCl
Na + Cl Na + Cl
The electron from Na moves over to
the Cl. Now both satisfy the octet rule.
Na becomes Na+ - a cation
Cl becomes Cl- - an anion
The + and - charges attract each other
and form an ionic bond.

19. Lewis dot structures of covalent compounds
In covalent compounds atoms share electrons. We can use Lewis structures to help visualize the molecules.
Lewis structures
Multiple bonds must be considered.
Will help determine molecular geometry.
Will help explain polyatomic ions.

20. H Cl Types of electrons Bonding pairs
Two electrons that are shared between two atoms. A covalent bond.
Unshared pairs
A pair of electrons that are not shared between two atoms. Lone pairs or nonbonding electrons. oo
Unshared
pair
H Cl oo oo oo
Bonding pair

21. H H H H C H F F H Single covalent bonds
Do atoms (except H) have octets?

22. Nonpolar and polar covalent bonds
When two atoms share a pair of electrons equally.
H H Cl Cl
Polar
A covalent bond in which the electron pair in not shared equally.
H Cl
Note: A line can be used to represent a shared pair of electrons. oo oo oo oo oo oo oo oo oo d+ d- oo oo oo

23. Polar molecules
Electrons in a covalent bond are rarely shared equally.
Unequal sharing results in polar bonds. oo
H F oo oo oo
Slight positive side
Smaller electronegativity
Slight negative
Larger electronegativity

24. Electronegativity
The ability of an atom that is bonded to another atom or atoms to attract electrons to itself.
It is related to ionization energy and electron affinity.
It cannot be directly measured.
The values are unitless since they are relative to each other.
The values vary slightly from compound to compound but still provide useful qualitative predictions.

25. Electronegativities Electronegativity is a periodic property.
Atomic number

26. Electronegativity
Relative ability of atoms to attract electrons of bond. At
1.9 I
2.2 Br
2.7 Cl
2.8 Po
1.8 Te
2.0 Se
2.5 S
2.4 Bi
1.7 Sb As P
2.1 Pb
1.5 Sn Ge Si F
4.1 O
3.5 N
3.1 Tl
1.4 Na
1.0 Cs
0.9 Rb K Ba Mg
1.2 Sr Ca In Ga Al H Li Be B C

27. Electronegativity
The greater the difference in electronegativity between two bonded atoms, the more polar the bond.
If the difference is great enough, electrons are transferred from the less electronegative atom to the more electronegative one.
- Ionic bond.
Only if the two atoms have exactly the same electronegativity will the bond be nonpolar.

28. Electronegativity
Determine the difference in electronegativity between the bonded atoms in the following compounds.
KCl
H2O
CH4
NO2

29. Electronegativity
Determine the difference in electronegativity between the bonded atoms in the following compounds.
KCl ENK = 0.9 ENCl = 2.8 D = 1.9
H2O
CH4
NO2

30. Electronegativity
Determine the difference in electronegativity between the bonded atoms in the following compounds.
KCl ENK = 0.9 ENCl = 2.8 D = 1.9
H2O ENH = 2.2 ENo = 3.5 D = 1.3
CH4
NO2

31. Electronegativity
Determine the difference in electronegativity between the bonded atoms in the following compounds.
KCl ENK = 0.9 ENCl = 2.8 D = 1.9
H2O ENH = 2.2 ENo = 3.5 D = 1.3
CH4 ENC = 2.5 ENH = 2.2 D = 0.3
NO2

32. Electronegativity
Determine the difference in electronegativity between the bonded atoms in the following compounds.
KCl ENK = 0.9 ENCl = 2.8 D = 1.9
H2O ENH = 2.2 ENo = 3.5 D = 1.3
CH4 ENC = 2.5 ENH = 2.2 D = 0.3
NO2 ENN = 3.1 ENO = 3.5 D = 0.4

33. Properties of ionic and covalent compounds
Ionic compounds
Held together by electrostatic attraction
Exist as 3-D network of ions
Empirical formula is used
Covalent compounds
Discrete molecular units
Atoms held together by shared electron pairs
Formula represents atoms in a molecule

34. Drawing Lewis structures
Write the symbols for the elements in the correct structural order.
Calculate the number of valence electrons for all atoms in the compound.
Put a pair of electrons between each symbol, the bond between each.
Beginning with the outer atoms, place pairs of electrons around atoms until each has eight (except for hydrogen).
If an atom other than hydrogen has less than eight electrons, move unshared pairs to form multiple bonds.

35. C-O-O O-C-O Lewis structures Example CO2 Step 1
Draw any possible structures
C-O-O O-C-O
You may want to use lines for bonds.
Each line represents 2 electrons.

36. Lewis structures
Step 2
Determine the total number of valence electrons.
CO2 1 carbon x 4 electrons = 4
2 oxygen x 6 electrons = 12
Total electrons = 16

37. C O O Lewis structures Step 3
Try to satisfy the octet rule for each atom
- all electrons must be in pairs
- make multiple bonds as required
Try the C-O-O structure
No matter what you
try, there is no way
satisfy the octet for
all of the atoms.
C O O

38. O C O O=C=O Lewis structures = is a double bond,
This arrangement needs
too many electrons.
O C O
How about making some double bonds?
O=C=O
That works!
= is a double bond,
the same as 4 electrons

39. Ammonia, NH3 H H N H H H N H Step 1 Step 2 3 e- from H 5 e- from N
8 e- total
Step 3
N has octet
H has 2 electrons
(all it can hold) H
H N H

40. Multiple bonds So how do we know that multiple bonds really exist?
The bond energies and lengths differ!
Bond Bond Length Bond energy
type order pm kJ/mol
C C
C C
C C

41. Resonance structures O - S = O O = S - O
Sometimes we can have two or more equivalent Lewis structures for a molecule.
O - S = O O = S - O
They both - satisfy the octet rule
- have the same number of bonds
- have the same types of bonds
Which is right?

42. Resonance structures O - S = O O = S - O O S O They both are!
This results in an average of 1.5 bonds between each S and O.

43. Resonance structures
Benzene, C6H6, is another example of a compound for which resonance structure must be written.
All of the bonds are the same length. or

44. Exceptions to the octet rule
Not all compounds obey the octet rule.
Three types of exceptions
Species with more than eight electrons around an atom.
Species with fewer than eight electrons around an atom.
Species with an odd total number of electrons.

45. Atoms with more than eight electrons
Except for species that contain hydrogen, this is the most common type of exception.
For elements in the third period and beyond, the d orbitals can become involved in bonding.
Examples
5 electron pairs around P in PF5
5 electron pairs around S in SF4
6 electron pairs around S in SF6

46. An example: SO42- S O 1. Write a possible arrangement.
2. Total the electrons.
6 from S, 4 x 6 from O
add 2 for charge
total = 32
3. Spread the electrons
around. S O - ||

47. Atoms with fewer than eight electrons
Beryllium and boron will both form compounds where they have less than 8 electrons around them. : :
:F:B:F:
:F:
:Cl:Be:Cl: : : : : :

48. Atoms with fewer than eight electrons
Electron deficient. Species other than hydrogen and helium that have fewer than 8 valence electrons.
They are typically very reactive species. F | B
F - + H
:N - H
F H
| |
F - B - N - H

49. Species with an odd total number of electrons
A very few species exist where the total number of valence electrons is an odd number.
This must mean that there is an unpaired electron which is usually very reactive.
Radical - a species that has one or more unpaired electrons.
They are believed to play significant roles in aging and cancer.

50. Species with an odd total number of electrons
Example - NO
Nitrogen monoxide is an example of a compound with an odd number of electrons.
It is also known as nitric oxide.
It has a total of 11 valence electrons: six from oxygen and 5 from nitrogen.
The best Lewis structure for NO is:
:N::O: : .