1. Thursday 10.27.16 Agenda Review POGIL exercise Do Now
Covalent Bonding notes
LDD for compounds
Do Now
Pull out your POGIL packets and a scrap sheet of paper
Homework

2. Covalent Bonding

3. Bonding Two methods to gain or lose valence electrons:
Transfer of Electrons = Ionic Bonding
Sharing of Electrons = Covalent Bonding
Ionic Bonding - attracted to each other, but not fully committed
Covalent Bonding - fully committed, and shares everything

4. Why do metals have multiple charges?
Transition metals have 2 s electrons in their valence shells
They also have ‘high energy’ d electrons
These atoms typically lose their s electrons to form a +2 ion, but also sometimes lose some of their d electrons
This is why there are two possible oxidation states for many transition element

5. Naming Covalent Compounds
Prefixes used to denote number of atoms of element
The suffix -ide is added to the more electronegative element
For the least electronegative (on the left), mono- is not used to denote one atom.
The first vowel is often dropped to avoid the combination of “ao” or “oo”. CO = carbon monoxide (monooxide)

6. Molecules Molecule = covalently bonded atoms
Diatomic molecule = two atoms
Molecular compounds – lower melting and boiling point than ionic compounds
Molecular formula
Doesn’t have to be lowest whole-number ratio
Does represent structure

7. Molecular Representations

8. 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.

9. Types of electrons H Cl 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

10. Covalent Bonding
When two similar atoms bond, none of them wants to lose or gain electrons
Share pairs of electrons to each obtain noble gas e- configuration.
Each pair of shared electrons = one covalent bond
Unshared Pairs = Pairs of e- not shared by all atoms
Show unshared pairs as dots N H

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

12. Sharing of Electrons Water - H2O - Covalent Bonding O H
H = 1e-, O = 6e-
Each H shares their electron with O Þ O = 8e- = [Ne]
O shares 1e- with each H Þ H = 2e- = [He]
Use dashes or dots to show covalent bonds H O

13. Multiple Covalent Bonds
Elements can share more than two electrons - creates double and triple bonds
Carbon Dioxide (CO2) - 4 electrons shared between the carbon and the oxygens (double bond) O C O C

14. Triple Covalent Bonds C H C H
Triple bond = Six electrons shared between the carbon atoms
Ethyne (C2H2), a.k.a acetylene
Bond length decreases from single to triple bonds
Bond strength increases from single to triple bonds C H C H

15. Coordinate Covalent Bonds
Coordinate covalent bond = one atom contributes both bonding electrons
Polyatomic ions contain coordinate covalent bonds
O = 8e-
C = 6e- C O
O = 8e-
C = 8e- C O
O must contribute one of its pairs

16. 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.

17. Lewis structures C-O-O O-C-O 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.

18. 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

19. Lewis structures C O O Step 3 Try the C-O-O structure
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

20. Lewis structures O C O O=C=O = 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

21. 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

22. Chapter 8

23. Resonance Structures
Structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms = Resonance Structures
Example: Ozone has two identical bonds whereas the Lewis Structure requires one single (longer) and one double bond (shorter).

24. Resonance Structures
Resonance Structures

25. 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?

26. Resonance structures O - S = O O = S - O O S O
They both are! But really…neither are.
O - S = O O = S - O
O S O
This results in an average of 1.5 bonds between each S and O.

27. Exceptions to Octet Rule
Three classes of exceptions to the octet rule for molecules
Odd number of electrons
One atom has less than an octet
One atom has more than an octet
Odd Number of Electrons
Few examples.
Molecules such as ClO2, NO, and NO2 have an odd number of electrons.

28. Exceptions to Octet Rule
Less than an Octet
Relatively rare.
Typical for compounds of Groups 13.
Most typical BF3.
More than an Octet
Atoms from the 3rd period onwards can accommodate more than an octet.
The d-orbitals are low enough in energy to participate in bonding and accept extra electron density.

29. An example: SO42- S O 1. Write a possible 2. Total the electrons.
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 - |

30. 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: : : : : :

31. 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: 6 from oxygen and 5 from nitrogen.
The best Lewis structure for NO is:
:N::O: : .

32. Electronegativity
Electronegativity = Ability to attract electrons in a chemical bond
Decrease as you move down a group
Increase as you move from left to right across a period.
Decreasing Electro-negativity
Increasing Electro- negativity

33. Polar Covalent Bonds
Sharing of electrons in a covalent bond does not imply equal sharing of those electrons.
In some covalent bonds - electrons located closer to one atom than the other
Unequal sharing Þ polar bonds.
Sharing based on electronegativity of elements in bond

34. Electronegativities
Chapter 8

35. Electronegativity and Polarity
Difference in electronegativity between atoms is a gauge of bond polarity
Difference
Type of Bond
Electrons
0 – 0.4
Nonpolar Covalent
Equal sharing
0.4 – 1.0
Moderately Polar Covalent
Unequal sharing
1.0 – 2.0
Very Polar Covalent
» 3
Ionic
Transfer

36. 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

37. Polar Dipoles O H δ- δ+ δ+
The positive end (or pole) in a polar bond is represented + and the negative pole - = dipoles (partial charges)
Polar molecules placed between electric (+/-) plates – become aligned with plates δ- H O δ+ δ+

38. Attractions between Molecules
Intermolecular attractions weaker than ionic or covalent bonds
these are attractions BETWEEN molecules not inside molecules like ionic or covalent bonds (intramolecular)

39. Types of Intermolecular Attractions
van der Waals forces
Dipole interaction
δ+ end of one molecule attracted to δ- end of another
Dispersion forces
Caused by movements of e-
Hydrogen Bonds
Hydrogen bonded with very electronegative element is also weakly bonded to an unshared pair on another electronegative atom

40. Characteristics of Ionic and Covalent Compounds
Ionic Compound
Covalent Compound
Representative Unit
Formula Unit
Molecule
Bond Formation
Transfer of e-
Sharing of e-
Type of elements
Metal + Nonmetal
Nonmetals
Physical State
Solid
Solid, liquid, or gas
Melting Point
High (usually above 300°C)
Low (usually below 300°C)
Solubility in water
High
High to low
Electrical conductivity of aqueous solution
Good conductor
Poor to nonconducting