1. Ionic Nomenclature Recap
____________ is ALWAYS written first.
If monoatomic, use the name of the element.
If polyatomic, use the name of the polyatomic ion.
____________ is ALWAYS written second.
If monoatomic, use -ide as the suffix.
Examples:
AgCl2
NaOH
NH4NO3

2. Recall
What 2 things are needed to create an IONIC bond: Therefore, what do you think creates a COVALENT bond:

3. Covalent Bonding

4. Covalent Bonds What happens to electrons in covalent bond?
Electrons ________ between atoms, not given or taken completely.
Atomic orbitals are combined to form molecular orbitals.
Generally bond between ______________ and _____________.
Why? 3 trends need to be taken into account.
Electron affinity, Ionization energy, and electronegativity differences between atoms are small.
Differences in trends aren’t large enough for one atom to completely take away electrons from the other, so electrons are shared.
Example: CH4

5. Identification of Bond Types Examples
1.) Na—Cl
11.) Al—Cl
2.) Mg—Cl
12.) N—O
3.) C—N
13.) Be—F
4.) O—Cl
14.) S—C
5.) Zn—Br
15.) Zn—N
6.) Ti—F
16.) I—I
7.) P—C
17.) P—O
8.) K—O
18.) Li—N
9.) F—C
19) C—Cl
10.) Li—I
20.) Ag—I

6. Covalent Compounds—Nomenclature
Rules—Binary Compounds
Element with smaller group # is always given first (similar to cation in ionic bonding).
Second element combines prefix with suffix –ide.
If second element begins with vowel, the –o- or –a- in the prefix is dropped.
Do not simplify subscripts in covalent compounds
Example— P4O10
Prefixes:
Mono = 1
Di = 2
Tri = 3
Tetra = 4
Penta = 5
Hexa = 6
Hepta = 7
Octa = 8
Nona = 9
Deca = 10
Prefix indicating number of atoms contributed by second element
Prefix needed if first element contributes more than one atom
Root name of second element + ide
Name of first element

7. Covalent Compounds—Nomenclature
As2O5
Carbon Tetrafluoride CO
Sulfur Trioxide

8. BELLRINGER (5 Minutes) SiI4 NO OF2 tetraphosphorus decoxide
dinitrogen tetroxide

9. Acids
2 types:
Binary Acids—contain H and another element, usually a halogen.
1.) Put hydro- as prefix for H.
2.) Use –ic as suffix for second element.
Example—HCl = hydrochloric acid
Oxyacids—contain H, O, and a third element (mostly H paired with a polyatomic ion).
1.) Identify anion.
2.) if the anion suffix is –ate replace it with –ic
3.) if the anion suffix is –ite replace it with –ous
Example—HNO3 = Nitric Acid

10. Practice in your notes…
H2SO3
H3PO4
Hypochlorous Acid
Arsenic Acid

11. Whiteboarding SiI4 NO HF OF2 Carbonic acid hydrosulfuric acid
tetraphosphorus decoxide
dinitrogen tetroxide

12. Hydrocarbons Hydrocarbons = the simplest organic compounds
Contain only __________ and _________________
3 types of straight-chain hydrocarbons
Alkanes—completely saturated hydrocarbons (no double or triple bonds)
Go by the formula CnH2n+2
End in suffix -ane
Alkenes—contain a double bond
End in suffix -ene
Alkynes—contain a triple bond
End in suffix -yne

13. Need to know first ten alkane hydrocarbons…
CH4 Methane
C6H14 Hexane
C2H6 Ethane
C7H16 Heptane
C3H8 Propane
C8H18 Octane
C4H10 Butane
C9H20 Nonane
C5H12 Pentane
C10H22 Decane

14. Bellringer (5 min): Determine type of bond that will be formed
C – O
Al – Cl
K – Br
F – F
Mn – O
Ca – Se
S – Cl
Li – S
P – Br
C – I

15. Formation of Covalent Bond
Nature favors chemical bonding…why?
Makes the atoms more _______________.
potential energy is _________________ when the atoms are bonded. Lower ___________ = more _____________.

16. Characteristics of a Covalent Bond
__________________ —the average distance between two bonded atoms.
Depends on type of bond: single, double, or triple.
Bond _____________ Energy —The energy required to break 1 mol of a specific chemical bond (always _______________).
Positive energy to break a bond (endothermic)
Negative energy to make a bond (exothermic)
Indicates strength of a bond and is directly related to bond length.

17. In Reactions…
For endothermic reactions– a greater amount of energy is required to break the existing bonds than is released when the new product bonds form.
For exothermic reactions – more energy is released forming new bonds than is required to break bonds of the reactants.

18. Molecule—a neutral group of atoms that are held together by covalent bonds.
Molecular formula—shows the types and numbers of atoms combined in a single molecule of a compound.
Example: CH4

19. Octet Rule
Atoms undergo bonding in order to satisfy the ______________.
Octet rule: Chemical compounds tend to form so that each atom has ___ electrons, either by ____________, _______________, or _____________ electrons. Atoms want to become “Noble-gas-like” with filled valence shells.
Diatomic molecule: A molecule in which there are only two atoms.
F2, Cl2, Br2, I2, H2, O2, N2

20. Exceptions to Octet Rule
Hydrogen—forms only one bond to have two valence electrons.
Group 13—Has three valence electrons. Tends to form three bonds.
Some elements can form an expanded octet if bound to highly electronegative atoms.
Example: SF6
Expanded octet involves empty d orbitals to fit extra electrons.

21. Lewis Structures What are they?? What are they used for??
A formula where atomic symbols represent nuclei and inner shell electrons, and dot pairs represent valence and bonded electrons.
What are they used for??
Gives us a way to visualize bonding between atoms
Represents where electrons are located
Gives relative bond strengths to establish reactivity of molecules.

22. Lewis Structures Six Steps: 1.) Determine types of atoms in molecule.
Example: CH3I
2.) Write electron dot notation for each atom.
3.) Determine the total number of valence electrons available.
4.) Arrange atoms with LEAST electronegative atom in the center (exception: H), and place one shared pair of electrons between each of the atoms.
5.) Fill in valence shells of atoms with unshared electrons (lone pairs).
6.) Count electrons to make sure all available valence electrons are accounted for.
***** Add all atoms (except H) first, then fill in molecule with Hydrogens

23. Practice with Lewis Structures
Draw the Lewis structures for the following molecules:
NF3
BH3
PCl5

24. Whiteboarding Practice
Draw the Lewis structures for the following molecules:
SiH Pl3
H2S CCl2F2
C2F SF6

25. Before we go any further…
Try the Lewis structure for ions:
Negative charge: add the electrons
Positive charge: remove the electrons
Ammonium Peroxide
Try the Lewis structure for C2H4.

26. Single and Multiple Covalent Bonds
Single bond – two electrons shared
Also called sigma bond, or  bond.
Examples: H2
Double bond – four electrons shared
Also called pi bond, or  bond.
Example: O2
Triple bond – six electrons shared
Consists of one  and one  bond.
Example: N2

27. Lewis structures with multiple bonds
Multiple bonds become evident in lewis structures when there are not enough valence electrons after adding lone pairs.
Examples:
CH2O
CO2
HCN

28. Relative Bond Lengths and Strengths
Bond length—the more shared pairs = the ___________ the bond.
Single > Double > Triple
Bond Strength—the more shared pairs = the __________ the bond.
Triple > Double > Single

29. Hydrocarbons Hydrocarbons = the simplest organic compounds
Contain only __________ and _________________
3 types of straight-chain hydrocarbons
Alkanes—completely saturated hydrocarbons (no double or triple bonds)
Go by the formula CnH2n+2
End in suffix -ane
Alkenes—contain a double bond
End in suffix -ene
Alkynes—contain a triple bond
End in suffix -yne

30. Need to know first ten alkane hydrocarbons…
CH4 Methane
C6H14 Hexane
C2H6 Ethane
C7H16 Heptane
C3H8 Propane
C8H18 Octane
C4H10 Butane
C9H20 Nonane
C5H12 Pentane
C10H22 Decane

31. Lewis Structure Bellringer
H2O
NH3
PCl5
BF3
SF6
CH4
CS2
HCl

32. Resonance Structures
In molecules with multiple bonds, electrons are ___________________ due to shared orbitals.
Since electrons are delocalized, a single lewis structure cannot account for all of the possible locations of electrons.
Only electron locations vary, NEVER atom arrangement.

33. Formal Charges
Formal Charges are used when there is more than one possible Lewis structure for a molecule (resonance structures).
They help us determine with Lewis structure will most likely exist in the real world.
The most stable Lewis structure is one where all charges equal zero or the negative charge resides on the most electronegative element.
FORMAL = VALENCE NON-BONDING - BONDING ELECTRONS
CHARGE ELECTRONS ELECTRONS
Example 1: Methane
Example 2: Sulfur dioxide

34. Polarity
Based on electronegativity differences.
_______ molecule = uneven distribution of electrons (non-symmetrical)
_________________ = even distribution (symmetrical)

35. Polar Covalent Bonds

36. Non-polar and Polar Covalent Bonds
Dipole—a molecule that contains both positively and negatively charged regions (_______________ sharing of electrons).
Nonpolar bonds do not contain a dipole (similar electronegativies).
Least Most
electronegative electronegative

37. Identify the dipole in the following molecules… and indicate if molecule is polar or non-polar
If geometry is symmetrical (NO lone pairs), dipoles will cancel
NH3
CH3F
SF6
AlCl3
CCl4

38. Identify the bond as polar or nonpolar
Remember:
C—H
Polar Covalent =
O—Cl
Nonpolar Covalent = 0-0.4
C—Cl
H—N
B—F
F—F