1. Covalent Bonding Day 1

2. Ionic Bonding vs Covalent Bonding
While we have been writing Ionic bonds as single formulas (ex. NaCl), there is no such thing as a single molecule of NaCl Ionic bonds keep repeating to form a crystal, with the cation and anion repeating again and again.
Ionic bonds keep repeating, forming a crystal, until the crystal breaks off. No such thing as 1 NaCl molecule.
Cation and Anion size

3. Ionic Bonding vs Covalent Bonding
The big difference between ionic and covalent is that covalent molecules are individual molecules
Is such thing as individual molecules with covalent bonds. Individual molecules are not chemically bonded to each other. Separate entities.

4. Covalent Bonds
Covalent bonds are formed when two or more atoms (usually non-metals) SHARE electrons Atoms share electrons instead of giving or taking because there is a small difference (or no difference) in electronegativity

5. Lewis Structures Shows the arrangement of valence electrons in a bond
Electrons are not transferred using arrows, but drawn in-between two atoms to show that they are shared
IONIC
COVALENT
Na* Cl*******  Na+ and Cl- which  NaCl whose electron configuration does Na+ have? Cl-?
Ca and Cl  CaCl2

6. Steps for Drawing Lewis Structures
Add up the total number of valence electrons from the formula.
Arrange the atoms so there is a central atom with the rest surrounding it.
Draw in single bonds between the atoms using 2 electrons per bond.
Now figure out how many electrons are left over. Fill in leftover electrons around each atom to give them each 8 electrons.
Having 8 Electrons is known as having a “full octet”
Hydrogen can only have 2 electrons- is still full
Pool together electrons, can think of it like a bank account

7. Example: CH4
Step 1: Add up the total number of valence electrons from the formula. CH4 = 1 C and 4 H’s C = 4 electrons H = 1 electron x 4 = 4 electrons Total = 8 Valence Electrons

8. Example: CH4
Step 2: Arrange the atoms so there is a central atom with the rest surrounding it.
This is usually the first atom in the formula
Hydrogen is NEVER a central atom H
H C H

9. Example: CH4
Step 3: Draw in single bonds between the atoms using 2 electrons per bond.

10. Example: CH4
Step 4: Figure out how many electrons are left over. Fill in leftover electrons around each atom to give them a full octet.
Started with 8 Electrons to work with
Used 8 Electrons for initial bonds
= 0 Electrons left
Don’t need to fill in any leftover electrons
Carbon has 8 electrons around it
Hydrogen has 2 electrons around it

11. Lewis Structures
If you have enough electrons to create all of the bonds and give every atom a full octet (2 for Hydrogen, 8 for the rest) with no electrons left over then you have the right structure.
Too few electrons– you will need a double or triple bond
Too many electrons– double check your math!

12. Practice
CF NH3 H2O

13. CCl4 CH2F2 HF PBr3 C2H6 H2 SiCl4 H2S

14. Covalent Bonds Day 2

15. Write the Lewis Structure for SO2
What’s the problem?

16. What if we don’t have enough?
If you do not have enough electrons, you need to add double bond(s) and/or triple bond(s)
Don’t go too fast – try ONE double bond first, and if that doesn’t work try another double bond or try a triple bond instead.
Let’s try SO2 again…

17. General Information
Always try a formula with single bonds first and count how many electrons you used…
If you have enough for all single bonds and dots you are fine
If you used 2 more electrons than you had available you need a double bond
If you used 4 more electrons then you had available you need 2 double bonds OR a triple bond
Hydrogen can only have 1 single bond, it will NEVER have a double bond or triple bond

18. Few more practice problems
GeO SiO2 C2H2Cl2

19. Polyatomic Ions
Still draw lewis structures the same way for polyatomic ions
The charge on the polyatomic ion gets added or subtracted to the total number of valence electrons you have to work with!
Polyatomic Ions with a negative charge: Add that many electrons to the total at the beginning
Add because negative charge means it has extra electrons
Example: SO Total Valence Electrons = 30
Valence Electrons: (add charge) +2
S = New Total = 32
O = 6 x 4 = 24

20. Polyatomic Ions
Polyatomic Ions with a positive charge: Subtract electrons from the beginning total.
Subtract because positive charge means it has lost electrons
Example: NH Total Valence Electrons = 9
Valence Electrons: (subtract charge) -1
N = New Total = 8
H = 1 x 4 = 4

21. Draw the lewis structure for… SO3 2-
Valence Electrons: S = O =