1. Chemical Bonding & Molecular Shapes

2. Electrons
The number of valence electrons determines many of the chemical properties an element has
Elements in the same group have similar chemical and physical properties because they have the same number of valence electrons

3. Electrons Group 1 = 1 valence electron Group 2 = 2 valence electrons
*skip the middle*
Group 13 = 3 valence electrons
Group 14 = 4 valence electrons
Group 15 = 5 valence electrons
Group 16 = 6 valence electrons
Group 17 = 7 valence electrons
Group 18 = 8 valence electrons

4. Lewis Dot Structures
Valence electrons can be represented as a series of dots surrounding an atomic symbol
This is called an electron-dot structure, or more commonly, a Lewis dot structure
Lewis dot structures immediately tell you two things about an element:
How many valence electrons the element has
How many of those electrons are paired

5. C B O Lewis Dot Structures
Completing Lewis dot structures must be done in a specific way in order to correctly show paired and unpaired electrons
After identifying the number of valence electrons an atom has:
Write down the element symbol.
Fill in the electrons ON EACH SIDE of the symbol until the correct number of electrons is reached
For example –
Carbon = 4 valence electrons
Boron = 5 valence electrons
Oxygen = 6 valence electrons B O

7. Paired & Unpaired Electrons
Paired valence electrons are relatively stable
Do not usually form chemical bonds
Referred to as nonbonding pairs
Unpaired valence electrons are much more likely to participate in chemical bonding
They become paired with electrons from another atom

8. Ions
Review: when an atom gains or loses electrons, there is a net electric charge
If electrons are lost, there are more protons than electrons—the ion’s charge is positive
Called a cation
Metals are more likely to lose electrons
If electrons are gained, there are more electrons than protons—the ion’s charge is negative
Called an anion
Nonmetals are more likely to gain electrons

9. Ions
Atoms tend to lose or gain electrons in order to have a full outermost shell (usually 8 electrons)
Noble gases already have 8 electrons in their outermost orbital, they are inert
Alkali metals (group 1) and Halogens (group 17) are very reactive because they are each only 1 electron away from having a full outermost shell

10. Ionic Bonding Ionic bonds results from a transfer of electrons
Happens between a metal and nonmetal
When a metal atom (wants to lose) is placed in contact with a nonmetal (wants to gain)—electrons are transferred and two oppositely charged ions are formed
The two ions are attracted to each other by the electric force of attraction

11. Ionic Bonding Compounds with ionic bonds are called ionic compounds
The properties of these compounds are completely different from the elements that form them

12. Naming Ionic Compounds
Identify the names of the elements in the compound
Identify which is a metal and which is a nonmetal
Write the name of the metal with a capital letter
Write the name of the nonmetal with a lowercase letter
Change the ending of the nonmetal to –ide
EXAMPLES
Sodium and Chorine
Lithium and Fluorine

13. Total positive charge = total negative charge
Ionic Compounds
Ions have different charges
Ionic compounds want to have an overall charge of zero
Total positive charge = total negative charge
From the charges on the ions, you can determine the formula of an ionic compound

14. Determining Formula of Ions
Example: BORON OXIDE
Write the symbols of both elements (cation 1st, anion 2nd)
Write the valence of each as a superscript
Drop the positive and negative signs
Crisscross the superscripts so they become subscripts
Reduce when possible (not possible here)
B O
B3+ O2-
B3 O2
B3 O2
B2O3

15. Transition Metals
Transition metals have a varying number of valence electrons
For example:
Copper can form Cu+ and Cu2+ ions
These are written as Copper (I) and Copper (II)
Iron can form Fe2+ or Fe3+ ions
These are called Iron (II) and Iron (III)
The Roman numeral represents the CHARGE of the ion

16. Transition Metals
You can use the ionic formula to determine the charge of the transition metal ion
If you are given CuCl2, what would be the charge on the copper atom?
Copper would have a 2+ charge because of the “2” subscript after chlorine

17. Polyatomic Ions Ions that have many atoms in one ion
You cannot break these ions apart—they always appear together
Ammonium
NH4+
Chlorate
ClO3-
Peroxide
O22-
Acetate
CH3CO2-
Perchlorate
ClO4-
Chromate
CrO42-
Nitrate
NO3-
Permanganate
MnO4-
Dichromate
Cr2O72-
Nitrite
NO2-
Carbonate
CO32-
Silicate
SiO32-
Hydroxide
OH-
Sulfate
SO42-
Phosphate
PO43-
Hypochlorite
ClO-
Sulfite
SO32-
Arsenate
AsO43-
Chlorite
ClO2-
Thiosulfate
S2O32-
Arsenite
AsO33-
Cyanide
CN-
Thiocyanate
SCN-
Borate
BO33-
Bicarbonate
HCO3-
Bisulfate
HSO4-
Bisulfite
HSO3-

18. Polyatomic Ions DO NOT MEMORIZE THE LIST!
Remember that polyatomic ions cannot be broken apart
Example: If you have a sodium ion (Na+) and a hydroxide ion (OH-) they will combine to form NaOH (read: sodium hydroxide)
If you see a name that ends in –ate, this is an indicator that a polyatomic ion is present

19. Covalent Bonds Atoms joined by covalent bonds share electrons
Covalent bonds are usually formed between two nonmetals

20. Naming Covalent Compounds
For covalent compounds of two elements, numerical prefixes are used to tell how many atoms of each element are in the molecule
If there is only one atom of the first element, the name does not get a prefix
The element that is farther to the right in the periodic table is named second, and the ending is changed to -ide
# of Atoms
Prefix 1
mono- 2
di- 3
tri- 4
tetra- 5
penta- 6
hexa- 7
hepta- 8
octa- 9
nona- 10
deca-

21. Covalent Compounds
In a Lewis dot structure, a straight line represents electrons being shared between the atoms
A single straight line = two electrons (each atom contributing one)
Electrons being shared are called the bonding pairs
Electrons that are not a part of the bond are called nonbonding pairs

22. Covalent Compounds
Covalent compounds can also have more than one bond between them
Double bonds occur when atoms share 4 electrons— shown by two straight lines
Triple bonds occur when atoms share 6 electrons— shown by three straight lines

23. Lewis Dot Structures Sum the valence electrons from all atoms
Every atom wants to have 8 electrons around it
8 electrons means it has a full octet, or outermost orbital, which makes the atom stable
Exception is hydrogen—will only have two electrons
To draw a dot structure for a compound –
Sum the valence electrons from all atoms
Use a pair of electrons to form a bond between each pair of bound atoms
Arrange the remaining electrons to satisfy the octet rule (or duet rule)

24. Write the Lewis dot structure for:
H2O
C2H6
H2O2

25. VSEPR Theory Stands for valence-shell electron pair repulsion
In 3-D, all atoms and electrons will try to stay as far away from each other as they possibly can
Electrons and the electron shells of other atoms are all negative, so similar charges repel one another
Remember: electrons that are shared between atoms are called bonding electrons; nonbonding electrons are also called lone pairs

26. Molecular Shape
VSEPR theory can predict the geometry of individual molecules using the number of electron pairs surrounding the central atom

27. PhET Colorado Simulation
Molecular Shape
1. Draw the Lewis dot structure
2. Determine the number of bonded pairs on the CENTRAL atom
3. Determine # of lone pairs on the CENTRAL atom
4. ADD these (STERIC #)
5. Use steric #, # of lone pairs, and the table to determine shape
PhET Colorado Simulation

28. Lone Pairs Steric #2 Steric #3 Steric #4 Steric #5 Steric #6 Linear
Linear
Trigonal Planar
Tetrahedral
Trigonal Bipyramidal
Octahedral 1
(Linear)
(Bent)
Trigonal Pyramidal
Seesaw
Square Pyramidal 2
n/a
Bent
T-Shaped
Square Planar 3

29. Draw the Lewis dot structure and determine its molecular geometry:
CCl4

30. Draw the Lewis dot structure and determine its molecular geometry:
NI3

31. Draw the Lewis dot structure and determine its molecular geometry:
SO2

32. Draw the Lewis dot structure and determine its molecular geometry:
H2O