Homework problems: 32, 33, 35, 36, 38, 41, 44, 48, 51, 53, 56, 58, 59, 64, 67, 69, 72,
73, 75, 78, 80, 82, 87, 90
Outline
I. Octet Rule
A. Rule of 8
1.
for Representative elements (not transition or inner transition metals)
2.
2s, 6p electrons
3.
noble gas electron configuration
4.
"octet" for hydrogen is 2 electrons (1s, but no 1p subshells)
B. Valence electrons
1. outer shell electrons
2. usually add electrons if only need a few to get full octet a. non-metals b. high electronegativity, high ionization energy c. easier to gain a few than lose a lot d. more electrons than protons > negative charge > anion
3. empty the outer shell to get a filled octet a. the next inner shell is full b. metals c. low electronegativity, low first ionization energy d. easier to lose a few than gain a lot e. more protons than electrons > positive charge > cation
C. Transition Elements
1. d shell area of periodic table > octet rule doesn't apply
2. transition metals can have more than one charge a. Cu, Fe, Co
3. fill order not the same as empty order a. s's empty before d's b. lots of +2 charge cations c. half filled d's also stable (Fe +3)
II. Nomenclature of Ions
A. Representative cations
1. name the same, add the word ion a. ex. sodium atom, Na vs sodium ion Na +1
B. Transition metal ions
1. usually more than one positive charge
2. -ic, -ous system a. only used when only two possible charges b. -ic for larger charge, -ous for smaller c.
Cu, Fe, Hg i. must remember iron can be +2 or +3, Cu can be +1 or +2
1

d. disadvantages
3. roman numeral system
C. Anions a. roman numeral is the charge on the ion
1. keep first part of element name, add -ide a. nitrogen > nitride, phosphorus > phosphide
D. Polyatomic ions
1. see table p3-43
2. difference between -ate and -ite
3. only hydroxide and cyanide has -ide ending, all others with -ide must be elemental anions
4. memorize ammonium, hydroxide, sulfate, nitrate, phosphate a. NH
4
+ , OH , SO
4
2, NO
3
, PO
4
3how to tell the charge on the metal ion (for roman numeral or –ic, -ous purposes): use an each total box (more below) and the charge on the non metal ion. ex. Cu
3
(PO
4
)
2
III. Ionic Bonds
A. How it happens
1. electron transfer between high electronegative non-metal and low ionization energy metal
2. electrostatic attraction between resulting anion and cation
B. Electronegativity
1. measure of how badly atom wants electrons
2. F most electronegative
3. size influences electronegativity - smaller size = greater electroneg.
4. N/A for noble gasses
C. Formulas for ionic compounds
1. net charge is zero
2.
use each/total box to determine formula (least common multiple method) ions Na Cl Mg Cl each
Fe Cl Fe S Cu PO
4
Cr OH total
Formula NaCl MgCl
2
FeCl
3
Fe
2
S
3
Cu
3
PO
4
Cr(OH)
6
2

D. Nomenclature of binary ionic compounds
1. for representative metal cations a. Metal + non-metal stem + ide
2. for transition metal cations a. metal stem + ic (or + ous) + non-metal stem + ide b. roman numerals: metal +(roman numeral) + non-metal stem + ide
E. Nomenclature with polyatomic
. a. metal stem + ic (or + ous) + name polyatomic anion b. roman numerals: metal +(roman numeral) + name polyatomic anion c. if ammonium is cation, no roman numeral or –ic, -ous needed
F. Ionic Crystals
1.
not a single discreet NaCl pair
2.
formula is least common multiple
IV. Covalent Bonds
A. How made
1. electrons shared
B. Polar vs Non-Polar
1. not always shared equally
2. dipole
3. two ways to indicate polar bonds a. arrows towards more electroneg atom b. δ+, δ-
O
H H
C. Lewis Structures p. 107 or...
1.
count up valence electrons
2.
add or subtract for charge
3.
figure out how many electrons needed to fill octets
4.
use steps 1 and 2 to determine how many electrons that you have
5.
subtract electrons you have (4) from electrons needed (3), result is electrons shared
6.
each shared electron pair equals one bond
7.
arrange all outer atoms around the central atom (usually least electronegative)
8.
put in shared pairs (bonds) – make sure everything has single bond before going to double bonds or triple bonds
9.
do not violate octet rules (no more than 4 bonds or lone pairs to any atom unless 3rd row or lower a. why? empty d shells to use
10. double check your electron count
D. Nomenclature
1. mono-, di-, tri-
2. first atom does not need mono- if only one of them
3. second atom always gets prefix, even mono-
3

4. second atom changes to –ide ending
N
2
O:
E. Resonance
CO
2
: CO:
1. movement of lone pair, double and triple bonds
O
O
CCl
H C O
H C O
F. VSEPR theory (Valence Shell Electron Pair Repulsion Theory)
1. shapes a.
2 atoms/lone pairs = linear
O C O
O
4
: b.
3 atoms/lone pairs = trigonal
H
C
H
Cl
C
O
Cl
H H
Cl c.
4 atoms/lone pairs = tetrahedral
Cl
G. Polar molecules vs polar bonds
1. need at least one polar bond to be polar molecule
2. only one polar bond => polar molecule
3. more than one polar bond, is it symmetrical? a. if yes, polar bonds cancel each other out

no polar molecule b. if no, polar bonds don’t cancel each other out 
polar molecule
Cl
Cl
Cl
Cl
C
Cl
Cl
Cl
C
H
4

Chapter Objectives
Knowledge
Recall the structure and names for 6 polyatomic ions (H
3
O
+
, OH
-
, NH
4
+
, NO
3
-
,
PO
4
3, SO
4
2)
Know that iron exists as a +3 and a +2 cation and that Cu can be +1 or +2
Remember the following terms: ion, cation, anion, intermolecular, covalent bond, ionic bond
Remember that non polar covalent bonds form between C and H (electronegativity values are virtually the same)
Recall that when naming or writing a formula, the cation comes first, then the anion
Comprehension
Recognize that ionic bonds formed by electron transfer, then attraction between opposite charges (usually a metal and a nonmetal, but sometimes involve polyatomic ions)
Understand that covalent bonds are electron sharing
Understand that electrons are not always shared equally (polar vs. non-polar)
Understand how the octet rule applies to some elements and why it doesn’t work for transition and inner-transition metals (fill octets to generate ion charges)
Predict electronegativity trends across the periodic table
Application
Use the octet rule to predict charges on representative atoms in the periodic table (not transition and inner-transition metals)
Interpret whether a bond is ionic or covalent from the position on the Periodic
Table (refer to metals and nonmetals).
Apply nomenclature rules to compounds to name them (mono, di, etc. for non-metals/ both the roman numeral and the –ic, -ous ending for metal compounds)
Given the name for a compound, determine the formula
Using Lewis structures and VSEPR theory, predict the shape of electron density around a central atom
Indicate polar covalent bonds by either of the methods talked about in class (arrow to more electronegative or

+,

-)
Analysis
Draw Lewis Structures for simple molecules
5