Chapter 7
REALLY Important!!!
7.1 – Ionic Compounds: Ions for s and p block elements:
Group 1
H
Li
Na
K
Rb
Cs
Fr
Group 2
Be
Mg
Ca
Sr
Ba
Ra
Group13
B
Al
Ga
In
Tl
Group 14
C
Si
Ge
Group 15
N
P
Group 16
O
S
Group 17
F
Cl
Br
I
At
________ Valence e- ________ Valence e- ________ Valence e- ________ Valence e- ________ Valence e- ________ Valence e- ________ Valence e-
Dot Diag:
Dot Diag:
Dot Diag:
Dot Diag:
Dot Diag:
Dot Diag:
Dot Diag:
Wants to _________
Wants to _________
Wants to _________
Wants to _________
Wants to _________
Wants to _________
Wants to _________
All ions are ________
All ions are ________
All ions are ________
All ions are ________
All ions are ________
All ions are ________
All ions are ________
Names:
Names:
Names:
Names:
Names:
Names:
Sn
Pb
As
Sb
Bi
Se
Te
Po
Dot Diag:
Dot Diag:
Dot Diag:
Can lose 2
Can lose 3
Can lose 4
Can lose 4
Can lose 5
Can lose 6
Names:
7.1 – Ionic Compounds: Ions – d Block
*Intro Classes can put Roman Numerals on ALL d-block elements
*Honors Class must know which elements NEED Roman Numerals
cobalt (II) = Co+2
cobalt (III) = Co+3
Cu+1
copper (I) =
copper (II) = Cu+2
chromium (II) = Cr+2
chromium (III) = Cr+3
chromium (VI) = Cr+6
iron (II) = Fe+2
iron (III) = Fe+3
Silver (I) = Ag+1
platinum (II) = Pt+2
platinum (IV) = Pt+4
mercury (I) = *Hg2+2 (Academic: Hg+1)
mercury (II) = Hg+2
Zinc (II) = Zn+2
cadmium (II) = Cd+2
manganese (II) = Mn+2
manganese (IV) = Mn+4
nickel (II) = Ni+2
nickel (III) = Ni+3
gold (III) = Au+3
7.1 – Polyatomic Ions
The Tough Stuff!
THE 8 -ATES:
• Carbonate
• Nitrate
• Sulfate
• Chlorate
• Chromate
• Bromate
• Phosphate
• Iodate
CO3-2
NO3-1
SO4-2
ClO3-1
CrO4-2
BrO3-1
PO4-3
IO3-1
Rules with –ates:
1 more oxygen than –ate =
per … ate
1 less oxygen than –ate =
…ite
2 less oxygens than –ate =
hypo … ite
7.1 – Other Polyatomic Ions
Ammonium
Hydronium
Peroxide
Hydroxide
Cyanide
NH4+1
H3O+1
O2-2
OH-1
CN-1
Ions Test Ends Here………
7.1 – Ionic Compounds – Names
and Formulas
Naming Binary Ionic Compounds:
A metal + a nonmetal = IONIC
Name = cation name anion name (w/-ide ending)
-Use Roman Numerals if needed
Examples:
7.1 – Ionic Compounds – Names
and Formulas
Writing Binary Ionic Formulas:
Remember Chapter 6?
New way: Use the swap technique –
Number on charge tells you how many of
the OTHER element
Examples:
More Examples
7.1 – Ionic Compounds – Names
and Formulas
Naming Ionic Compounds w/more than 2
elements:
Name = cation name anion name
(at least one will be a polyatomic ion)
Examples:
7.1 – Ionic Compounds – Names
and Formulas
Writing Ionic Formulas for Ionic Compounds
with More Than 2 Elements:
Use the swap technique
Examples:
Ionic Compound Review
Formula of:
Calcium bromide
Copper (II) sulfide
Nickel (II) phosphate
Strontium hydroxide
Barium nitrite
Name of:
CaO
MgCl2
CrO3
Fe3(PO4)2
Al2(CO2)3
7.1 – Molecular Compounds
2 Nonmetals Bonded
USE PREFIXES!
Mono=
Di=
Tri=
Tetra=
Penta=
Hexa=
Hepta=
Octa=
Nona=
Deca=
7.1 - Molecular Compounds
Name = prefix first element prefix second element-ide
Prefix is the quantity of that element
Mono is not needed in front of the FIRST element only
Examples:
7.1 – Review
Practice Problems
-Ions
-Naming Ionic Compounds
-Writing Ionic Compound Formulas
-Naming Molecular Compounds
-Writing Molecular Compound Formulas
7.2 – Oxidation Numbers
-Show distribution of electrons
-Negative means “stronger element / taking electrons” and positive
means “weaker element / losing electrons”
Rules:
1. Uncombined element has oxidation number = 0
2. Monatomic ion has oxidation number = charge
3. If in a compound:
A. Start with the element on the right. It has oxidation number =
charge if it was a negative ion.
B. If more than 2 atoms, go next to the element on the left, it has
oxidation number = charge if it was a positive ion.
C. Calculate the remaining element (It may not match what you
memorized as the ion charge)
4. Sum of all oxidation numbers in a neutral compound = 0
5. Sum of all oxidation numbers in a polyatomic ion = charge on ion
7.2 – Oxidation Numbers
Practice Problems:
7.3 – Molar Mass
Molar Mass = Sum of Average Atomic
Masses for all elements in a compound.
Unit = g/mol
Examples:
End after first example
Using Molar Mass
Can be used as a conversion factor
O2 = 32 grams / 1 mole OR 1 mole / 32 grams
Examples:
Percent Composition
Percent by mass of an element in a
compound
(Mass element / molar mass) x 100 = % composition
Examples:
EXPERIMENTALLY Determining
Formulas
Empirical Formula = The simplest (most reduced)
formula
Example: The empirical formula of glucose
(C6H12O6) is….
Steps:
1. Find moles of each element
2. Write formula with subscripts
3. Divide all moles by the smallest number of moles
4. If .5s multiply all by 2
Empirical Calculations
Examples:
EXPERIMENTALLY Calculating
Molecular Formulas
This is the NON-Reduced formula (ex – glucose = C6H12O6)
Need to have the Empirical Formula and the
molecule’s Molar Mass
Examples:
-If the empirical formula is BH3, and the molar mass
of the compound is 27.7 g/mol, what is the
Molecular Formula?