1. Chapter 5
Naming Ionic Compounds
A. Type I compounds
1. Metal in compounds forms only one type of ion
2. Most main group metals form type I compounds
Naming type I binary ionic compounds
1. Name of cation (metal) + (base name of anion + ide)
2. Example, NaCl is sodium chloride
B. Type II compounds
1. Metal forms more than one type of ion
2. Transition metals usually, but not exclusively, form type II compounds
Naming type II binary ionic compounds
1. Name of cation + (charge of cation) + (base name of anion + ide)
2. Charge of cation given in roman numerals
3. FeCl3 is iron (III) chloride
C. Naming Molecular Compounds BINARY(Greek prefixes!!)
A. Formed between 2 or more nonmetals
B. (Prefix)name of first element + (prefix) base name of second element + ide
C. 1 = mono; 2 = di; 3 = tri; 4 = tetra…
D. Example, N2O is dinitrogen monoxide

2. 5.5 . Naming ionic compounds containing a polyatomic ion
1. Use the same procedure as ionic compounds
2. Use name of polyatomic ion, not constituent atoms
3. Example, NaNO3 is sodium nitrate
5.6 Naming Acids
A. Molecular compounds that dissolve in water to form H+ ions
B. Binary acids
1. Hydrogen and nonmetal
2. (Hydro + base name of nonmetal + ic) + acid
3. Example, HCl is Hydrochloric acid
C. Oxyacids
1. Hydrogen and polyatomic oxyanion
2. Oxyanions ending with –ate
a. (Base name of oxyanion + ic) + acid
b. Example, HNO3 is nitric acid
3. Oxyanions ending with –ite
a. (Base name of oxyanion + ous) + acid
b. Example, HNO2 is nitrous acid
5.7 Formulas and Nomenclature Summary
A. Ionic compounds
B. Molecular compounds
C. Acids

3. Chemical Equations 6 6.1-6.2 The Chemical Equation
The Chemical Equation
A. Reactants go to products
B. Cannot create or destroy atoms in a chemical equation
C. Reaction must be balanced
Evidence of a Chemical Reaction
A. Color change
B. Formation of solid or gas
C. Heat/light absorption or emission
D. transfer of electron
6.3 Balancing Chemical Equations
How to Write Balanced Chemical Equations
A. Write reactant and product species, including physical states (s, l, g, aq.)
B. If an element occurs only once on each side of the reaction, balance that first
C. Balance free reactants last
D. Clear non-integer coefficients
Practice from Lab manual Exercise on Balancing chemical equation

4. 8.4

5. Reaction in Aqueous Solution 7
Identify Evidence of a Chemical Reaction
Aqueous Solutions and Solubility: Compounds Dissolved in Water
A. Aqueous: mixture of substance in water
B. Soluble: dissolves in water Li+, Na+, K+, NH4+,NO3-, C2H3O2-
C. Insoluble: does not dissolve in water (CO3 2–, PO43– except Li+, Na+, K+, NH4+ etc)
D. Solubility rules –Check in the back of lab manual
Important: Practice with Single double displacement :Labs
7.3 Writing equations:
1) First break reactants into ions with the correct charges
2) Then swap anions and cations, and come up with correct formula for products
3) Then identify physical states (s, l, g, aq.),
4) Last step BALANCE!; Ignore total and net ionic equations
7.4 Acid - Base and Gas Evolution Reactions (H2CO3, H2SO3, NH4OH)
Reactions of Acids and Bases-Forms salt and water
Acids produce H+ , bases produce OH- and they react to form water
B. Gas evolution – many different possible products (CO2 , SO2 ,NH3, H2S)
(H2CO3 splits as H2O and CO2 )

6. 7.5 Single displacement reaction- also called redox reaction
Remember the activity series of metal
1A strongest metals, then IIA metals, Fe, …Zn…..
Pb H Cu, Ag and Au least reactive……
A + BC B + AC A is stronger than B
7.6 Other ways to Classify Types of Chemical Reactions
A. Synthesis or combination reactions A + B  AB
B. Decomposition reactions AB  A + B
C. Single Displacement reactions A + BC  AC + B
D. Double-displacement reactions AB + CD  AD + CB

7. 8.1 Stoichiometry—What is it?- Quantities in a chemical reactions
STOICHIOMETRY and MOLE
8.1 Stoichiometry—What is it?- Quantities in a chemical reactions
Mole: Counting Atoms by the Gram
A. Pair =2; dozen = 12; 1 mole = 6.02 x 1023
B. 1 mole of Au atoms = 6.02 x 1023 Au atoms
C. The periodic table gives the mass of 1 mole of that atom in grams
1 mole of S atoms = g S ; 1 mole of C atoms = g C
D. Unit conversions between grams and moles of atoms
Moles = grams/Molar mass
Extend idea to Molecules by the Gram
A. Mass of 1 mole H2O = mass of 2 moles of H + 1 mole O
B. Calculations performed in a similar fashion
C. Learn to calculate molar mass of compounds by adding units
Chemical Formulas as Conversion Factors
A. Molecular formula gives ratios of atoms in molecule
B. Converting moles of atoms to moles of molecules
C. Converting grams of atoms to grams of molecules

8. 8.6 Learn the % composition- EASY!!!
Mass Percent Composition of Compounds
A. Percentage of molecule's total mass that is due to atom X
B. Convert between mass of element and mass of compound
C. Mass Percent Composition from a Chemical Formula
Calculating Empirical Formulas for Compounds
A. Empirical formula is the smallest whole number ratio of atoms
B. Mostly different from molecular formula
C. Many different compounds may have the same empirical formula
Calculating Molecular Formulas for Compounds
A. Empirical formula molar mass
B. Compare molar mass to empirical formula mass
C. Ration= (Molar mass/emp. Mass) = (molecular formula/Emp. Formula)
D. Ratio n is the nearest whole number
4 steps in Empirical formula problem
1) % to grams
2) Grams/Molar mass = Moles
3) Divide all moles by Smallest moles
4) Use appropriate rounding rules to get Empirical formulas
For finding molecular formula Ratio the molecular mass and empirical mass to concert empirical formula to molecular formula

9. EXAM POINTS
Part Multiple Choice –Show calculations for partial/full credit
(60 points) BRING SCANTRON
Part 2, Complete and Balance equations (10 points)
Part 4 Moles, empirical formula and % compositions (30 points)

10. Sample questions—Exam 3
Part I
MULTIPLE CHOICE on all sections, nomenclature
Part II and III
Ba(NO3) K2CO3 
HCl NaOH 
SO O  SO3 -Balance!
Part IV Moles to grams to # of atoms/molecules
empirical formula problems, % compositions