1. Predicting Products of Chemical RXNS
Make sure you have a whiteboard and dry erase marker

2. Recap: Chemical Equations
Coefficient: large number before the element/compound.
Indicates # of the element/compound used in rxn
Subscript: small number next to element symbol
Indicates the number of each element present in the compound/molecule.
States of Matter:
Solid (s), Liquid (l), Gas (g), Aqueous (aq)
New:
 = reversible reaction
= catalyst (in this case, Pt is the catalyst)
or = heat is added

3. Recap: Balancing
List all the element (Single Cap. Letter) on the reactant side of the equation.
Count # of atoms on reactant(left of arrow) & product(right of arrow) side for EACH element.
Use coefficient to balance the equation
Multiply: coefficient x subscript to change # of atoms.

4. Lets review a few of the Problems: __FeCl3 + __Be3(PO4)2  __BeCl2 + __FePO4 __AgNO3 + __LiOH  __AgOH + __LiNO3 __Mg + __Mn2O3  __MgO + __Mn

5. Packet Practice

6. Recap: Rxn Types Synthesis: A + B  AB Decomposition: AB  A + B
2+ reactants  1 product
Decomposition: AB  A + B
1 reactant  2+ products
Single Replacement: A + BX  B + AX
element + compound  element + Compound
Double Replacement: AX + BY  AY + BX
compound + compound  compound + compound
Combustion: AB + O2  H2O + CO2 or Oxide
O2 on REACTANT SIDE to produce CO2 & H2O or an oxide

7. 1. Synthesis (redox)
Two or more elements or compounds that combine to form a single substance either ionic or covalent.
General equation: A + B  AB
EX. Iron and chlorine combine to produce Iron chloride.
2Fe (s) + Cl2 (g)  2FeCl3 (s)
Neutral Neutral Iron +3 oxidized, Chlorine -1 reduced
Iron is the reducing agent, Chlorine is the oxidizing agent

8. Synthesis Special Cases
Non-metal oxide + water  acid
Use oxidation numbers to find the oxidation number of the non-metal.
Use the oxidation number of the non-metal to determine the oxyanion of the acid.
Must balance charges from the ions when you write the formula for the acid.
Example: SO2 + H2O  H2SO3
Non-metal oxide + metal oxide  salt
Use the oxidation number of the non-metal to determine the oxyanion of the salt. The cation is the metal ion.
Must balance charges from the ions when you write the formula for the salt.
Examples:
CO2 + metal oxide  metal carbonate
SO2 + metal oxide  metal sulfite
SO3 + metal oxide  metal sulfate
metal oxide + water  base
The metal is the cation of the base.
The anion is OH-1 (definition of a base).
Must balance charges from the ions when you write the formula for the base.
Example: Al2O3 + H2O  Al(OH)3

9. Practice 1. CaO(s) + H2O(l)  Ca(OH)2(aq) 2. H2(g) + Cl2(g)  2HCl(g)
3. CO2(g) + H2O(l) 
4. K2O(s) + CO2(g) 
Ca(OH)2(aq)
2HCl(g)
H2CO3(aq)
K2CO3(s)

10. 2. Decomposition
A single compound is broken down into two or more substances.
Opposite of synthesis.
General equation: AB A + B
EX. Hydrogen peroxide decomposes into water and oxygen.
2H2O2 (aq)  O2 (g) + 2H2O (l)

11. Decomposition Special cases
acid  Non-metal oxide + water
salt  Non-metal oxide + metal oxide
Uses the oxidation steps from the combination reactions in reverse.
Find the oxidation number of the non-metal in the oxyanion.
Use the oxidation number of the non-metal to determine the formula of the non-metal oxide.
Must balance charges from the ions when you write the formula for the metal oxide. The metal cation is the metal from the salt.
Metal halate  metal halide + oxygen gas
Metal carbonate  metal oxide + carbon dioxide
Metal peroxide  metal oxide + oxygen gas

12. Practice 1. BaCO3  BaO + CO2 2. Ni(ClO3)2  NiCl2 + 3O2 3. Ag2O 
4. H2SO4 
BaO + CO2
NiCl2 + 3O2
Ag + O2
H2O + SO3

13. 3. Single replacement
One element replaces a second element in a compound.
-the more reactive metal will replace the least reactive metal.
-the more reactive nonmetal will replace the least reactive nonmetal.
General equation: A + BX  AX + B (metal)
Q2 + RS  RQ + S2 (nonmetal)
EX. 2AgNO3 (aq) + Cu(s)  2Ag(s) + Cu(NO3)2 (aq)
Silver nitrate reacts with copper to produce silver and copper

14. Activity Series
A chart of metals listed in order of declining relative reactivity. The top metals are more reactive than the metals on the bottom.
The top metals will always replace metals below it in a chemical reaction

15. Practice 1. Au(s) + AgNO3(aq) → No Reaction 2. Fe(s) + Cu(NO3)2(aq) →
3. Ca(s) + H2O(l) →
4. Cl2(g) + KI(s) 
No Reaction
Cu(s) + Fe(NO3)2(aq)
CaO(s) + H2(g)
I2(g) + KCl(s)

16. 4. Double replacement
Involves an exchange of positive ions between two reacting compounds. Reactants are two ionic compounds
One of the following applies:
One product precipitates from solution or
One product is a gas and bubbles out of solution or
One product is a molecular compound such as water.
Use Solubility chart to determine precipitate!
AX + BY  A Y + B X
Ex. AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq)

17. Double Replacement Special cases:
Acid-Base neutralization reactions:
Acid + Base  Salt + water
Acid: Cation = H+
Base: Anion = OH-1
Salt: Anything that can be made from an acid and a base that is not water.
Gas forming reactions:
Acid + metal carbonate  salt + water + carbon dioxide
*Double replacement reaction that produces H2CO3
- H2CO3 breaks down as soon as it forms. Follows the
rules for metal carbonate decomposition.
Acid + metal sulfate  salt + water + sulfur dioxide
-Same type of problem as the acid + metal carbonate.

18. Practice 1. Ca(C2H3O2)2 + Na2CO3  2. NaCl + AgNO3 
3. HNO3 + Na2CO3 
4. HBr + KOH 
CaCO3(s) +2NaC2H3O2(aq)
AgCl(s) + NaNO3(aq)
2NaNO3 + H2CO3
H2O(l) + CO2(g)
H2O(l) + KBr(aq)

19. 5. Combustion
An element or compound reacts with oxygen, often producing energy as heat or light.
CxHyOz + O2  CO2 + H2O
Some synthesis (combination) reaction fall under this category.
Metal reacting to form metal oxides
General form…you get the oxide of every element you burn (except oxygen).
Complete combustion is excess oxygen…gives you CO2 as the oxide of carbon.
Incomplete combustion is limited oxygen…gives you CO as the oxide of carbon.

20. Practice 1. hexane(C6H14) burns in air. 2. ethene (C2H4) burns in air.
3. benzene (C6H6) burns in air.
4. Magnesium ribbon burns in air.
2C6H O2  12CO2 + 14H2O
C2H4 + 3O2  2CO2 + 2H2O
2C6H6 + 15O2  12CO2 + 6H2O
2Mg + O2  2MgO

21. Net Ionic Equations
Net ionic equations show only the species actually involved in the reaction. As you first learn to write net ionic equations, you will write three different equations for each reaction.

22. Do not ionize solids, pure liquids, or gases.
Steps for Net Ionic Equations
Write the complete molecular equation. (This is the type of equation that you are used to writing.)
Write the complete ionic equation. To do this, you must ionize everything that is soluble and ionized in solution. Everything else is left together.
Do not ionize solids, pure liquids, or gases.
Write the net ionic equation. To do this, cancel out all ions that are not participating in the reaction (spectator ions) and rewrite the equation.

23. Ex. Sodium chloride + silver nitrate
NaCl(aq) + AgNO3(aq) 
Na+ + Cl- + Ag+ + NO3-  Na+ + NO3- + AgCl(s)
Cl- + Ag  AgCl
NaNO3(aq) + AgCl(s)
Ex. HCl + Ba(OH)2 
1. 2HCl(aq) + Ba(OH)2(aq) 
2. 2H+ + 2Cl- + Ba2+ + 2OH-  2H2O + Ba2+ + 2Cl-
3. 2H OH-  2H2O
H+ + OH-  H2O (reduce)
2H2O(l) + BaCl2(aq)