1. Predicting Products of Reactions
Double and Single Replacement

2. Double replacement Two main types a. formation of a precipitate
b. acid/base reaction (produces a salt plus water)
(c. or both!)
use solubility rules and electrolytic properties
only break apart into ions if soluble, in solution, and a strong electrolyte

3. Must have a driving force that removes ions from the solution
a precipitate forms and/or
a gas forms and/or
a weak/nonelectrolyte (water usually) forms

4. Formation of a precipitate
A precipitate is an insoluble compound (a solid)
Use solubility rules to determine if one or both of the products is a precipitate
Example
3AgNO3 (aq) + K3PO4 (aq)  Ag3PO4 (s) + 3KNO3 (aq)
insoluble soluble
precipitate
Net ionic eqn: 3Ag+(aq) + PO4-3(aq)  Ag3PO4(s)

5. Formation of a gas
Must know common gases
If you get H2CO3 when you do the double replacement it really gives H2O + CO2 (g)
If you get H2SO3 when you do the double replacement it really gives H2O + SO2 (g)
H2S is a gas

6. Na2CO3(aq)+ H2SO4(aq)Na2SO4(aq)+ H2CO3(aq) REALLY IS:
For example
Na2CO3(aq)+ H2SO4(aq)Na2SO4(aq)+ H2CO3(aq)
REALLY IS:
Na2CO3(aq)+H2SO4(aq)Na2SO4(aq)+H2O(l)+CO2(g)
Net ionic eqn: CO3-2(aq) + 2H+(aq)  H2O(l) + CO2(g)

7. Formation of a nonelectrolyte
Water is a common nonelectrolyte and is a result of an acid-base reaction
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Net ionic eqn: H+ (aq)+ OH- (aq) H2O(l)

8. Redox Reactions
Some synthesis, some decomposition, single replacement, and other reactions are oxidation-reductions
OXIDATION-REDUCTION REACTIONS involve electron transfer

9. Terms to Know:
OIL RIG – oxidation is loss, reduction is gain (of electrons)
Oxidation – the loss of electrons, increase in charge (becomes more positive)
Reduction – the gain of electrons, reduction of charge (becomes more negative)
Oxidation number – the assigned charge on an atom
Oxidizing agent (OA) – the species that is reduced and thus causes oxidation
Reducing agent (RA) – the species that is oxidized and thus causes reduction
*note that some in some reactions, the same species can be oxidized and reduced; these are called disproportionation reactions

10. Rules for Assigning Oxidation States
1. The oxidation state of an atom in an element is ZERO including allotropes [i.e. N2, P4, S8].
2. The oxidation state of a monatomic ion is the same as its charge.
3. In its compounds, fluorine is always assigned an oxidation state of -1.
4. Oxygen is usually assigned an oxidation state of -2 in its covalent compounds, such as CO, CO2, SO2, and SO3.
Exceptions to this rule include peroxides (compounds containing the O22- group), where each oxygen is assigned an oxidation state of -1, as in hydrogen peroxide (H2O2),
and OF2 in which oxygen is assigned a +2 oxidation state.
5. In its covalent compounds with nonmetals and acids, hydrogen is assigned an oxidation state of +1.
Metal hydrides are an exception; H is at the end of the chemical formula since it has an oxidation state of -1. (such as NaH)
6. The sum of the oxidation states must be zero for an electrically neutral compound. For a polyatomic ion, the sum of the oxidation states must equal the charge of the ion.

11. There can be non-integer oxidation states like in Fe3O4
There can be non-integer oxidation states like in Fe3O4. There’s a -8 for the 4 oxygens divided across 3 iron ions, therefore Fe’s charge is Fe8/3+

12. Single Replacement Redox
How can we predict if a single replacement reaction will occur?
Reading the reduction potential chart
 elements that have the most positive reduction potentials are easily reduced (in general, non-metals)
 elements that have the least positive reduction potentials are easily oxidized (in general, metals)
Can also be used as an activity series. Metals having less positive reduction potentials are more active and will replace metals with more positive potentials. 
The MORE POSITIVE reduction potential gets to indeed be reduced IF you are trying to set up a cell that can act as a battery.

14. For example,
Will the following reaction occur?
Mg(s) + 2KCl(aq)  2K(s) + MgCl2(aq)
Step1: Determine which element is reduced by assigning oxidation numbers before and after
So K was reduced b/c it went from +1 to 0

15. Step 4: Does this reaction occur?
Step 2: Look up the reduction potentials of the elements whose oxidation #’s changed
K= -2.92
Mg=-2.37
Step 3: Determine if the element with the most +reduction potential was reduced
NO! K was reduced in the reaction as written, but Mg has the most positive reduction potential!
K cannot be reduced by Mg, but Mg could be reduced by K
Step 4: Does this reaction occur?
NO! (but note that means that the reverse reaction would occur!)