Replacement & Double Replacement Notes
Add hydrolysis of salt RXN: Na2S+H2O  NaOH + H2S?
AP Test: Reaction question directions:
For each of the following three reactions, write a balanced equation for the reaction in part (i) and answer the question about
the reaction in part (ii). In part (i), coefficients should be in terms of lowest whole numbers. Assume that solutions are
aqueous unless otherwise indicated. Represent substances in solutions as ions if the substances are extensively ionized. Omit
formulas for any ions or molecules that are unchanged by the reaction. You may use the empty space at the bottom of the
next page for scratch work, but only equations that are written in the answer boxes provide will be scored.
Net Ionic Reactions:
Consider the following reaction:
potassium chromate + barium nitrate  barium chromate + potassium nitrate
K2CrO4
+
Ba(NO3)2 
BaCrO4
+
2 KNO3
Most ionic compounds are solid (think table salt). In order for these reactions to occur the reactants must be dissolved in
water.
In this reaction, the barium chromate is a precipitate (insoluble). To indicate what is happening in this reaction we sometimes
include the physical state of the reactants and products. This equation is called the molecular equation. It shows the
complete formulas of all reactants and products.
(aq) = aqueous or dissolved in water
(g) = gas
(l) = liquid
(s) = solid
K2CrO4 (aq)
+ Ba(NO3)2 (aq) 
BaCrO4 (s) +
2 KNO3 (aq)
When ionic compounds dissolve in water they break apart into their ions. If the compound is a precipitate that means it does
not dissolve and does not break apart into ions. Molecules do not break apart into ions either (well actually some do – but
we’re not going to worry about this in this class).
To give an even better indication of exactly what is happening, we separate the ions in the soluble compounds to give us what
is called the complete ionic equation as written below:
2 K+1 (aq) + CrO4-2 (aq) + Ba+2 (aq) + 2 NO3-1 (aq)  BaCrO4 (s) + 2 K+1 (aq) + 2 NO3-1 (aq)
Notice that there are K+1 and NO3-1 ions on both sides of the reaction. Since these ions are on
both sides of the reaction, they are not actually part of the reaction and are called spectator ions.
The ions that actually participate in the chemical reaction are only chromate and barium. This equation is called the net ionic
equation and includes only the ions that participate in the reaction (making the precipitate or molecule.)
Net Ionic Reaction:
Example 2:
potassium hydroxide + sulfuric acid  water + potassium sulfate
A little lie…..Not all acids formed in ionic reactions indicate a reaction …. only WEAK acids!
When we write ionic or net ionic reactions, all strong electrolytes are written as ions.
Reactions that produce one of the 7 strong acids may not occur.
Ex: NaCl (aq) + HNO3(aq) NaNO3 (aq) + HCl(aq)
Na+ + Cl- + H+ + NO3-  Na+ + NO3- + H+ + ClNo new substances are formed = no reaction
Oxidation-Reduction (Redox) Reactions
A process where electrons are transferred from one substance to another.
One substance is oxidized while another substance is reduced
Oxidation
•
___________________________________
•
an __________________ in oxidation number
Reduction
- ____________________________
- a _________________ in oxidation number
How can you tell when a redox reaction is taking place?
1) Assign oxidation numbers to atoms in substances
2) Compare oxidation numbers before and after reaction to determine if atom has lost or gained electrons
Rules For Assigning Oxidation Numbers - A bookkeeping system
1.
The oxidation number of a ________________________________.
Examples: Na, H2, Br2, S8, Ne
Ox. #
2.
The oxidation number of a monatomic ion is ______________________.
Examples: Na+1, Ca+2, Al+3, Cl-1, O-2
Ox #
3. The sum of the oxidation numbers of all atoms in a neutral compound is _________.
4. The sum of the oxidation numbers of all atoms in an ion is equal to the ____________________________.
5. In compounds, fluorine is always assigned an oxidation number of ___________.
6. Hydrogen’s oxidation number will be
+1 when bonded to a nonmetal (HCl)
-1 when bonded to a metal (NaH)
Examples:
NaH
CaH2
HCl
Na—H
H—Ca—H
H—Cl
H2S
H—S—H
7.
Oxygen usually has an oxidation number of -2. Exceptions: in peroxides, oxygen will be -1 and when combined with F,
it will be +2 and in O2 it will be 0
Examples:
H2O
CaO
H2O2
H—O—H
Ca—O
H—O—O—H
O2-2
[O—O]-2
8.
OF2
F—O—F
Halogens usually have an oxidation number of -1. Exception is when chlorine, bromine, iodine are combined with
oxygen
Examples:
NaCl
Na—Cl,
MgI2
I—Mg—I,
OCl2
Cl—O—Cl,
HOBr
H—O—Br
** If none of the above rules help you get started…look for an atom with a known charge, and use that charge as its oxidation
number
CdS:
Cd-S
Examples:
CO2
H2SO4
FeSO4
S2O3-2
AlH3
ClO4-1
Na2Cr2O7
NH4+1
Fe2(SO4)3
Once oxidation numbers have been assigned, compare them before and after the reaction
2 Fe2O3 (s) + 3 C (s)  4 Fe (s) + 3 CO2 (g)
Fe is reduced, going from +3 to 0 and C is oxidized, going from 0 to +4, the O undergoes no change
Reducing agent
Causes reduction
Loses electrons
Undergoes oxidation
Oxidation number of atom increases
Oxidizing agent
Causes oxidation
Gains electrons
Undergoes reduction
Oxidation number of atom decreases
Assign oxidation numbers, indicate what is oxidized and reduced, indicate what is the oxidizing agent and reducing agent
1) Ca (s) + 2 H+1  Ca+2 (aq) + H2 (g)
Ca is ___________________________
Ca is the ________________________
H+1 is __________________________
H+1 is the _______________________
2) 2 Fe+2 (aq) + Cl2 (aq)  2 Fe+3 (aq) + 2 Cl-1 (aq)
Fe+2 is ___________________________
Fe+2 is the ________________________
Cl2 is ____________________________
Cl2 is the _________________________
In general, metals act as reducing agents (are oxidized) and nonmetals act as oxidizing agents (are reduced).
Replacement Reactions
Are redox reactions!
(So are many of the reactions we studied in chapter 3!)
Example: It is possible for metals to be oxidized in the presence of a salt (in solution):
Fe(s) + Ni(NO3)2(aq)  Fe(NO3)2(aq) + Ni(s)
OR
In this reaction iron has been oxidized to Fe2+ while the Ni2+ has been reduced to Ni.
The Activity Series
We can list metals in order of decreasing ease of oxidation.
This list is the activity series.
The metals at the top of the activity series are called active metals and are easily oxidized.
The metals at the bottom of the activity series are called noble metals and NOT easily oxidized.
A metal in the activity series can only be oxidized by a metal ion below it.
(Higher will replace lower)
Example (Dime Lab): If we place Cu into a solution of Ag+ ions, then Cu2+ ions can be formed because Cu is above Ag in the
activity series:
Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)
or
Cu(s) + 2Ag+(aq)  Cu2+(aq) + 2Ag(s)
Decompositon of acids:
*This is NOT a redox reaction – the oxidation number on the nonmetal must stay constant. Use this knowledge to
predict the formula for the nonmetal oxide.
H2SO3 
Composition Reactions:
These are NOT redox reactions…oxidation numbers stay constant. Use that knowledge to predict the product!
SO2 + H2O 
P4O10 + CaO 