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Reaction Predictions
Most Commonly Used
Cations and Anions
Hydrogen H+
 Sodium Na+
 Potassium K+
 Calcium Ca+²
 Magnesium Mg+²
 Iron (Ferrous) Fe+²
 Iron (Ferric) Fe+³
•Hydroxide OHˉ
•Chloride Clˉ
•Sulfide Sˉ²
•Bicarbonate HCOзˉ
•Carbonate COзˉ²
•Sulfate SO4ˉ²
•Phosphate PO4ˉ³
Cations/ Anions, contd.
You can figure out the charge of an ion by
using the periodic table. For Example:
 Alkali metals such as Lithium can easily
lose an electron to become stable (just like
a Noble gas) so taking away an electron
give Lithium a +1 charge.
 On the other hand Halogens can easily
accept an electron to become stable.
Accepting an electron gives halogens a -1
What is the oxidation state of Oxide?
What is the oxidation state of Iodide?
What is the oxidation state of a Calcium
What is the oxidation state of a Lithium
Net Ionic Equation
To create a net ionic equation, you break
apart all ionic molecules in a balanced
molecular equation into their ions if they
are soluble.
 If there are spectator ions, ions that
appear on both sides of the equation, they
cancel each other.
Net Ionic Example
Silver nitrate is mixed with potassium
2AgNO3 + K2CrO4 → Ag2CrO4 + 2KNO3
Molecular Equation
 2Ag+ + 2NO3ˉ + 2K+ + CrO4-2 → Ag2CrO4 +
2K+ + 2NO3-2
Complete ionic equation
 2Ag+ + CrO4-2 → Ag2CrO4
Net Ionic Equation
Solubility Rules
all nitrates are soluble
CH3COO or C2H3O2
all acetates are soluble except AgCH3COO
all chlorates are soluble
Cl all chlorides are soluble except AgCl, Hg2Cl2,
Br all bromides are soluble except AgBr, PbBr2,
Hg2Br2, and HgBr2
I all iodides are soluble except AgI, Hg2I2, HgI,
and PbI2
Solubility Rules, contd.
SO4¯² all sulfates are soluble except
BaSO4, PbSO4, Hg2SO4, CaSO4, AgSO4
and SrSO4
 Alkali metal, cations, and NH4 – all are
 H+
all common inorganic acids and low
molecular mass organic acids are soluble
(In)Soubility Rules, contd.
CO3 ²
all carbonates are insoluble except those of
alkali metals and NH4
CrO4 ² all chromates are insoluble except those of
alkali metals, NH4, CaCrO4, and SrCO4
all hydroxides are insoluble except those of the
alkali metals, NH4, Ba(OH)2, Sr(OH)2, and Ca(OH)2
PO4 ³
all phosphates are insoluble except those of
alkali metals and NH4
SO3 ²
all sulfites are insoluble except those of alkali
metals and NH4
S ²
all sulfides are insoluble except those of alkali
metals and NH4
Synthesis occurs when two or more reactants combine
to form a single product. There are several common
types of synthesis reaction.
You know it happens when you have:
-A metal combines with a nonmetal to form a bianary
-A piece of lithium metal is dropped into a container
of nitrogen gas.
6Li+ N2  2Li3N
-Metal oxide and water forms a base (metallic
-Solid sodium oxide is added to water.
Na2O + H2O 2NaOH
Synthesis, contd.
Nonmetal oxide and water forms acids.
Nonmetal retains its oxidation number.
-Carbon dioxide is burned in water.
CO2 + H2O  H2CO3
 Metallic oxides and nonmetallic oxides
form salts.
-Solid sodium oxide is added to carbon
Na2O + CO2  Na2CO2
Occurs when a single reactant is broken
down into two or more products.
 The reactions react to form basic
compounds or elements.
 When a compound is heated or
electrolyzed, it means that it is broken up
into its ions.
 AB A+B
Examples of Decomposition
A sample of magnesium carbonate is
MgCO3  MgO + CO2
 Molten sodium chloride is electrolyzed.
2NaCl  2Na + Cl2
 A sample of ammonium carbonate is
(NH4)2CO3  2NH3 + H2O + CO2
Single Replacement
Reactions that involve an element
replacing one part of a compound. The
products include the displace element and
a new compound. An element can only
replace another element that is less active
than itself. (Look a activity series/ AP
 A +BX B+AX
Single Replacement Rules
Active metals replace less active metals
from the less active metals’ compounds
in aqueous solutions
ex. 3Mg+ 2FeCl3—> 2Fe + 3MgCl2
Active metals replace hydrogen in water
ex. 2Na + 2H2O—> H2 + 2NaOH
Active metals replace hydrogen in acids
ex. 2Li + 2HCl —> H2 + 2LiCl
Single Replacement Rules,
4. Active nonmetals replace less active
nonmetals from their compounds in
aqueous solutions
ex. Cl2 + 2KI —> I2 + 2KCl
5. If a less reactive element is combined
with a more reactive element in compound
form, there will be no reaction
ex. Cl2 + KF —> no reaction*
* On the AP test reactions will ALWAYS
have products; it will never be “no
Activity Series (Single Replacement)
Li, Ca, Na, Mg, Al, Zn, Fe, Pb, [H2], Cu, Ag, Pt
F2, Cl2, Br2, I2,
More active
Less Active
Double Replacement
Two compounds react to form two new compounds. No
changes in oxidation numbers occur.
Each cation pairs up with the anion in the other
The “driving force” in these reactions is the removal of at
least one pair of ions from solution.
This removal of ions happens with the formation of a
precipitate, gas, or molecular species.
When a double replacement reaction doesn’t go to
completion, it is a reversible reaction (no ions have been
How do you know a double
replacement reaction occurs?
The reactants will contain a(n):
-insoluble precipitate
-molecular species
*Remember– on the AP test the reaction will
always occur
Common Gases Released (Dbl. Repl.)
H2S Any sulfide plus any acid forms
H2S and a salt.
 CO2 Any carbonate plus any acid form
CO3, water, and a salt.
 SO2 Any sulfite plus any acid form SO2,
water, and a salt.
 NH3 Any ammonium plus a soluble
hydroxide form NH3, water, and a
Acid/ Base Reactions (Dbl.
An acid and a base will react and form
water and a salt.
 Hydrochloric acid is added to sodium
HCl + NaOH  NaCl + H2O
Hydrolysis (Dbl. Repl.)
It is the reverse of neutralization and results when a salt plus a water molecule yields
an acid plus a base.
Salt + water  acid + base
Key things to know about hydrolysis reactions:
 Salts of a strong acid plus a weak base will hydrolyze into an acidic solution.
NH4+ +Cl- +H2O → H+ +Cl- + (NH)4OH
Salts of a weak acid and a strong base will always hydrolyze to give a basic
+F- +H2O → K+ +OH- +HF
Salts of a strong acid and a strong base will never undergo hydrolysis and
therefore make a neutral solution.
Na+ +Cl- +H2O → Na+ +OH- +H+ Cl-
Salts of a weak acid plus salts of a weak base may hydrolyze as an acid, base,
or a neutral solution; the final result depends on the Ka’s and Kb’s of the acids
and bases formed during the hydrolysis process.
Disclaimer!! The spectator ions were not removed 
Examples of Dbl. Replacement
Solutions of potassium bromide and silver
nitrate are mixed.
KBr + AgNO3  AgBr + KNO3
 A solution of sodium sulfate is added to a
solution of hydrochloric acid.
Na2SO3 + 2HCl  2NaCl + H2SO3
Hydrolysis Sample Problems
Try these:
An aqueous solution of manganese (II) sulfate
undergoes hydrolysis.
 Ammonium fluoride and water are mixed
Hydrolysis answers
MnSO4 + 2H2O → H2SO4 + Mn(OH)2
 NH4F + H2O → HF + NH4OH
Combustion (Organic Reacs.)
An organic compound reacts with O2 to
form water and carbon dioxide.
 If something is burned there is a
combustion reaction.
 Methanol is burned in oxygen gas.
2CH3OH + 3O2  4H2O + 2CO2
Addition (Organic Reacs.)
A halogen or hydrogen is added to an
alkene or alkyne, breaking apart the
double or triple bonds and forming single
 Fluorine is added to ethene
F2 + CH2=CH2  CH2F-CH2F
Substitution (Organic Reacs.)
An atom attached to a carbon is removed
and something else takes its place.
 Bromine is added to methane
Br2 + CH4  CH3Br + HBr
Oxidizing Agents (Redox
Common Oxidizing Agents
MnO4¯ in acidic solution
MnO2 in acidic solution
MnO4¯ in neutral or basic solution
Cr2O7ˉ² in acidic solution
HNO3, concentrated
HNO3, dilute
H2SO4, hot, concentrated
Metallic ions (higher oxidation #)
Free halogens
H 2 O2
Products Formed
Metallous ions (lower oxidation #)
Halide ions
Reduction Agents (Redo
Common Reducing
 Halide ions
 Free metals
 Sulfite ions or SO2
 Nitrite ions
 Free halogens, dilute
basic solution
 Free halogens,
concentrated basic
 Metallous ions (lower
oxidation #)
Products Formed
Free halogen
Metal ions
Sulfate ions
Nitrate ions
Hypohalite ions
Halite ions
Metallic ions
(higher oxidation #)
Electrolysis (Redox Reacs.)
An electrolysis reaction is a reaction in which a nonspontaneous redox reaction is brought about by the
passage of current under sufficient external electrical
potential. The devices in which electrolysis reactions
occur are called electrolytic cells.
In theory, E° values (Standard Reduction Potentials) can
be used to predict which element will plate out at a
particular electrode when various solutions are
(B&L text)
Rules for Predicting Cathode
Reactions (Reduction)
When a direct electric current is passed
through a water solution of an electrolyte,
two possible reduction processes may
occur at the cathode.
 The cation may be reduced to the
corresponding metal.
Mn+ + ne-  M(s) (reaction 1)
n = (charge of cation)
 Water molecule may be reduced to
elementary hydrogen
2H2O + 2eˉ  H2 + 2OHˉ (reaction 2)
Rules for Predicting Cathode
Reactions, contd.
For salts containing transition metal cations,
which are relatively easy to reduced compared
to water, reaction #1 will occur at the cathode
(and the transition metal will plate out).
Mn+ + ne-  M(s)
If the cation is representative metal, the water
molecules will be easier to reduce compared to
the cation, and reaction #2 will occur at the
cathode, producing hydrogen gas and hydrogen
2H2O + 2eˉ  H2 + 2OHˉ
Rules for Predicting Anode
Reaction (oxidation)
The oxidation process that occurs at the anode of an
electrolytic cell operating in aqueous solution may be
one of two oxidation processes.
The anion may be oxidized to the corresponding
- 2Xˉ  X2 + 2eˉ (reaction 1)
Water molecules may be oxidized to elementary oxygen.
- HOH  ½ O2 + 2H+ + 2eˉ (reaction 2)
Rules for Predicting Anode
Reactions, contd.
For salts containing iodide, bromide, or chloride
ions, it is usually easier to oxidize these
nonmetals rather than water. It will be found that
the nonmetal is formed at the anode.
When the anion present is any other ion that is
more difficult to oxidize than water, Reaction #2
will occur at the anode producing elementary
oxygen and aqueous hydrogen ions.
Example Electrolysis Reactions
Copper (II) chloride in water
Cu+2 + 2Clˉ  Cu + Cl2
2. Copper (II) sulfate in water
Cu+2 + HOH  Cu + ½ O2 + 2H+
3. Sodium chloride in water
2Clˉ + 2HOH  H2 + Cl2 + 2OHˉ
4. Sodium sulfate in water
2HOH  2H2 + O2
Metals w/ Multiple Oxidation Levels
(Redox Reacs.)
These metals can change their oxidation state in a redox reaction
 Antimony (III) or (V)
 Bismuth (III) or (IV)
 Cerium (III) or (IV)
 Chromium (II) or (III)
 Cobalt (II) or (III)
 Copper (I) or (II)
 Gallium (I) or (II) or (III)
 Germanium (II) or (IV)
 Gold (I) or (III)
 Iron (II) or (III)
 Lead (II) or (IV)
 Mercury (I) or (II)
 Nickel (II) or (III)
 Thallium (I) or (III)
 Thorium (II) or (IV)
 Tin (II) or (IV)
Tin (II) sulfate is added to iron (III) sulfate
SnSO4 + Fe2(SO4)3  Sn(SO4)2 + 2FeSO4
Complex Ion Reactions
Nomenclature is on pages 23-27 of The Ultimate
Chemical Equations Handbook
There are a lot of very complicated types of these
reactions, but, for all intensive purposes and for the AP
test, you only need to be familiar with those reactions
pertaining to ammonia and water.
In a complex ion reaction, ligands will attach to a
transition metal ion.
There will usually be twice as many ligands as the
metals oxidation number
Complex Ion Reactions, contd.
These reactions usually occur in a
concentrated solution of the ligand.
 Copper chloride (II) is added to a
concentrated solution of ammonia
Cu2+ +NH3  [Cu(NH3)4]2+
Common Reaction Terms
Electrolysis: Electricity is run through a compound,
resulting in a change of oxidation states.
Hydrolysis: The reaction of a salt with water to form
molecular species. Salts of a strong acid + a weak base
will always hydrolyze to give an acidic solution.
Neutralization: Acid and base react to form a salt and
Catalyst: A molecule that speeds that speeds a reaction
but that does not appear in the reaction.
Oxidation number: the charge that it would have if all the
ligands (atoms that donate electrons) were removed
along with the electron pairs that were shared with the
central atom
Common Reaction Terms,
Precipitate: an insoluble substance formed
by the reaction of two aqueous
 Anode: the electrode where oxidation
an ox
 Cathode: the electrode where reduction
red cat
By: Will Lambert, Adam Robinson,
Michelle Klassen, and Tori Waldron
 (APChem ‘06-’07)