Key Points for
REACTION
PREDICTIONS
Review Nomenclature
 Know symbols/names
 Positive monatomic ions’ names
are _____________.
 Negative monatomic ions/names
end in __________.
 ____________ ion is named first.
Polyatomic Ions
 Know “home base” ion (-ate ion) and
charge. Others can be figured out
from there.
 “Home base” ion plus one oxygen
________________
 “Home base” ion minus one oxygen
________________
 “Home base” ion minus two oxygens
______________
You
Gotta
Know…
1
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
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


Bromate
Chlorate
Iodate
Nitrate
Acetate
Azide
Cyanide
Hydroxide
Permanganate
Thiocyanate
2





Carbonate
Chromate
Sulfate
Oxalate
Dichromate
Peroxide
3Arsenate
Borate
Phosphate
Citrate
1+
Ammonium
Hydronium
Other Hints:
 Bi- or hydrogen in the name means
an extra hydrogen is present—Add H
and add +1 to the charge.
 Thio- means a sulfur replaces an
oxygen—Add S & remove O; charge is
the same.
Acid Names
 Derive from negative
monatomic or polyatomic ion
 Monatomic ion plus
hydrogen ________________
 -ate becomes ______
 -ite becomes ______
General Information
 Uncombined elements have an
oxidation number of zero.
 Learn polyatomic ions and their
charges!!
 Think about reasonable
possibilities for given reactants.
Equations
 Word:
Sodium chloride + silver nitrate  sodium nitrate +
silver chloride
 Formula:
NaCl + AgNO3  NaNO3 + AgCl
 Complete Ionic:
Na + + Cl- + Ag+ + NO3-  AgCl + Na+ + NO3 Net Ionic:
Cl- + Ag+  AgCl
Net Ionic Equations
 Dissociate all ions when in
solution. (Is it soluble???)
 Cancel ions that appear on
both sides of an equation
(spectator ions).
Solutions of nickel
(II) nitrate and
cesium hydroxide are
mixed.
Equal volumes of
equimolar sulfuric
acid and sodium
hydroxide are mixed.
REACTION PREDICTIONS
ON THE A.P. EXAM
General Information
 Will always appear on the exam
 Three equations to write & three
questions to answer
 Must write a net ionic equation—NO
SPECTATOR IONS
 Balance the equations—atoms &
charges.
 All reactions will occur. Don’t worry
about activity series.
Reaction Predictions
One of a few general types:
 Double Replacement or Metathesis
 Single Replacement (also redox)
 Redox
 Synthesis/Decomposition
 Complex Ions
 Combustion
Double Replacement or
Metathesis
Starting point: two compounds, often in
solution
 Look for an acid/base reaction
 Look for the formation of an insoluble
compound (precipitate)
 Both can happen simultaneously.
Acid/Base
 A solution of hydrochloric acid is combined
with a solution of sodium hydroxide.
 A solution of hydrofluoric acid is combined
with solid zinc hydroxide.
Precipitates
 A solution of barium chloride is combined
with a solution of potassium sulfate.
 Solutions of cobalt II chloride and lithium
sulfite are combined.
Both
 Sulfuric acid solution in combined
with solid calcium hydroxide.
Single Replacement
Starting point: one compound
(often in solution) and one
uncombined element
 Like will replace like in a reaction.
 Oxidation numbers must change
since an uncombined element’s
oxidation # is always zero.
Single replacement
 Fluorine gas is bubbled through a
solution of potassium chloride.
 A piece of solid aluminum is placed in
a solution of copper II sulfate.
Solid lithium metal
is added to water.
Aqueous solutions of
oxalic acid (H2C2O4)
and excess potassium
hydroxide are mixed.
Synthesis/Decomposition
 Synthesis—two reactants combine
 Synthesis—starting point will be two
separate substances—possibly
elements
 Decomposition—one reactant breaks
apart
 Starting point will be only one
reactant—look for heat or electricity
Synthesis
 Be familiar with diatomics
 Sulfur and phosphorus occur as S8
and P4.
 Nonmetal oxides + water make
acids.
 Metal oxides + water make bases.
Synthesis
 Sulfur is burned in oxygen.
 Sulfur (VI) oxide is added to water
 Calcium oxide is added to water.
Decomposition
 Know some special types:
 Metal carbonates metal oxides &
CO2
 Metal hydroxides  metal oxides and
water
 Metal chlorates metal chlorides and
oxygen gas
 Hydrogen peroxide  H2O & O2
Decomposition
 Solid potassium chlorate is heated.
 Solid aluminum oxide is heated.
 An electric current is passed through
molten sodium chloride.
Solid potassium
chlorate is strongly
heated.
Redox
Several starting points:
 Key compounds or ions: MnO4-, H2O2,
Cr2O72-, HNO3
 Metals with multiple oxidation states:
Sn2+, Sn4+, Cr2+, Cr3+, Cr6+, etc.
 Acidic or basic conditions
 Sometimes halogens, i.e. I-, IO-, IO2-,
etc.
Redox
 In redox, one reactant gains
electrons and one reactant loses.
 LEO says GER:
 Losing electrons = oxidation
 Gaining electrons = reduction
 One process cannot occur without
the other!
Reducing Agents
 Cause reduction in something else
while being oxidized themselves
 Electron donors (any species that
loses electrons)
 Metal atoms
 Negative ions
 Positive metal ions that may still be
able to lose more electrons
Oxidizing Agents
 Cause oxidation in something else
while being reduced themselves
 Electron acceptors
 Nonmetal atoms
 Positive ions
 Permanganates, peroxides and
dichromates
Autooxidation
 Some species such as peroxide are
self-oxidzining
 One species is both oxidized and
reduced
 Ex: H2O2  H2O + O2
Acidic or Basic Conditions
 See page 833
 Reactions may be different under
acidic or basic conditions
 Use the reduction potential table
for reference
To Predict:
 Single replacement—follow normal
pattern
 Synthesis—follow normal pattern
 All others:
 Think about the most stable state of a
substance (i.e. K vs K+)
 Use the standard reduction potential
table to help
An acidic solution of potassium
dichromate is added to a
solution of iron II nitrate
 Find acidified dichromate as a reactant
(oxidizing agent—It is reduced.)
 Find iron II as a product. (It is oxidized.)
 Reverse the one that is oxidized (iron II).
 Combine and balance.
A strip of copper is immersed
in dilute nitric acid
 Copper begins with ox. # of zero; can
only form positive ions; must be
oxidized; reverse reaction.
 NO3- + 4H+  NO + 2H2O
SUMMARY
 Two uncombined reactants—
synthesis
 Single reactant—decomposition
 Water as a reactant—metals &
metal oxides produce bases;
nonmetals and nonmetal oxides
produce acids.
SUMMARY
 Acid & base reactants (including
salts)—neutralize to salt and water
 Two salt solutions—look for a
precipitate or a gas produced
 Combustion of hydrocarbon—
produce CO2 & H2O
 Solid metal placed in solution—
single replacement/redox rxn
SUMMARY
 Transition metal with ammonia
(NH3), hydroxide (OH-), cyanide
(CN-), or thiocyanate (SCN-)—
form complex ions in which ions
attach to metals
 Doesn’t matter how many—just get
charge correct
Special Notes
 If a product would be carbonic acid
(H2CO3), it will break down into CO2
and H2O.
 If a product would be ammonium
hydroxide, it would break down into
NH3 and H2O.
 Ammonium carbonate decomposes
into CO2, NH3 & H2O.
EXAMPLES
 A piece of aluminum metal is added
to a solution of silver nitrate.
 Al + 3Ag+  Al3+ + 3Ag
 A piece of solid bismuth is strongly
heated in oxygen.
 4Bi + 3O2  2Bi2O3 (or another
oxide)
EXAMPLES
 An excess of sodium hydroxide solution

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is added to a solution of magnesium
nitrate
Mg2+ + 2OH-  Mg(OH)2
Solid lithium hydride is added to water
LiH + H2O  LiOH + H2
A concentrate solution of ammonia is
added to a solution of zinc nitrate
Zn2+ + 4NH3  Zn(NH3)42+
Examples
 Concentrate hydrochloric acid is
added to solid manganese II sulfide.
 2 H+ + MnS  H2S + Mn2+
 Excess chlorine gas is passed over hot
iron filings.
 3Cl2 + 2Fe  2FeCl3 (possibly FeCl2)
Examples
 Solid ammonium carbonate is
heated
 (NH4)2CO3  CO2 + NH3 + H2O
 Equal volumes of 0.1 M sulfuric acid
and 0.1 M potassium hydroxide are
mixed
 H+ + OH-  H2O
EXAMPLES
 Propanol is burned completely in air
 2CH3CH2CH2OH +9 O2 6 CO2 + 8H2O or

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2C3H7OH +9 O2 6 CO2 + 8 H2O
A solid sample of magnesium carbonate is
heated strongly.
MgCO3  MgO + CO2
Ethene gas is bubbled through a solution
of bromine.
C2H4 + Br2  C2H4Br2