Reactions
8
Every AP test has a few reaction questions. Given some reactants, you must write a balanced, net
ionic equation. We will start with unbalanced equations.
Net Reactions
Most materials in AP reactions are dissolved. You don't need to figure this out; the problem says
so. In such cases all you write is the ion which reacts. Thus when sodium chloride solution reacts
with silver nitrate solution, you write:
Ag+ + Cl–  AgCl
The sodium and the nitrate ions are spectators and are not included.
Sometimes materials are added as solids. Generally this is stated explicitly, as in: "solid zinc
nitrate," or "a suspension of copper(II) hydroxide." Sometimes, however, you know it's a solid
because it isn't a solution, as in "excess silver acetate is added to potassium chloride solution." In
such cases you write out the whole compound, and not just the part that reacts.
AgC2H3O2 + Cl–  AgCl + C2H3O 2–
Note that the potassium is a spectator and is omitted. However, the acetate ion has to be included
on the right because it was part of the solid on the left.
Although we are starting with unbalanced equations, any element which is on one side of an
equation must also appear on the other. The following equation is incorrect because it has no
oxygen on the right:
CrO42– + Fe2+  Cr3+ + Fe3+
But you can't just add O2. That would mean than oxygen was going from an oxidation state of -2
(in chromate) to an oxidation state of 0 in oxygen gas. Instead you have to add water to the right.
This, then, means you have to add H+ to the left, giving you:
CrO42– + Fe2+ + H+  Cr3+ + Fe3+ + H2O
To make things easier, the reactions have been divided into categories. These categories are
arbitrary, and could have been done differently. But I doubt that it would be VERY different.
Writing reactions when you know the category is fairly easy. You will start by learning reactions
within their category. Then, after you have had time to practice, we will start mixing the
categories.
Precipitation Reactions
These are reactions in which a precipitate (an insoluble material) forms. To do them you must
know the solubility rules. (A handout which you have received). What is important here is not
what is soluble but what is insoluble, since it is the insoluble materials that form precipitates. You
need to know the insoluble materials in Rule 3 (silver and lead halides), Rule 4 (calcium, barium
and lead sulfates), and rule 5 (most hydroxides, sulfides, sulfites, carbonates, phosphates and
chromates). So if barium nitrate solution is mixed with potassium sulfate solution, Rule 4 says that
barium sulfate is insoluble and:
-1-
Ba2+ + SO42–  BaSO4
People worry about what to do with the slightly soluble (Rule 5) calcium and barium hydroxides.
Don't. When they are reactants, the question will always make clear whether they should be
treated as soluble or insoluble. Thus, mixing solutions of calcium hydroxide and sodium phosphate
would give insoluble (Rule 5) calcium phosphate. Since it says "...solutions of calcium
hydroxide..."
Ca2+ + PO43–  Ca3(PO4)2
When calcium or barium hydroxides are products, they are considered insoluble.
Some Examples
Solutions of potassium carbonate and zinc nitrate are mixed.
Zn2+ + CO32–  ZnCO3
A solution of copper(II) sulfate is added to a solution of barium hydroxide.
Cu2+ + SO42– + Ba2+ + OH–  Cu(OH)2 + BaSO4
Hydrogen sulfide gas is bubbled through a solution of lead(II) nitrate.
H2S + Pb2+  PbS + H+
Acid/Base Reactions
You need to know the strong acids -- HCl, HBr, HI, HNO3, H2SO4. These are considered to be
completely dissociated. All other acids are weak and are written undissociated. You also need to
know that an acid and a base react to form water
HF + OH–  H2O + F–
and that the anion from a weak acid is protonated by water (or H+).
NO–2 + H2O  HNO2 + OH–
Some reactions specify certain numbers of moles of acid and base. This usually leads to the partial
protonation of a polyprotic acid. Thus if sodium phosphate is mixed with an equal number of
moles of hydrochloric acid solution:
H+ + Na3PO4  Na+ + HPO42–
You should know that a carbonate (or bicarbonate) reacts with acid to form carbon dioxide and
water.
CaCO3 + H+  Ca2+ + CO2 + H2O
This is because protonation of the carbonate ion forms H2CO3, which comes apart to form CO2 and
H2O. Sulfides react with acid to form hydrogen sulfide:
Fe2S3 + HC2H3O2  Fe+3 + C2H3O–2 + H2S
Finally, the oxides and hydroxides of certain metals, called amphoteric metals, can react as if they
are either acids or bases. The only amphoteric metals you need to remember are zinc and
aluminum. Thus zinc hydroxide dissolves in sodium hydroxide solution:
Zn(OH)2 + OH–  Zn(OH)42–
Every once in a while the AP exam asks about a gas phase Lewis acid/base reaction:
BF3 + NH3  F3B–NH3
Every question I've seen has used a boron trihalide or trihydride as the Lewis acid.
-2-
Some More Examples
A solution of sodium hydroxide is added to a solution of ammonium chloride.
OH– + NH4+  NH3 + H2O
Dilute hydrochloric acid is added to a potassium carbonate solution.
H+ + CO32–  H2O + CO2
Sodium hydroxide solution is added to a precipitate of aluminum hydroxide.
Al(OH)3 + OH–  Al(OH)63–
Dilute solutions of acetic acid and ammonia are mixed.
HC2H3O2 + NH3  NH4+ + C2H3O–2
Redox Reactions
In a redox reaction one species is oxidized (its oxidation state becomes more positive) and another
species is reduced (its oxidation state becomes less positive). To do these problems, you must
know the common oxidizing and reducing agents mentioned in Table 3 -- Some Useful Reactions.
You should realize that oxidizing agents are themselves reduced, while reducing agents are
oxidized. It sounds funny but it makes sense. It helps if you are familiar with the common
oxidizing and reducing agents. For example, hydrogen peroxide is usually an oxidizing agent (and
is therefore reduced):
H2O2 + Fe2+  H2O + Fe2+
However, it can occasionally be a reducing agent (and be oxidized):
H2O2 + Cr2O72–  O2 + Cr3+ + H+
Look at the way the oxidation state of oxygen changes in these reactions.
Permanganate, chlorate, bromate, iodate, chlorine, bromine or oxygen are always oxidizing
agents. Other oxidizing agents are less certain because they can do other things besides oxidize.
You must further remember that permanganate forms Mn2+ in acidic solutions and MnO2 in basic
solutions.
MnO–4 + I– + H+  Mn2+ + I2 + H2O
MnO–4 + SO32– + OH–  MnO2 + H2O + SO42–
If the acidity or alkalinity is not specified, assume an acid solution.
Some More Examples
A solution of tin(II) sulfate is mixed with acidified potassium dichromate.
Sn2+ + Cr2O72– + H+  Sn4+ + Cr3+ + H2O
Hydrogen peroxide solution is added to a solution of iron(II) sulfate
H2O2 + Fe2+  H2O + Fe3+
In grading: "RWO" will mean reduction without oxidation; "OWR" will mean oxidation
without reduction.
-3-
Metal Displacement Reactions
These are a particular type of redox reaction in which a metal, in its elemental state is oxidized.
Usually this combines with another metal being reduced. (Thus one metal displaces another.) If a
piece of copper is put into a silver nitrate solution, the copper replaces the silver:
Cu + Ag+  Cu2+ + Ag
In real life you would have to know which metal is more active in order to determine whether the
replacement reaction goes. But here, all the reactions go. So you need not worry about the activity
series. Also I've included the reaction of metals with acid or water, since the reaction is the same
except now it's hydrogen which is reduced. So the reaction of aluminum with hydrochloric acid is:
Al + H+  Al3+ + H2
And the reaction of calcium with water is:
Ca + H2O  Ca(OH)2 + H2
Metal displacement reactions are easy to recognize because they involve a piece of metal.
Dissolving metals in nitric acid or in hot sulfuric acid is more complicated, since these acids also
act as oxidizing agents. In each case the nitrogen or sulfur is reduced to the oxide of a lower
oxidation state. This question usually involves copper, lead, or silver. These metals are less
reactive than H+ and, therefore, too inert to dissolve in ordinary acids (e.g. HCl). They need the
additional oxidizing power of the (acidified) nitrate or the (hot, acidified) sulfate.
H+ + NO–3 + Ag  Ag+ + H2O + NO (or NO2)
H+ + SO42– + Cu  Cu2+ + H2O + SO2
Some More Examples
A piece of aluminum metal is added to silver acetate solution.
Al + Ag+  Al3+ + Ag
Combustion
These are easy because the questions tells you that something is being burned. You take the
material being burned, add oxygen gas, and convert the carbon to carbon dioxide and the hydrogen
to water.
C2H5OH + O2  CO2 + H2O
If they are present, chlorine is converted to HCl and most other elements are converted to their
oxides.
Some More Examples
Carbon disulfide vapor is burned in oxygen.
CS2 + O2  CO2 + SO2
Gaseous silane (SiH4) is burned in oxygen.
SiH4 + O2  SiO2 + H2O
-4-
Reaction with Water
We are principally concerned about oxides reacting with water. Metal oxides react to form
hydroxides. In fact, forming basic hydroxides is considered to be a property of metal ions and is
shown on some periodic tables by a blue color. So adding water to potassium oxide would give:
K2O + H2O  K+ + OH–
Non-metal oxides, however, react with water to form acids. They are often called "acid
anhydrides" and are shown by a red color on some periodic tables. Several oxides which form
acidic solutions are shown on the "Some Useful Reactions" sheet, but one example is sulfur
dioxide:
SO2 + H2O  H2SO3
If you are paying attention, you will have no trouble recognizing a non-metal oxide and knowing
that it forms an acid. But which acid does it form? Sulfur dioxide was easy because you added one
molecule of sulfur dioxide to one molecule water to form one molecule of sulfurous acid. But
suppose the question asked about N2O5? You know two nitrogen oxy-acids, HNO3 and HNO2.
Which one does it form? This type of question is answered by determining the oxidation state of
the non-metal and then finding which acid has the non-metal in the same oxidation state. In N2O5
nitrogen is in the +5 oxidation state. This matches with the N(+5) in HNO3. Using similar logic:
P4O10 + 6 H2O  4 H3PO4
SO3 + H2O  H2SO4
Although oxides are the most important cases, nitrides react similarly:
Mg3N2 + H2O  Mg(OH)2 + NH3
This actually makes sense if you realize that the nitride anion reacts like the anions of other weak
acids. It doesn't matter whether you deal with oxide, nitride, hydride, fluoride, acetate or any other
such anion. The anion removes protons from water, leaving hydroxide ions. Thus
LiH + H2O  Li+ + H2 + OH–
KF + H2O  K+ + HF + OH–
And calcium carbide, which unexpectedly contains the C22– anion, reacts as shown below.
CaC2 + H2O  Ca2+ + C2H2 + OH–
Remember this!
Some More Examples
Carbon dioxide is bubbled through a sodium hydroxide solution.
CO2 + OH–  CO32– + H2O
Sodium phosphide is dissolved in water.
Na3P + H2O  Na+ + H3P + OH–
-5-
Complexation
A complex is a compound which forms between a metal and an electron donor called a "ligand." In
AP reactions, the ligand is usually either ammonia or cyanide. Thus, if solutions of sodium cyanide
and silver nitrate are mixed:
Ag+ + CN–  Ag(CN)–2
The negative charge on the complex ion is the sum of one plus (for the Ag+) and two minuses (for
the CN–).
If a concentrated ammonia solution is added to a suspension of zinc hydroxide:
Zn(OH)2 + NH3  Zn(NH3)42+
Notice that the number of ligands, called the coordination number, varies. If you don't know
otherwise you should assume the coordination number is twice the metal's charge.
One special complexation reaction you should know is the reaction of iron(III) with thiocyanate
(SCN–):
Fe3+ + SCN–  Fe(SCN)2+
There's really more to this reaction than you see, but this is enough.
Still another special case is the amphoteric metals, aluminum and zinc, dissolving in base. The
products of this reaction can be viewed as hydroxide complexes:
Al(OH)3 + OH–  Al(OH)63–
Complexes in which the ligand can be protonated to form a weak acid (e.g. NH3 and CN–) come
apart in acid solution. For example:
Cu(NH3)42+ + H+  Cu2+ + NH4+
Some More Examples
A suspension of copper(II) hydroxide in water is treated with excess ammonia.
Cu(OH)2 + NH3  Cu(NH3)42+ + OH–
Electrolysis
The good thing about electrolysis reactions is that the questions tells you that electrolysis is
occurring. The bad thing is that even then some of the reactions are hard to figure out. The most
important thing to remember, though, is that you need both a reduction and an oxidation. Don't
forget! If a solution of potassium iodide is electrolyzed, you get:
I– + H2O  I2 + H2 + OH–
Oxidizing iodide is easy. But what is reduced? Potassium is much too difficult to reduce and
hydrogen is all that is left.
If a dilute solution of sulfuric acid is electrolyzed, you get:
H2O  H2 + O2
Here there is nothing which can be oxidized other than oxygen. Actually the sulfuric acid isn't a
reactant; it just makes the solution conductive. But the sulfate confuses many students. You must
realize that sulfate (where S is +6) cannot be oxidized and is usually not reduced.
-6-
Some Examples
A solution of copper(II) sulfate is electrolyzed using platinum electrodes.
Cu2+ + H2O  Cu + O2 + H+
Again, there is nothing in the solution except for oxygen which could be oxidized. The platinum is
there to show you that the electrodes are non-reactive.
Decomposition
When a compound reacts all by itself (usually with heat) it is decomposing. Some carbonates and
hydroxides do this. Thus, if calcium carbonate is heated:
CaCO3  CaO + CO2
If magnesium hydroxide is heated:
Mg(OH)2  MgO + H2O
But what happens when potassium nitrate is heated? Neither nitrogen oxides nor K3N are very
stable. So:
2 KNO3  2 KNO2 + O2
The 2002 AP test asked what happened when sodium bicarbonate was heated. This was tricky
because sodium carbonate does not decompose easily. Even knowing that, you still had to
balance the equation to figure out what happens:
NaHCO3  Na2CO3 + H2O + CO2
Most of the decomposition reactions which you will see involve ionic solids. They generally form
other ionic solids and NOT the ions themselves. Separate ions (e.g. Ca2+) exist only in solution.
Look at the examples above and you will what this means.
Some Examples
Hydrogen peroxide solution is warmed.
H2O2  H2O + O2
Combination
If heating a solid in an atmosphere of a specific gas causes it to react, the reaction is probably a
combination reaction. Thus heating calcium oxide in sulfur dioxide:
CaO + SO2  CaSO3
Similarly if magnesium metal is heated in nitrogen gas:
Mg + N2  Mg3N2
If titanium metal is heated in an atmosphere of oxygen:
Ti + O2  TiO2
Some Examples
Excess chlorine gas is passed over hot iron filings.
Fe + Cl2  FeCl3
Also acceptable would be FeCl2.
-7-
Some Other Important Reactions
Know the reactions in the last section of the "Some Useful Reactions" handout. They don't fit into
a category well, but they are important and have each been asked several times.
Some Examples
Chlorine gas is bubbled into a cold, dilute solution of potassium hydroxide.
Cl2 + OH–  OCl– + Cl– + H2O
Copper(II) oxide is heated in an atmosphere of hydrogen.
CuO + H2  Cu + H2O
Aluminum metal is placed in a solution of potassium hydroxide.
Al + OH–  Al(OH)6-3 + H2
Aluminum really does react with base the same way it reacts with acid. This was once a common
laboratory method for generating hydrogen.
Catalysis
If your reaction employs a catalyst, don't let it worry you. You can leave it out of the reaction,
because catalysts are not consumed. Thus, if potassium chlorate is heated in the presence of a
manganese dioxide catalyst:
2 KClO3  2 KCl + 3 O2
-8-