Download TIPS for NET-IONIC EQUATIONS A.P. Chemistry (long form)

TIPS for NET-IONIC EQUATIONS
(long form)
A.P. Chemistry
Traditionally students have done consistently poor on the section of the A.P. exam
that deals with writing net-ionic reactions. To succeed in this section requires
patience, practice, and a good deal of memorization. However, it is one of the only
sections that we KNOW will be tested on the AP exam. To prepare yourself for
this section gives you a competitive advantage on AP exam day.
On the actual exam you will have 10 minutes (calculator-free) to write 3 net-ionic
equations and answer a question for each. A survey of previous A.P. exams indicates
that most reactions given fall into one of these 5 rather broad categories:
1. double replacement
2. redox or single replacement
3. composition or decomposition
4. combustion or organic
5. complex ion formation
6. other*
* Mr. Schertz, you said there were “5 rather broad categories”. What gives? Well,
if you find a reaction that falls into the “other” category, remember, you can’t know
everything. Write reactants, guess products, move onto the question – whatever?!
BASIC INFO BEFORE WE GET STARTED:
Beginning at the beginning we state the "obvious". Most compounds can be divided
into about 6 categories:
1. acids (formulas begin with H except for organic acids which can be written
starting with H, but are often written ending in -COOH)
2. bases (ionic compounds ending in OH except for ammonia, NH3, and organic
bases which are similar to ammonia and contain nitrogen)
3. metal oxides (binary compounds of a metal and oxygen)
4. non-metal oxides (binary compounds of a non-metal and oxygen)
5. salts (compounds of metals or ammonium that are NEITHER bases NOR oxides)
6. other (many compounds are “other”)
What should you know right now about the categories above?
ACIDS : "strong" acids are HCl, HBr, HI, HClO4, HNO3, H2SO4 (only first H+)
BASES : "strong" bases are also few in number and should be learned: alkali metal
hydroxides AND large alkaline earth metal hydroxides Ca(OH)2, Sr(OH)2, Ba(OH)2
SALTS : learn and be able to recall solubility rules
This is the beginning. It is the “vocabulary” of chemical reactions. If you don’t
know it, you will fail to learn how to write chemical reactions.
DOUBLE REPLACEMENT REACTIONS
These reactions begin with two reactant compounds and produce two product
compounds. They typically occur when the reactants are acids, bases or salts.
The products are predicted by simply switching the cations of the two reactants.
EXAMPLES
HCl + NaOH → NaCl + H2O
AgNO3 + NaCl → AgCl + NaNO3
The AP exam requires net-ionic equations for full credit. The intent of the netionic equation is to represent reactant and product species as accurately as
possible, eliminating spectators that do not participate in the reaction. Thus the
two examples above in net-ionic form would be:
NET-IONIC
H+ + OH- → H2O
Ag+ + Cl- → AgCl
[it is worth noting that on that AP exam there are never "no-reaction" options]
PRACTICE:
1. dilute sulfuric acid is added to a solution of barium acetate
2. solutions of sodium phosphate and calcium chloride are mixed
3. hydrogen sulfide gas is bubbled through a solution of silver nitrate
4. manganese(II) nitrate solution is mixed with sodium hydroxide solution
5. solutions of zinc sulfate and sodium phosphate are mixed
6. solutions of silver nitrate and sodium chromate are mixed
7. dilute solutions of lithium hydroxide and hydrobromic acid are mixed
8. a solution of ammonia is added to a dilute solution of acetic acid
9. a solution of sulfuric acid is added to a solution of barium hydroxide
10. ammonia gas and carbon dioxide are bubbled into water
SINGLE REPLACEMENT REACTIONS
Single replacement reactions are a type of redox reaction in which an
element reacts with an ionic compound or water. In that compound there must be
an element similar in some way to the reacting element but less reactive. The more
reactive element takes the place of the less active element in the compound and
the less active element is "replaced":
EXAMPLES
Cl2+ 2 KBr → 2 KCl + Br2
Zn + CuSO4 → ZnSO4 + Cu
NET IONIC
Cl2 + 2 Br-  2 Cl- + Br2
Zn + Cu2+  Zn2+ + Cu
IMPORTANT NOTE: balancing redox is a little tricky – balance elements AND
charges. If you forget how to do ½ reactions or it takes too long, just write the
formulas and skip balancing.
Note that the activity of a given element is most easily predicted with a standard
reduction potential table. For metal displacements the metal with the more
positive reduction potential will be replaced. For halogens, the displacement order
follows the periodic table, fluorine being the most active.
PRACTICE:
1. calcium metal is added to a dilute solution of hydrochloric acid
2. liquid bromine is added to a solution of potassium iodide
3. magnesium turnings are added to a solution of iron(III) chloride
4. hydrogen gas is passed over hot iron(III) oxide
5. small chunks of solid sodium are added to water
REDOX (not single replacement) REACTIONS
We have already looked at simple reactions which involve changes in oxidation
number. The non-trivial sort require knowledge of common oxidizer/reducer pairs
or at the very least some common-sense elimination of unlikely products and then
guessing. Memorizing the "common" pairs may help but you will probably get
farther by trying to reason through the process since there is no guarantee that
the pairs you memorize will be used on the exam.
IMPORTANT NOTE: Redox will be especially difficult because of balancing – don’t
dwell too long on the concept. If you can quickly balance with ½-reactions, then do so.
If you forget or it takes too long, then write the formulas and skip the balancing.
Some essential principles to keep in mind are these:
 elements in their highest positive oxidation state can ONLY be reduced
 elements in their lowest oxidation state (0 for metals, negative for non-metals,
corresponding to distance from noble gases) can ONLY be oxidized
 intermediate oxidation states can go either way!!!
 “acidified” or “acidic” solutions: H+ should be a reactant; water is a product
 “basic” conditions or strong base: OH- should be a reactant; water is a product
 occasionally the acid anion or base cation may precipitate with a product ion
(know solubility rules)
common oxidizers [remember, oxidizers will become reduced]
MnO4- (in acid) → Mn2+
MnO4- (in neutral or basic) → MnO2
MnO2 (in acid) → Mn2+
Cr2O72- (in acid) → Cr3+
NO3- (from conc. nitric acid) → NO2
NO3- (from dilute nitric acid) → NO
H2SO4 (hot, concentrated) → SO2 [if not hot and conc., this acts like HCl or other
strong acids]
metal cations → lower charge cations or free metals
free halogens → halide ions
common reducers [remember, reducers will become oxidized]
halide ions → free halogens
free metals → metal cations
metal cations → higher charge cations
O22- → O2
SO32- (or SO2) → SO42NO2- → NO3free halogens (dil. basic) → hypohalite ions [like XO-]
free halogens (conc. basic) → halate ions [like XO3-]
That seems like a lot to remember, but if you look at the list carefully you
will see that none of the species are really exotic (with the possible
exception of the hypohalite and halate ions). If you don't recognize a common
oxider or reducer in a redox reaction, then it’s probably NOT redox, or you
probably deserve to miss the question – skip it.
How important are these? On the real exam there are 3 reactions. Here is the
number of non-trivial redox reactions for the past few years of exams. Judge for
yourself. OF COURSE, I will quiz and test you on them!!
1998 – none
2002 – none
2006 – none
1999 – 1
2003 – none
2007 – none
2000 – none
2004 – 1
2008 – none
2001 - 1
2005 – none
2009 - 1
PRACTICE:
1. hydrogen peroxide is added to an acidified solution of potassium dichromate
2. sulfur dioxide gas is bubbled through an basic solution of potassium
permanganate
3. a solution containing tin(II) ions is added to acidified potassium dichromate
solution
4. powdered iron is added to a solution of iron(III) sulfate
5. copper(II) sulfide is added to dilute nitric acid
6. concentrated hydrochloric acid is added to solid manganese(IV) oxide and heated
7. chlorine gas is bubbled through a solution of dilute sodium hydroxide
8. solutions of potassium iodide and potassium iodate are mixed
9. solid silver is added to dilute nitric acid
COMPOSITION or DECOMPOSITION REACTIONS
Composition (synthesis) reactions can refer to the very simple reaction of two
elements (a redox process) to form a compound. Net-ionic form is usually not
required because typically water would be absent:
EXAMPLES
S + O2 → SO2
3 Mg + N2 → Mg3N2
However, some compounds also combine to form other compounds. These are
sometimes less obvious. There are some common categories which appear on the
exam. It would be helpful to recognize them:
1. metal oxide + water → a base (non-redox)
2. non-metal oxide + water → an acid (non-redox)
3. metal oxide + non-metal oxide → salt (non-redox)
EXAMPLES (net ionic)
CaO + H2O → Ca2+ + 2 OHSO2 + H2O → H2SO3
CaO + SO2 → CaSO3
Decomposition reactions follow a similar pattern (although reversed). When a
simple binary compound is decomposed the elements result. These are not a
problem. When a ternary compound decomposes (usually by heating) there are some
general rules:
1. metal hydroxide → metal oxide + water
2. acid containing oxygen → non-metal oxide + water
3. salt containing oxygen → metal oxide + non-metal oxide
EXAMPLES (net ionic)
Ca(OH)2 → CaO + H2O
H2CO3 → CO2 + H2O
CaCO3 → CaO + CO2
[NOTE: the 2nd example above occurs under standard conditions. “H2CO3” should
never be written as a reactant or product on the reactions section of the AP exam.
Always write “CO2 + H2O”. Think of it as the “Coke rule”.]
PRACTICE:
1. solid calcium oxide is exposed to a stream of carbon dioxide gas
2. a mixture of solid calcium oxide and tetraphosphorus decaoxide is heated
3. sulfur trioxide gas is added to excess water
4. solid magnesium carbonate is heated
5. solid potassium chlorate is heated in the presence of a manganese(IV) oxide catalyst
6. a solution of hydrogen peroxide is heated
COMBUSTION REACTIONS
For hydrocarbon combustions, the simple rule is that burning (combining with O2)
produces carbon dioxide and water. This is true whether the hydrocarbon contains
oxygen or not. However, if halogens or nitrogen are present, oxides of these
elements may form as well. These are difficult to predict, although in the case of
nitrogen, NO or NO2, would always be a reasonable guess.
Other combustion reactions tested on the AP exam are typically metals in the
presence of oxygen – treat like a composition reaction.
EXAMPLES
2 C6H14 + 19 O2 → 12 CO2 + 14 H2O
4 Li + O2 → 2 Li2O
Note that net-ionic versions of these reactions really don't exist since liquid water
is generally not present.
PRACTICE:
1. propanol is burned completely in air
2. acetic acid is burned in an excess of oxygen
3. calcium metal is heated strongly in the presence of oxygen
ORGANIC REACTIONS
The subject of organic reactions is very large and rather complicated. The
complicated topics are NOT the type of things you should be worrying about at this
level. Later this semester we will learn more detailed rules for NAMING organic
compounds that WILL be helpful on the AP test, though:
Information to know by AP exam day (won’t be quizzed or tested in January):
1) general formulas for alkanes, alkenes and alkynes AND how Greek prefixes are
used to name them according to the length of the carbon chain
2) some common functional groups and how they are attached (usually) to the
carbon chain
Historically, some types of organic reactions that have shown up on AP exams:
1. combustion (products are always the same, CO2 and H2O)
2 C6H14 + 19 O2→ 12 CO2 + 14 H2O
2. addition (requires alkenes or alkynes, halogens, halogen acids or just hydrogen)
C3H6 + Cl2→ C3H6Cl2
3. substitution (a hydrogen atom in the molecule is "displaced" by something)
CH4 + 4 Br2 → CBr4 + 4 HBr
4. esterification (carboxylic acid + alcohol → ester + water)
CH3OH + CH3COOH → CH3COOCH3 + H2O
5. saponification ( ester + base → salt + alcohol)
CH3COOC2H5 + OH- → CH3COO- + C2H5OH (an argument could be made for
writing this last one in molecular form, but the NaOH is most likely aqueous
and therefore water would be present)
PRACTICE:
1. a sample of 2-butene is treated with hydrogen bromide gas
2. an excess of chlorine gas is added to acetylene gas (ethyne)
3. acetic acid is refluxed with ethanol for several hours
4. propene and bromine gases are mixed
5. propanol is burned completely in air
6. methane is mixed with excess bromine gas
7. propene reacts with water in the presence of a catalyst
COMPLEX ION REACTIONS
Reactions of coordination compounds and ions are not covered in depth on the exam
but you will sometimes see them in the reaction-writing section and they are easy
enough to complete with a few basic principles in mind. Most can be recognized by
the choice of reactants: generally a transition metal ion or compound (also the
amphoteric species from Group 13, such as Al) and a source of ligands. The
most common ligands involved in questions are ammonia and the hydroxide ion. Key
to such questions is the word "excess", indicating that enough of the complexation
agent has been added to eliminate the possibility of precipitation of lessercoordinated species.
One of the hurdles to get over is some knowledge of the likely coordination
numbers for metal ions. It may be helpful to know that often the coordination
number is typically twice the cation charge. In any case, you will seldom lose points
because you used an incorrect coordination number . . . as long as the overall charge
of the complex ion agrees with the oxidation number of the metal/ligand complex!!
EXAMPLES
1. complexation of a soluble salt
a concentrated solution of ammonia is added to a solution of copper(II) chloride
4 NH3 + Cu2+ → Cu(NH3)42+
2. complexation of an insoluble salt
excess concentrated potassium hydroxide is added to zinc hydroxide
2 OH- + Zn(OH)2 → Zn(OH)423. destruction of a complex by acid/base neutralization
dilute hydrochloric acid is added to a solution of diamminesilver nitrate
2 H+ + Cl- + Ag(NH3)2+ → AgCl + 2 NH4+
PRACTICE:
1. excess dilute nitric acid is added to a solution of tetramminecadmium(II) ion
2. an excess of ammonia gas is bubbled through a solution saturated with silver
chloride
3. a concentrated solution of ammonia is added to a suspension of zinc hydroxide
4. a solution of ammonium thiocyanate is added to a solution of iron(III) chloride
“OTHER” REACTIONS (hydrolysis examples)
Didn’t I say you could SKIP this section?? Well, some of you need more . . . I know.
We will not cover any “other” reactions in class. In fact, on this sheet I will only
cover the most common “other” types of reactions that might show up on the AP
exam.
Hydrolysis Reactions are generally understood as reactions of something with
water. Reaction should not be confused with dissolving. Many compounds dissolve
in water and are essentially unchanged. In a hydrolysis reaction, some new
substance or species forms that is not found in the original compound.
The most common examples of hydrolysis are reactions of the anions of weak acids
or the cations of weak bases with water. These are typical of the processes which
occur when salts of these compounds enter water.
You will better understand this type of reaction when we get to acid/base
chemistry and the behavior of weak acids/bases. The net ionic reaction for
something as innocent as the dissolving of sodium acetate in water:
NaCH3COO + H2O  CH3COOH + OH- + Na+
(the double arrow is not necessary, but would be impressive!)
The general form for such a reaction might be written:
anion of weak acid + water  weak acid + hydroxide ion
A similar reaction could be written for the cations of weak bases, although there is
really only one common one (NH4+):
cation of weak base + water  weak base + hydronium ion
Historically, non-metallic halides have been popular as reactants in hydrolysis
reactions. Chlorides of phosphorus or sulfur are common. You can just skip these,
or you can hazard a guess. In general, if water is thought of as HOH, the H+ will
combine with the more (or most) electronegative atom in the non-metal halide. This
gives one product. The other product consists of the remaining elements.
EXAMPLES
phosphorus trichloride + water →
PCl3 + 3 H2O → 3 H+ + 3 Cl- + H3PO3 [the H+ and Cl- are HCl in net-ionic form]
phosphorus pentachloride + water →
PCl5 + 4 H2O → 5 H+ + 5 Cl- + H3PO4
How can you tell which phosphorus acid forms????? The oxidation number of
phosphorus remains constant. This is not a redox reaction.
sulfur tetrachloride + water →
SCl4 + 3 H2O → 4 H+ + 4 Cl- + H2SO3
Anyway, here are some “other” PRACTICE:
1. phosphorus tribromide is added to water
2. gaseous hydrogen chloride is dissolved in water
3. phosphorus oxytrichloride is added to water
4. solid sodium sulfite is added to water
5. sodium phosphate crystals are added to water
6. solid sodium cyanide is added to water
7. solid aluminum sulfide is added to water
Answers to Practice Equations
DOUBLE REPLACEMENT
1. H+ + HSO4- + Ba2+ + 2 CH3COO- → BaSO4 + 2 CH3COOH
2. 3 Ca2+ + 2 PO43-→ Ca3(PO4)2
3. H2S + 2 Ag+ → 2 H+ + Ag2S
4. Mn2+ + 2 OH- → Mn(OH)2
5. 3 Zn2+ + 2 PO43- → Zn3(PO4)2
6. 2 Ag+ + CrO42- → Ag2CrO4
7. H+ + OH- → H2O
8. NH3 + CH3COOH → NH4+ + CH3COO9. H+ + HSO4- + Ba2+ + 2 OH- → BaSO4 + 2 H2O
10. NH3 + CO2 + H2O → NH4+ + HCO3-
SINGLE REPLACEMENT
1.
2.
3.
4.
5.
Ca + 2 H+ → Ca2+ + H2
Br2 + 2 I- → I2 + 2 Br3 Mg + 2 Fe3+ → 3 Mg2+ + 2 Fe OR (Mg + 2 Fe3+ → Mg2+ + 2 Fe2+)
3 H2 + Fe2O3 → 2 Fe + 3 H2O
2 Na + 2 H2O → 2 Na+ + OH- + H2
REDOX (not single replacement)
1.
2.
3.
4.
5.
6.
7.
8.
9.
3 H2O2 + Cr2O72- + 8 H+ → 2 Cr3+ + 3 O2 + 7 H2O
3 SO2 + 2 MnO4- + 4 OH- → 2 MnO2 + 3 SO42- + 2 H2O
3 Sn2+ + 14 H+ + Cr2O72- → 3 Sn4+ + 2 Cr3+ + 7 H2O
Fe + 2 Fe3+ → 3 Fe2+
3 CuS + 8 H+ + 2 NO3- → 3 S + 2 NO + 3 Cu2+ + 4 H2O
4 H+ + 2 Cl- + MnO2 → Mn2+ + Cl2 + 2 H2O
Cl2 + 2 OH- → H2O + ClO- + Cl6 H+ + 5 I- + IO3- → 3 I2 + 3 H2O
3 Ag + 4 H+ + NO3-  3 Ag+ + NO + 2 H2O
COMPOSITION and DECOMPOSITION
1.
2.
3.
4.
5.
6.
CaO + CO2 → CaCO3
6 CaO + P4O10 → 2 Ca3(PO4)2
SO3 + H2O → H+ + HSO4MgCO3 → MgO + CO2
2 KClO3 → 2 KCl + 3 O2
2 H2O2 → 2 H2O + O2
COMBUSTION
1. 2 C3H7OH + 9 O2 → 6 CO2 + 8 H2O
2. CH3COOH + 2 O2 → 2 CO2 + 2 H2O
3. 2 Ca + O2 → 2 CaO
ORGANIC
1.
2.
3.
4.
5.
6.
7.
C4H8 + HBr → C4H9Br
C2H2 + 2 Cl2 → C2H2Cl4
CH3COOH + C2H5OH → CH3COOC2H5 + H2O
C3H6 + Br2 → C3H6Br2
2 C3H7OH + 9 O2 → 6 CO2 + 8 H2O
CH4 + 4 Br2 → CBr4 + 4 HBr
C3H6 + H2O → C3H7OH
COMPLEX IONS
1.
2.
3.
4.
4 H+ + Cd(NH3)42+ → Cd2+ + 4 NH4+
2 NH3 + AgCl → Ag(NH3)2+ + Cl4 NH3 + Zn(OH)2 → Zn(NH3)42+ + 2 OHSCN- + Fe3+ → FeSCN2+ [other species up to (SCN)6 accepted]
“OTHER” REACTIONS (hydrolysis examples)
PBr3 + 3 H2O → H3PO3 + 3 H+ + 3 BrHCl + H2O → H3O+ + Cl- [Arrhenius version also accepted]
POCl3 + 3 H2O → H3PO4 + 3 H+ + 3 ClNa2SO3 + H2O → 2 Na+ + HSO3- + OHNa3PO4 + H2O → 3 Na+ + OH- + HPO42[OR Na3PO4 + 2 H2O → 3 Na+ + 2 OH- + H2PO4-]
6. NaCN + H2O → Na+ + HCN + OH7. Al2S3 + 6 H2O → 2 Al(OH)3 + 3 H2S
1.
2.
3.
4.
5.