AP Chemistry
A part of the Advanced Placement Chemistry Examination on which the performance of candidates has been
disappointing through the years has been the question that asks candidates to provide formulas for the names of
reactants and then to write formulas for the products obtained as each indicated reaction takes place. In 2002,
for example, the average score on this part of the examination was only 5.5 out of a possible 15. Yet college
chemistry professors and AP Chemistry teachers are in agreement that the material involved is very important
for students of chemistry to be able to handle.
AP level textbooks cover this topic but it is scattered throughout the text; not confined to a single chapter. What
follows are a few suggestions on how to successfully predict the products to a chemical reaction. The College
Board requests that the reaction equations not be balanced, nor are physical states necessary. For purposes of
clarity the article will include physical states.
For brevity purposes this booklet does NOT cover nomenclature (writing and naming chemical compounds).
Students should review the nomenclature of salts, acids, bases, complex ions and basic organic compounds.
A useful classification of compounds is the following: (The definitions are reasonable for a water world such as
the one in which we live.)
1. Acids - compounds with formulas that begin with H
Examples: HCl, HNO3, H2SO4, H2CO3
2. Bases - compounds with formulas that end with OH
Examples: NaOH, Ca(OH)2, Fe(OH)3
3. Metal oxides - binary compounds of a metal and oxygen
Examples: CaO, Fe2O3, Li2O
4. Nonmetal oxides - binary compounds of a nonmetal and oxygen
Examples: SO2, NO3, P2O5
5. Salts - compounds of metals that are NOT bases or metal
oxides. Examples: NaCl, MgS, ZnSO4
6. Other compounds (Most compounds belong here.)
Examples: CH4, NH3, PCl3
Three of these categories need some elaboration to make further steps in our presentation reasonable:
1. ACIDS. The number of ôfrequentlyö encountered strong acids (acids that are present in solution very largely
as ions rather than as molecules) is small and one should know them by name and formula: HCl, HBr, HI,
HNO3, H2SO4, HClO4 As a first approximation, all other acids may be considered weak (present largely as
molecules).
2. BASES. The number of strong bases (bases that are present in solution largely as metal ions and hydroxide
ions rather than as molecules) is not large, and these substances should also be learned: LiOH, NaOH, KOH,
CsOH, RbOH, Ca(OH) 2, Sr(OH) 2, Ba(OH) 2. All other bases should be considered weak.
3. SALTS. The salts that are soluble in water include all of the salts of lithium, sodium, potassium, and
ammonium cations and of nitrate and acetate anions. All chlorides are soluble except those of silver, lead, and
mercury(I) ions. All sulfates are soluble except those of lead, calcium, strontium, and barium. All other salts
should be considered only slightly soluble.
I suggest encouraging students to develop their own devices for learning the lists in the notes above since a
studentÆs searching for devices that are useful will be much more effective in fixing the information than
having the devices presented by a teacher. And the student can have the pleasure of discovery.
When a student is representing a reaction that occurs in water solution, the substances that should be written as
ions are the strong bases, strong acids, and the soluble salts.
The reactions themselves can also be put into categories. A student who is able to decide about the category
into which a reaction fits by studying the nature of the reactants is the student who is most likely to be able to
predict products of reactions that have not been learned.
A. Reactions involving no changes in oxidation states.
Double replacement reactions - reactions where two compounds dissolved in water switch cations to form two
products; one of which is either a nondissociating, is insoluble or a gas.
General formula: AX + BY
BX + AY
a. Acid/base neutralization reaction. These reactions go to completion as long as one of the reactants is
ôstrong.ö The best answers eliminate ôspectator ionsö (ions that are not involved in the chemical reactions and
are not chemically changed from reactant to product side.)
(1) Strong acid + Strong base
salt + water
General reaction: HX(aq) + MOH(aq)
MX(aq) + H2O
Net Ionic reaction:
H2O
H+(aq) + OH-(aq)
Example: Hydrochloric acid + sodium hydroxide
HCl(aq) + NaOH(aq)
H+(aq) + OH-(aq)
NaCl(aq) + H2O
H2O
(2) Strong acid + weak base
salt + water
General reaction: HX(aq) + MOH(aq)
Net Ionic reaction:
H+(aq) + MOH(aq)
MX(aq) + H2O
M+(aq) + H2O
Example: Hydrochloric acid + magnesium hydroxide
HCl(aq) + Mg(OH)2(aq)
MgCl2(aq) + H2O
H+(aq) + Mg(OH)2 (aq)
Mg2+(aq) + H2O
(3) Weak acid + strong base
salt + water
General reaction: HX(aq) + MOH(aq)
Net Ionic reaction:
MX(aq) + H2O
HX(aq) + OH- (aq)
X- (aq) + H2O
Example: Hydrofluoric acid + potassium hydroxide
HF(aq) + KOH(aq)
HF(aq) + OH-(aq)
KF(aq) + H2O
F-(aq) + H2O
Precipitation reactions - one or more of the products is insoluble in water and reaction goes to completion. You
must be familiar with the solubility rules for compounds to successfully complete these predictions.
Solubility Rules for Salts in Water
All salts containing nitrate(NO3-) and chlorates (ClO3-) are soluble.
All salts containing Group 1 (alkali metals) ions are soluble.
All salts containing ammonium (NH4+) ion are soluble.
All acetate (C2H3O2-) salts are soluble except silver acetate
(Ag C2H3O2).
Chloride, bromide and iodide salts are soluble. Exceptions are salts containing the ions Ag+, Pb2+, and Hg22+
Most sulfate salts are soluble. Exceptions are BaSO4, PbSO4, HgSO4, Hg2SO4, Ag2SO4 and CaSO4
Hydroxide compounds are not soluble. Group I and ammonium hydroxides are soluble. Ba(OH)2, Sr(OH)2
and Ca(OH)2 are slightly soluble.
Sulfides (S2-),sulfites (SO32-) carbonates (CO32-), chromates (CrO42-), oxides (O2-) and phosphates (PO43-)
are not soluble. Group 1 and ammonium compounds of these ions are soluble.
General reaction: AX(aq) + BY(aq)
BX(aq) + AY(s)
Remember to remove all spectator ions from the final reaction equation.
Net Ionic reaction: A+(aq) + Y-(aq)
AY(s)
Example: Silver nitrate + sodium chloride
Reaction equation: AgNO3(aq) + NaCl(aq)
Net Ionic equation: Ag+(aq) + Cl-(aq)
NaNO3(aq) + AgCl(s)
AgCl(s)
2. Some combination reactions with a single product.
a. Metal oxide + water
heated
a base, the metal in the same oxidation state as in the oxide. Not reversed when
Group 1 metal oxides: M2O(s) + H2O
M+(aq) + OH-(aq)
Example: Sodium oxide + water
Na2O(s) + H2O
Na+(aq) + OH-(aq)
Group 2 metal oxides: MO(s) + H2O
M2+(aq) + OH-(aq)
Example: calcium oxide + water
CaO(s) + H2O
Ca2+(aq) + OH-(aq)
b. Nonmetal oxide + water
an acid, the nonmetal in the same oxidation state as in the oxide. Keep in
mind that these reactions are easily reversed at higher temperatures.
Example: Sulfur dioxide + water
SO2(g) + H2O
H2SO3(aq) a weak acid
Example: Sulfur trioxide + water
SO3(g) + H2O
2 H+(aq) + SO42-(aq) a strong acid
Example: Carbon dioxide + water
CO2(g) + H2O
H2CO3(aq) a weak acid
c. Metal oxide + nonmetal oxide
salt, with the nonmetal appearing in a radical where it has the same
oxidation state as in the oxide. Easily reversed at higher temperatures.
Example: Calcium oxide + sulfur trioxide
CaO(s) + SO3(s)
CaSO4(s)
Example: sodium oxide + carbon dioxide
Na2O(s) + CO2(s)
3. Metal oxides in acids
Na2CO3(s)
metal ion and water
Example: hydrochloric acid is added to solid calcium oxide
H+(aq) + CaO(s)
Ca2+(aq) + H2O
4. Some decomposition reactions - (the reverse of many composition reactions above) where a compound
decomposes into two or more simpler compounds or elements.
a. Metallic carbonates, when heated, form metallic oxides and carbon dioxide. They are readily reversed at
room temperature.
Example: Calcium carbonate solid, when heated
CaCO3(s)
CaO(s) + CO2(s)
b. Many metallic hydroxides, when heated, decompose into metallic oxides and water. Common exceptions
are NaOH and KOH, which will not decompose upon heating. They are readily reversed at room temperature.
Example: Calcium hydroxide solid, when heated
Ca(OH)2(s)
CaO(s) + HOH(g)
Example: Aluminum hydroxide solid, when heated
Al(OH)3(s)
Al2O3(s) + HOH(g)
c. Metallic chlorates, when heated, decompose into metallic chlorides and elemental oxygen.
Example: potassium chlorate solid, when heated
KClO3(s)
KCl(s) + O2(g)
d. Some acids, when heated, decompose into nonmetallic oxides and water. This includes carbonic acid,
sulfuric acid, and sulfurous acid. They are readily reversed at room temperature.
Example: carbonic acid, when heated
H2CO3(aq)
CO2(g) + HOH(l)
e. A few metallic oxides, when heated, decompose.
Example: mercury(II) oxide, when heated
HgO(s)
Hg(l) + O2(g)
Example: lead(IV) oxide, when heated
PbO2(s)
PbO(s) + O2(g)
f. Many compounds can be decomposed into their elemental
state by electricity (electrolysis).
Example: electrolysis of molten sodium chloride
NaCl(l)
Na(l) + Cl2(g)
Example: electrolysis of water
H2O(l)
H2(g) + O2(g)
5. Hydrolysis reactions:
Many salts react with water, particularly nonmetallic halides. If the water is written as HOH, combining the H
from the water with the more (or most) electronegative element from the other compound usually gives the
formula for one of the products. The other product contains the remaining elements. The formula for this
second compound usually needs to be rearranged in order to make clear its acidic properties.
Example: Phosphorus trichloride reacts with water
PCl3(g) + HOH
H+(aq) + Cl-(aq) + H3PO3(aq)
Example: Phosphorus(V) oxytrichloride is added to water
POCl + HOH
H+(aq) + Cl-(aq) + H3PO4(aq)
Group 1 & 2 metal oxides react with water to form bases:
MO + H2O
M ion(aq) + OH-(aq)
Example: Solid sodium oxide is added to water
Na2O(aq) + H2O
Na+(aq) + OH-(aq)
Nonmetal oxides react with water to form acids:
Example: Sulfur trioxide is bubbled into water.
SO3(g) + H2O
H+(aq) + SO42-(aq)
a. Reactions between an anion base and water. This involves the hydrolysis of the conjugate base of a weak
acid.
Example: Carbonate ion in water
CO32-(aq) + H2O
HCO3-(aq) + OH-(aq)
b. Reactions between a cation acid and water. This involves the hydrolysis of the conjugate acid of a weak
base.
Example: Zinc ion in water
Zn2+(aq) + H2O
ZnOH+(aq) + H+(aq)
6. LIGANDS - The reactions of coordination compounds and ions. The number of different reactions falling
into this category that are considered in general chemistry courses is rarely large. The ligands most frequently
considered, attached to a central atom that is usually a metal ion, are the ammonia molecule and the hydroxide
ion. Writing the reactions that form the coordination ions requires (1) recognition of the possibility of the
formation of the coordination ion when an ample source of the ligands is available and (2) knowledge of the
metal ions likely to form coordination ions, of what the common ligands are, and of the likely number of the
ligands to be attached to the central ion. It may be useful to keep in mind that the number of ligands attached to
a central metal ion is sometimes twice the oxidation number of the central metal: Ag(NH3)2+ , Zn(OH)42+.
The breakup of these coordination ions if frequently achieved by adding an acid. The products are the metal ion
and the species formed when hydrogen ions from the acid react with the ligand (NH4+ from NH3 and HOH
from OH-).
Al(OH)4-(aq) + H+(aq)
Al3+(aq) + H2O
Nearly all of the coordination ions are transition metal ions, such as Cu+, Cr3+, Fe3+, Zn2+, Pt2+, Ag+, Co3+,
Pd2+, etc.
The ligands vary in their ability to attach to coordination ions, with stronger ligands replacing weaker ligands
from the coordination ion:
CN- > NO2- > NH3 > H20 > OH- > F- > Cl- > Br- > Istrong ligand
weak ligand
Example: Aqueous diamine silver ion is mixed with cyanide ion solution
Ag(NH3)2 + 2CN-
Ag(CN)2- + 2 NH3
7. Reactions based on nonwater definitions of acids and bases. Both Bronsted and Lewis definitions of acids
and bases can be illustrated by the writing of equations. Recognizing that an acid and a base are the reactants
according to one of the definitions and knowing how they react is the best approach for the writing of such
reactions. Bronsted reactions involve the transfer of a proton while Lewis reactions involve the formation of a
coordinate covalent bond by donating a pair of electrons.
Examples: Ammonia gas reacts with hydrogen chloride gas
NH3 + HCl
NH4Cl Bronsted acid/base reaction
Boron trifluoride reacts with ammonia
BF3 + :NH3
B.
F3B:NH3 Lewis acid/base reaction
Reactions involving changes in oxidation states.
1. Some combination reactions:
Two elements react to form a binary compound of the two. The symbol for the more electropositive element is
written first and valence relations are used to obtain the formula.
Examples: Sodium metal reacts with chlorine gas
Na + Cl2
2NaCl
Aluminum metal reacts with oxygen gas
Al
+
O2
Al2O3
2. Reactions between an oxidizer and a reducer:
Products from such reactions can usually be predicted from knowledge about a limited number of oxidizers and
reducers. A useful list of such reagents follows:
Important oxidizers Formed in the reaction
MnO4- in acid solution . . . . . . . . . . . . . . . . . . Mn2+ + H2O
MnO2 in acid solution . . . . . . . . . . . . . . . . . . . Mn2+ + H2O
MnO4- in neutral. . . . . . . . . . . . . . . . . . . . . . MnO2
MnO4- in basic solution . . . . . . . . . . . . . . . . . MnO42Cr2O72- in acid solution. . . . . . . . . . . . . . . . . Cr3+ + H2O
HNO3 concentrated. . . . . . . . . . . . . . . . . . . . . NO2 + H2O
HNO3 dilute. . . . . . . . . . . . . . . . . . . . . . . . . . . NO + H2O
H2SO4 hot, concentrated . . . . . . . . . . . . . . . SO2 + H2O
metal-ic ions . . . . . . . . . . . . . . . . . . . . . . . . . metal-ous ions
free halogens . . . . . . . . . . . . . . . . . . . . . . . . . halide ions
Na2O2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . NaOH
HClO4 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ClH2O2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . O2 + H2O
S2- . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . S
halate ion . . . . . . . . . . . . . . . . . . . . . . . . . . . . halide ion
Important reducers Formed in the reaction
halide ions . . . . . . . . . . . . . . . . . . . . . . . . . . . free halogen
free metals . . . . . . . . . . . . . . . . . . . . . . . . . . . metal ions
SO32- ions (or SO2) . . . . . . . . . . . . . . . . . . . . SO42- ions
NO2- ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . NO3- ions
free halogen in dilute basic solution . . . . . . . . . hypohalite ions
free halogen in conc. basic solution . . . . . . . . . halate ions
metal-ous ions . . . . . . . . . . . . . . . . . . . . . . . . . metal-ic ions
C2O42- . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . CO2 + H2O
To predict products of a reaction that fits into this category a student looks at the reagents given in the question
to see if there are available both an oxidizer and a reducer. This step may involve recognizing the ions that are
parts of the compounds listed as the reagents. Then one can write the appropriate products from the oxidizer
and the reducer present. One keeps in mind the acid or the base present if an acid or a base is listed as a
reactant. In acidic solutions, any metal ions formed can combine with the anion of the acid to form salts.
Keeping in mind the solubilities of salts, one can then predict whether the products include a precipitated salt or
whether ions are the appropriate products. In basic solutions, any anions present can combine with the cation of
the base present to form salts. Again, the solubility considerations indicate whether formulas of salts or of ions
should be written as products.
Examples: Aqueous potassium dichromate and iron(II) in acid
solution
Cr2O72- + Fe2+ + H+
Cr3+ + Fe3+ +
H2O
Aqueous potassium permanganate and sodium chloride in acid solution
MnO4- + Cl- + H+
Mn2+ + Cl2 + H2O
A piece of iron is added to a solution of iron(III) sulfate
Fe +
Fe3+
Fe2+
Solid silver is added to a dilute nitric acid (6M) solution
Ag + NO3- + H+
Ag+ + NO + H2O
Hydrogen peroxide is added to an acidified solution of sodium bromide
H2O2 + Br-
Br2 + O2 + H2O
Copper(II) sulfide is oxidized by dilute nitric acid
S2- + NO3- + H+
S + NO + H2O
Sodium iodide is added to an acidified solution of potassium iodate
I- + IO3- + H+
I2 + H2O
3. Single Replacement Reactions
The more reactive element will replace the less reactive element from the compound. The products are the less
reactive element and a compound containing the more reactive element.
General reaction equation:
Metal + salt where A and B are metals and A is the more reactive metal. Metals above H can replace H from
acids while metals above Mg can replace H from water.
A + BX
Nonmetal + salt
X + AY
AX + B
where X and Y are nonmetals and X is the more reactive nonmetal.
AX + Y
Use the Activity Series to predict the products. Any element higher in the Activity Series will replace any
element, lower in the series, from a compound.
Activity Series
Metal Halogens
Li M F2
K o Cl2
Ca s Br2
Mg t I2
Al
Zn t
Cr o
Fe
Ni L
Sn e
Pb a
Hs
Cu t
Hg
Ag
Au
Examples:
Lead foil is immersed in silver nitrate solution.
Pb + Ag+
Pb2+ + Ag
Chlorine gas is bubbled into a solution of potassium iodide
Cl2 + I-
Cl- +
I2
Zinc foil is placed into a solution of tin(II) chloride
Zn + Sn2+
Zn2+ + Sn
A strip of magnesium is added to a solution of silver
nitrate
Mg + Ag+
Mg2+ + Ag
Liquid bromine is added to a solution of sodium iodide
Br2 + I-
Br-
+
I2
ORGANIC CHEMISTRY - the basics
A significant numbers of questions on the AP Chemistry Exam involve organic reactions. We will cover only
the basics of organic chemistry; just enough to successfully answer organic reaction questions.
NOMENCLATURE (names of compounds)
Prefix Carbons Example Formula
meth (methyl) 1 carbon methane CH4
eth (ethyl) 2 carbons ethane C2H6
prop (propyl) 3 carbons propane C3H8
but (butyl) 4 carbons butane C4H10
Alkanes, alkenes and alkynes:
C to C General
Suffix Bond Formula Example Formula
-ane single CnH(2n+2) ethane C2H6
-ene double CnH2n ethene C2H4
-yne triple CnH(2n-2) ethyne C2H2
Common functional Groups:
Functional General
Class Group Formula Example Formula
Alcohols -OH R-OH methanol CH3OH
Ketones -C- R-C-RÆ propanone CH3CCH3
O
O (acetone)
O
Carboxylic -COOH R-COOH ethanoic acid CHCOOH
acids
Esters -COO- R-COO-RÆ ethyl acetate
CH3COOCH2CH3
Reactions:
Burn a hydrocarbon in air:
Example: butane burns in air
produces CO2 and H2O
C4H10 + O2
CO2 + H2O
Halogen substitution reactions with alkanes: (X replaces H)
CnH(2n+2) + X2
CnH(2n+1)X + HX
Ethane reacts with chlorine gas
C2H6 + Cl2
C2H5Cl + HCl
Addition reactions with alkenes and alkynes:
Alkenes react with hydrogen gas to form alkanes:
CnH2n + H2
CnH(2n+2)
Propene reacts with hydrogen gas to form propane
C3H6 + H2
C3H8
Alkynes react with hydrogen gas to form alkenes or alkanes:
CnH(2n-2) + H2
CnH2n or CnH(2n+2)
Propyne reacts with excess hydrogen gas
C3H4 + 2 H2
C3H8
Equal moles of propyne and hydrogen gas are mixed
C3H4 + H2
C3H6
Alkenes react with halogens:
CnH2n + X2
CnH2nX2
Propene reacts with bromine gas to form propane
CnH2n + X2
CnH2nX2
C3H6 + Br2
C3H6Br2
Alkenes react with HX
CnH2n + HX
CnH(2n+1)X
Ethene reacts with hydrogen iodide gas
C2H4 + HI
C2H5I
Alkenes react with water vapor to form alcohols
CnH2n + HOH
CnH(2n+1)OH
Ethene reacts with water vapor to form ethanol
C2H4 + HOH
C2H5OH
Alkynes react with halogens:
CnH(2n-2) + X2
CnH(2n-2)X2 or CnH(2n-2)X4
Propyne reacts with excess fluorine gas
C3H6 + 2 F2
C3H6F4
Equal moles of propyne and fluorine gas are mixed
C3H6 + F2
C3H6F2
Special case: Benzene, C6H6, substitutes a hydrogen atom with another atom or group.
C6H6 + Cl2
C6H5Cl + HCl
Carboxylic acids react with alcohols to form esters and water, similar to an acid-base reaction.
R-COOH + RÆ-OH
R-COO-RÆ + H2O
Ethanoic acid (acetic acid) reacts with methanol (methyl alcohol)
CH3COOH + CH3OH
CH3COOCH3 + H2O
How to Get the Highest Possible Score on Question 4
Write the reactants in symbol form for all eight reactions, showing each reactant in net ionic form as follows:
Strong acids, strong bases and soluble salts are written as ions.
Weak acids, weak bases, and insoluble salts are written as molecules.
Classify the reactions as:
(A/B)
Acid/Base - Look for H or OH or salts which could act as weak acids or weak bases.
Precipitation - Look for insoluble salts as products. Learn the solubility rules!
(R) Redox - Look for substances that change oxidation states. Look
for common oxidizing agents (NO3-, MnO4-, Cr2O72-,
H2O2) and reducing agents (Cl-, Br-,I*- and elemental
metals). If itÆs an element reacting with a compound,
itÆs redox.
(H) Hydrolysis û Look for a metal oxide, metal hydrides, or nonmetal
oxide in water.
Decomposition û If a single compound is heated it must
decompose. Look for chlorates, carbonates, sulfates,
sulfites and ammonium compounds. Ammonium
hydroxide and carbonic acid decompose in water. DonÆt
forget reverse decomposition where the products of
decomposition form a larger compound.
Organic û Look for hydrocarbon compounds (CH).
Ligands û Look for complexes of transition metal ions and other
heavy metal ions. Typical ligands are OH-, CN-, NH3.
Miscellaneous û Those reactions that donÆt fit the seven
categories listed above.
3. You only answer five of the eight choices, so eliminate the three you have the least confidence in
completing. Try to predict the products of the remaining five choices. Get all the reactants correct and two or
three products correct and youÆve got a FANTASTIC score!
David T. Lee, chemistry teacher for 35 years (retired).
Mountain Lakes High School, Mountain Lakes, NJ 07046
Email address * HYPERLINK "mailto:dtleemlhs@aol.com" **dtleemlhs@aol.com*
COMMON POSITIVE IONS
Name Symbol
Name Symbol
aluminum ion Al3+ mercury(II) or mercuric ion Hg2+
ammonium ion NH4+ mercury(I) or mercurous ion Hg22+
barium ion Ba2+ nickel(II) or nickelous ion Ni2+
bismuth ion Bi3+ potassium ion K+
calcium ion Ca2+ rubidium ion Rb+
cadmium ion Cd2+ silver ion Ag+
cesium ion Cs+ strontium ion Sr2+
chromium(III) or chromic ion Cr3+ tin(IV) or stannic ion Sn4+
chromium(II) or chromous ion Cr2+ tin(II) or stannous ion Sn2+
cobalt(III) or cobaltic ion Co3+ zinc ion Zn2+
cobalt(II) or cobaltous ion Co2+
copper(II) or cupric ion Cu2+
copper(I) or cuprous ion Cu+
hydrogen ion H+
iron(III) or ferric ion Fe3+
iron(II) or ferrous ion Fe2+
lead(II) or plumbous ion Pb2+
lithium ion Li+
magnesium ion Mg2+
COMMON NEGATIVE IONS
Name Symbol
Name Symbol
acetate ion C2H3O2- hydrogen sulfite or bisulfite ion HSO3bromate ion BrO3- hydroxide ion OHbromide ion Br- hypochlorite ion ClOcarbonate ion CO32- iodate ion IO3chlorate ion ClO3- iodide ion Ichlorite ion ClO2- monohydrogen phosphate ion HPO42chloride ion Cl- nitrate ion NO3chromate ion CrO42- nitrite ion NO2cyanide ion CN- oxalate ion C2O42dichromate ion Cr2O72- oxide ion O2dihydrogen phosphate ion H2PO4- perchlorate ion ClO4fluoride ion F- periodate ion IO4hydride ion H- permanganate ion MnO4hydrogen carbonate or HCO3- peroxide ion O22-
bicarbonate ion
hydrogen oxalate or HC2O4- phosphate ion PO43bioxalate ion
hydrogen sulfate or HSO4- sulfate ion SO42bisulfate ion
hydrogen sulfide or HS- sulfide ion S2bisulfide
sulfite ion SO32thiocyanate ion SCN-
Predicting Products of Reactions
AP Chemistry