.Guessing Reaction Products
This handout presents some guidelines on how to predict reactions. The products one gets by
using these rules are those normally expected. Exceptions to these are the difficult part of
chemistry, and one spends a much more time on learning the exceptions than those expected.
In order to understand and utilize this handout, you must be able to:
1) determine oxidation numbers
2) identify, recall and write the common polyatomic ions and
3) construct Lewis dot structures.
If you are weak in these areas, review until you are competent.
There are two general classes of reactions, redox (oxidation-reduction) and non-redox. Redox
reactions are those where there is a change in oxidation number during the reaction. The
following are special definitions used here.
"expected" and "normal" - These are those products of redox reactions that one would obtain by the rules presented here.
"acceptable Lewis dot structure" - These are those structure which are possible by the
rules given in the Laboratory Manual
.
WARNING - Never balance a reaction before you are certain of the reaction products.
Never distort the answer for the product merely to balance a reaction. Any reaction can be
balanced, regardless of the complexity of the reactants or products.
Non-redox Reactions
I. Production of bases or acids by reaction with water
A. Reaction of water with normal metal oxides:
Examples:
normal oxide + water ⇌ base
Li2O + H2O ⇌ 2LiOH
Na2O + H2O ⇌ 2NaOH
CaO + H2O ⇌ Ca(OH)2
MgO + H2O ⇌ Mg(OH)2
B. Reaction water with non-metal oxides for which there is an acceptable Lewis dot structure:
Examples:
SO3 + H2O ⇌ H2SO4
SO2 + H2O ⇌ H2SO3
P4O6 + 6H2O ⇌ 4H3PO3
P4O10 + 6H2O ⇌ 4H3PO4
N2O5 + 2H2O ⇌ 2HNO3
N2O3 + 2H2O ⇌ 2HNO2
C. Exceptions involving redox violate the caveats given, but yield bases or acids with the possibility of changes in oxidation number:
Na2O2 + H2O ➞ NaOH + H2O2 - a reaction of a peroxide (not normal) to give hydrogen peroxide.
4KO2 + H2O ➞ 4KOH + O2 - a reaction of a superoxide (not normal) to give oxygen gas
3NO2 + H2O ➞ 2HNO3 + NO - a reaction of an oxide that does not have an acceptable Lewis dot structure. I. E. NO2 has a Lewis dot structure with an odd number of electrons or one unpaired electron.
II. Acid-base reactions, by any of the definitions: Arrhenius, Brønsted-Lowry or Lewis.
Examples:
A. Arrhenius:
Acid + Base ➞ Salt + (water optional)
HCl + NaOH ➞ NaCl + H2O
HCl + NH3 ➞ NH4Cl
H2SO3 + KOH ➞ KHSO3 + H2O
H2SO3 + 2KOH ➞ K2SO3 + 2H2O
B. Brønsted-Lowry:
Acid 1 + Base 2 ➞Acid 2 + Base 1
HCl + H2O ➞ Cl– + H3O+
NH3 + H2O ➞ NH4+ + OH–
CH3COOH + H2O ➞ CH3COO– + H3O+
NH4 + H2O ➞ NH3 + H3O+
CN– + H2O ➞ HCN + OH–
H2O + H2O ➞ H3O+ + OH–
NH3(l) + NH3(l) ➞ NH4+ + NH2–
HCO3– + H2O ➞ H2CO3 + OH–
C. Lewis:
Acid + Base ➞ compound
or Acid + Base ➞ complex ionBF3 + NF3 ➞ BF3NF3
Ag+ + 2NH3 ➞ Ag(NH3)2+
Cu2+ + 2NH3 ➞ Cu(NH3)22+
Zn2+ + 4NH3 ➞ Zn(NH3)42+
Zn2+ + 4OH– ➞ Zn(OH)42–
III. Reactions of hydroxides or normal oxides with CO2 to yield carbonates, CO32– or HCO3– and dissolution of CO2 in H2O and reaction leading to HCO3–
CaO + CO2 ⇌ CaCO3
Ca(OH)2 + CO2 ⇌ CaCO3 + H2O
Na2O + CO2 ⇌ Na2CO3
2KOH + CO2 ⇌ K2CO3 + H2O
CO2 + H2O ⇌ H2CO3
H2CO3 + OH– ⇌ HCO3– + H2O (see above)
Redox Reactions
IV. Reactions of Metals with non-metals
The oxidation number for metals will become the principal oxidation number.
The oxidation number for non-metals (including hydrogen) will become the principal oxidation
number. (Recall that hydrogen as a non-metal is –1)
Examples:
2Na + F2 ➞ 2NaF
Ca + F2 ➞ CaF2
16Na + S8 ➞ 8Na2S
2Ca + O2 ➞ 2CaO
2Mg + O2 ➞ 2MgO
4Y + 3O2 ➞ 2Y2O3
Sc + Cl2 ➞ ScCl3
4Al + 3O2 ➞ 2Al2O3
3Mg + N2 ➞ Mg3N2
2Na + H2 ➞ 2NaH
2Li + H2 ➞ 2LiH
Ca + H2 ➞ CaH2
2Sc + 3H2 ➞ 2ScH3
V. Reactions of hydrogen with non-metals
The oxidation number for hydrogen will become +1.
The oxidation number for non-metals will become the principal oxidation number.
Examples:
H2 + Cl2 ➞ 2HCl
2H2 + O2 ➞ 2H2O
3H2 + N2 ➞ 2NH3
12H2 + P8 ➞ 8PH3
16H2 + S8 ➞ 8H2S
H2 + Br2 ➞ 2HBr
VI. Reactions of non-metals with non-metals
The oxidation number for the most electronegative non-metal will become the principal oxidation number.
The oxidation number for the other non-metal will become either The IUPAC group number - 10 and/or 8 – the group number (which is –[principal oxidation number]).
The caveat is that a Lewis dot structure must be possible.
Example:
Cl2 + F2 ➞ 2ClF Cl has an oxidation number of (8 - group number) = 1
Cl2 + 7F2 ➞ 2ClF7 Cl has an oxidation number of the (group number – 10) - however, an allowed Lewis dot structure cannot be created. Therefore ClF7 does is not formed.
Examples: (Note that O is more electronegative than Cl.)
I2 + Cl2 ➞ 2ICl
S8 + 12F2 ➞ 8SF6
O2 + 2F2 ➞ 2OF2
S8 + 12O2 ➞ 8SO3
N2 + 3F2 ➞ 2NF3
P4 + O2 ➞ P4O6
P4 + O2 ➞ P4O10
C + 2F2 ➞ CF4
C + O2 ➞ CO2
Cl2 + 7O2 ➞ 2Cl2O7
Cl2 +O2 ➞ 2ClO
2N2 + 3O2 ➞ 2N2O3
2N2 + 5O2 ➞ 2N2O5
VII. Some important exceptions:
Note, there exists other possible compounds. This is especially true with oxygen. You must learn these compounds. Amongst these exceptions are:
Compounds of N and O which form:
N2O
NO2 with an unpaired electron (called a free radical)
NO with an unpaired electron (called a free radical)
N2O4 the dimer of NO2
(The expected compounds N2O3 and N2O5 also form.)
Compounds of S and O:
SO2 which has the same Lewis dot structure as ozone O3 does form.
SO is not a common oxide of sulfur, i.e. expected but not formed.
(SO3 which is predicted does form.)
Compounds of peroxides and superoxides:
In dry oxygen the following are common products of the reaction between O2 and the metal given. In wet conditions, the normal oxide (expected oxide) forms.
Na + O2 ➞ Na2O2 - sodium peroxide
K + O2 ➞ KO2 - potassium superoxide (an ion with an unpaired electron)
Rb + O2 ➞ RbO2 - rubidium superoxide - same for the rest of group 1
Ca + O2 ➞ CaO2 - calcium peroxide - same for metals of group 2 below calcium