Chapter 4
Chemical Formulas, Reactions, Redox and Solutions
Terms to Know:
Solubility
the amount of substance that dissolves in a given volume of solvent at a given
temperature.
Solute
a substance dissolved in a liquid to form a solution
Solvent
the dissolving medium in a solution
Solution
a homogeneous mixture of a solute in a solvent
Molarity, M
Moles of solute per volume of solution in Liters (Expression of concentration )
M
mole of solute
Liter of solution
Conductivity
the ability to conduct electricity in an aqueous solution
Electrolyte
Ionic compounds that conduct electricity in H2O
Strong- excellent conductor, fully ionized
Weak- poor conductor, partially ionized
Nonelectrolyte
non-conductors, no ions present to conduct
Precipitation Reaction
a reaction in which an insoluble substance forms and separates from the solution
Spectator ions
ions present in solution that do not participate directly in a reaction. Appear as both
reacts and products.
Reducing agent
electron donor; a reactant that donates electrons to another substance to reduce the
oxidation state of one of its atoms
Oxidizing agent
electron acceptor; a reactant that accepts electrons from another reactant
Reduction
a gain of electrons (a decrease in oxidation state)
Oxidation
a loss of electrons (an increase in oxidation state)
Redox reaction
a reaction where electrons are gained (reduction) or lost (oxidation)
Oxidation state (number) -the apparent charge of an atom
-a concept that provides a way to keep track of electrons in oxidation-reduction
reactions according to certain rules
Net-ionic equation
an equation for a reaction in a solution, where strong electrolytes are written as ions,
showing only those components that are directly involved in the chemical change.
Ionic equation
an equation for a reaction in solution, where strong electrolytes are written as ions
Concepts
I. Assigning Oxidation Numbers
Rules to calculate the oxidation number of any atom in any molecule.
1. In molecules group 1 and group 2 metals are +1 and +2 respectively.
2. Fluoride is always –1.
3. O is always -2 except in peroxides when it is -1 and OF2 where it is +2.
4. H is always +1 with nonmetals and –1 with metals.
5. The sum of all the oxidation number of all the atoms equals the overall charge
N2
N =0
Na1+
Na = +1
NH41+
H = +1
N = -3
SO32O = -2
S = +4
Chapter 4
Chemical Formulas, Reactions, Redox and Solutions
II. Oxidation – Reduction (Redox) (Single replacement reactions are redox rxns.)

Involve a transfer of electron(s)
o Oxidation- loss of electrons; the atom that is oxidized is the reducing agent
o Reduction- gain of electrons; the atom that is reduced is the oxidizing agent
Example: In the following rxn, identify the oxidized atom, reduced atom, oxidizing and reduction agents.
Fe3+ + Cu1+ → Cu2+ + Fe2+
Solution: First determine what is occurring with each atom. Then apply the definitions.
Fe: +3 → +2
Fe3+ + 1e- → Fe2+
reduction, oxidation agent
Cu: +1 → +2
Cu1+
→ Cu2+ + 1eoxidation, reduction agent
Note: In half-rxns, when reduction occurs, e’s are reactants and when oxidation occurs, e’s are products.
III. Balancing Redox Rxns
{Note: Any time a question says a rxn occurs under certain conditions, it is most likely a redox & ½ reaction prob.}
There are two methods that both arrive at the same answer. Your choice, choose the method that works for you.
Method 1
Method 2
Under acidic conditions
1st Write the two half reactions.
Then for each ½ reaction:
i
Balance the mass
1–Balance non- H and O atoms.
2–Add H2O to balance O’s
3–Add H+’s to balance H’s
ii Balance the charge by adding e-’s.
nd
2 Cross multiply to cancel e-’s.
3rd Add the two half-rxns to recreate the full rxn.
Under acidic conditions
1st Write the two half reactions.
Then for each ½ reaction:
Balance non–H and O atoms.
Add e’s to the appropriate side to match the
change in the oxidation number.
2nd
Cross multiply to cancel e-’s.
rd
3
Add the two half-rxns to recreate the full rxn.
4th
Add H2O to balance O’s
th
5
Add H+’s to balance H’s
Under basic conditions
1st Balance ½ reactions as if acidic.
2nd Add OH–’s to each side to cancel the H+’s
OH– + H+→ H O. Cancel extra H2O’s
Under basic conditions
1st Balance ½ reactions as if acidic.
2nd Add OH1-’s to each side to cancel the H+’s
OH– + H+→ H O. Cancel extra H2O’s.
Example: Balance the following reaction:
Fe2+ + MnO41- → Mn2+
+ Fe3+
Under Acidic Conditions:
1st Fe2+
→ Fe3+ + 1e+
8H + MnO41- + 5e- → Mn2+ + 4H2O
Example: Balance the following reaction:
Fe2+ + MnO41- → Mn2+
+ Fe3+
Under Acidic Conditions:
1st Fe2+
→ Fe3+ + 1e[Fe: +2 to +3]
1MnO4 + 5e→ Mn2+ [Mn: +7 to +2]
2nd 5( Fe2+
→ Fe3+ + 1e- )
1(8H+ + MnO41- + 5e- → Mn2+ + 4H2O)
3rd 5Fe2+ + 8H+ + MnO41- → 5Fe3+ +Mn2+ + 4H2O
2nd 5( Fe2+
→ Fe3+ + 1e- )
1(8H+ + MnO41- + 5e- → Mn2+ + 4H2O)
3rd 5Fe2+ + MnO41- → 5Fe3+ +Mn2+
4th 5Fe2+ + MnO41- → 5Fe3+ +Mn2+ + 4H2O
5th 5Fe2+ + 8H+ + MnO41- → 5Fe3+ +Mn2+ + 4H2O
Under basic conditions:
acidic cond’s
neuralize H+
basic cond’s
5Fe2+ + 8H+ + MnO41- → 5Fe3+ +Mn2+ + 4H2O
8OH + 5Fe2+ + 8H+ + MnO41- → 5Fe3+ +Mn2+ + 4H2O + 8OH8H2O + 5Fe2+
+ MnO41- → 5Fe3+ +Mn2+ + 4H2O + 8OH2+
4H2O + 5Fe
+ MnO41- → 5Fe3+ +Mn2+ + 8OH-
Chapter 4
Chemical Formulas, Reactions, Redox and Solutions
IV. Ionic Solutions
1.
Soluble ionic compounds dissociate when dissolved in water and are always written as ions.
ex. AbBa(s)  Aa+(aq) + Bb- (aq)
2.
Soluble ionic compounds “redissociate” when formed in water leaving no net reaction because all ions cancel.
ex. Aa+(aq) + Bb- (aq)  Aa+(aq) + Bb-(aq)
3.
Insoluble ionic compounds precipitate when formed in H 2O and do not dissolve in H2O.
ex. Aa+(aq) + Bb-(aq)  AbBa(s)
V. Ionic Reactions
When solutions of ionic compounds are mixed, a chemical reaction may or may not occur. A reaction occurs if
anything remains after cancelling the spectator ions. You will be asked to predict if a chemical rxn occurs and
if it does to write the balanced chemical equation for the rxn.
Things you should be aware of…
 These reactions involve 2 solid ionic compounds dissolved in H2O, Aa+Bb- + Cc+Dd
When two opposite ions attract they will form either a soluble or insoluble compound (apply the rules).
–A soluble compound re-dissociates into ions so soluble compounds are written as ions
–An insoluble compound precipitates and is written as an empirical formula.

Ions that appear as both reactants and products are spectator ions and are removed from the ionic
chemical equation.

Once the spectator ions are removed, we write the net ionic equation.
Example: A solution of MgCl2 (aq) is mixed with a solution of AgNO3. Predict the reaction that occurs.
To solve:
Solution:
1st in the question, “connect” cation to anion to predict the potential products
2nd apply solubility rules – soluble are ions; and insoluble are formulas
3rd write the equation cancelling the spectator ions.
Potential products are MgNO3 (soluble) and AgCl (insoluble).
Mg2+ + Cl1– + Ag1+ + NO31– → Mg2+ + NO31– + AgCl
So, Ag1+ + Cl1- → AgCl
VI. Conducting
Electrolytes – ionic compounds that conduct electricity in H2O or ionized covalent compounds.
 Strong- excellent conductor, typically because they are 100% ionized
 Weak- poor conductors, typically because they are partially ionized
Non-Electrolytes – nonconductors, cause no ions present to conduct, typically covalent compounds.
Example: Which Conduct More?
1. Sugar vs. NaCl
2. Acetic acid vs. NaCl
3. 1 M NaCl vs. 0.5 M NaCl
4. 1 M BaCl2 vs. 1 M NaCl
(1M has more ions)
(1M with 3 ions vs 1M with 2 ions)
Chapter 4
Chemical Formulas, Reactions, Redox and Solutions
VII. Review of Formula Writing and Ionic Equation Writing
TABLE–1
IONS → FORMULAS
To write Neutral Binary Formulas from Ions
1
Write the two ions side-by-side, cation first,
without the charges.
Example
Aa+ + Bb- → ________
1
Aa+ + Bb- → A B
2
2
Aa+ + Bb- →
3
Aa+ + BOxb- →
3
Determine how many of each ion are needed to
create a neutral formula, write that number as a
subscript. [Typically this is done by criss-crossing
and reducing the charges.]
Note that if a polyatomic ion is used, the subscript
is written outside of ( )’s
TABLE–2
FORMULAS → IONS
Dissociating Neutral Binary Formulas Into Ions
note that you do not have to write the 1’s, but it helps
___ + ___ or
Ab(BOx)a → __ + __
Ab(BOx)a →
1
Split the formula into two “halves”. Often this is 1
done by writing the first symbol as the first half and
the remaining symbols as the second half. Do NOT
write any subscripts unless they are inside ( )’s or
follow two capital letters (see example)
AbBa →
A + B
2
The first ion is positive the second is negative.
AbBa →
A
3
Now determine the magnitude of the charge. Do
the opposite of “Ions to formulas” by criss-crossing
the subscripts into the position of charges.
3/4 AbBa →
4
Ab(BOx)a
Al3+ + OH1- → Al1(OH)3
Example
AbBa →
2
AbBa
+
+ B
or
-
Aa+ + Bb-
A + BOx
+
A
or Ab(BOx)a →
Aa+ + BOxb-
+ BOx
To make sure you are right (and often you are not)
you need to check your charges with the periodic 4 Examples of Checking Charges
table or with the charges on the polyatomic ions youCu2O → Cu1+ + O2were to memorize.
Note the charge for O matches the periodic table. So we assume
the charges are correct.
Trend For Charges
For simple compounds, any element within the vertical
family headed by the following s and p block elements
will have the corresponding charge.
H1+, Be2+, B3+, C4+/4-, N3-, O2-, F1-, He no charge
K2SO4 →
-
or Ab(BOx)a →
K1+ + SO42-
Note the charge for K and SO4 matches the periodic table and the
memorized charge, so we assume the charges are correct.
Also note, the subscript “4” was not criss-crossed. The invisible
“1” outside the invisible “( )’s” was criss-crossed. The “4” follows
two capital letters and stays put.
For more complex compounds (AP Chem), the rules for CuO → Cu1+ + O1determining oxidation numbers must be followed.
Note that the charge for O should be a 2- from the periodic
.
table. So, we must double all charges to get it from 1- to 2-.
So...
CuO → Cu2+ + O2TABLE–3
NAMES → FORMULAS
To write Neutral Binary Formulas from Names
Example: Write the formula of sodium sulfate
sodium = Na1+
1 Write the symbols for the corresponding cation and anion 1,
sulfate = SO422 Write the formula using Table–1.
2
Na1+ + SO42- → Na2SO4
Writing, Dissociating and Naming Ionic Formulas
TABLE–4
NAMING IONS and BINARY IONIC COMPOUNDS
Naming Binary Ionic Compounds
Example
You never name formulas. Name ONLY ions! So…
If you are given a formula to name you must
first dissociate it into ions and then name the ions.
or
If you are given a name and asked to write a
formula, you must first write the ions and then
write the formula.
Naming Cations
if s- or p-block elements
The name is the elemental name.
if d-block elements
The name is the elemental name followed by the
charge as a Roman numeral in parentheses.
Naming Cations
if s- or p-block elements
Li1+
lithium.
Al3+
aluminum
if d-block elements
Ni1+
nickel (I).
Ni3+
nickel (III)
Naming Anions
if mono-atomic (only one elemental symbol)
Then the name is the elemental name, but with the
ending changed to –ide.
Naming Anions
if monatomic:
F1fluoride (not fluorine)
O2oxide
(not oxygen)
As3- arsenide (not arsenic)
if polyatomic (more than one elemental symbol)
Then the name is the memorized name.
if polyatomic: These are the names to memorize
C2H3O21- acetate
CO32carbonate
1CN
cyanide
OH1hydroxide
NO31nitrate
1NO2
nitrite
SO42sulfate
SO32sulfite
3PO4
phosphate
PO33phosphite
Note that there exists one polyatomic cation that should
be memorized.
NH41+ is ammonium
All other cations will be one elemental symbol.
To name a formula
1
2
Dissociate into ions, then
Name the cation followed by the anion.
To name a formula
ex. Write the name of CuCO3
1
CuCO3 → Cu2+ + CO322
Cu2+ = copper (II)
CO32- = carbonate
so, the name is copper (II) carbonate
VIII. Solubility Rules
The following rules must be memorized. These rules are applied to a solution of an ionic compound and tell you whether
that compound is soluble and so must be written as ions. OR If the compound is insoluble and so must be written as an
empirical formula.
Aside from the very first rule, the rules are applied by looking at the anion and then checking for exceptions.
Example:
What are the solubilities of each compound?
Solution:
MgNO3
AgCl
K2SO4
K2C2O4
Soluble
Insoluble
Soluble
Soluble
All nitrates are soluble, no exceptions so you would write Mg2+ + NO31Most chlorides are soluble, this is an exception, so you would write AgCl
Most sulfates are soluble or first rule, so you would write K1+ + SO42Most oxalates are insoluble this is an exception or first rule, so write K1+ + C2O41-
Solubility Rules Must Be Memorized!
TABLE–1 SOLUBLE ANIONS
Soluble Compounds
Exceptions
FIRST RULE Most salts containing alkali metal ions and
the ammonium ion are soluble.
Salts of nitrate, NO31chlorate, ClO31perchlorate, ClO41acetate, CH3CO21Most salts of Cl1-, Br1-, and I1-
Halides of Ag1+, Hg22+, and Pb2+
Compound containing fluoride, F1-
Fluorides of Mg2+, Ca2+, Sr2+, Ba2+, and Pb2+
Salts of sulfate, SO42-
Sulfates of Ca2+, Sr2+, Ba2+, and Pb2+
TABLE–2 INSOLUBLE ANIONS
Insoluble Compounds
All salts of
carbonate, CO32-
Exceptions
Salts of NH41+ and the alkali metal cations.
phosphate, PO43oxalate, C2O42chromate, CrO42Most metal sulfides, S2-
Salts of NH41+ and the alkali metal cations.
most metal hydroxides, OH1-, and oxides, O2-
Salts of NH41+ and the alkali metal cations.