AP Chemistry Note Outline
Chapter 4: Reactions
Tuesday, September 24 - TEST
Strong and Weak Electrolytes
Formation of a Gas
Types of Reactions Introduction
Precipitation Reactions
To be able to predict the
Acid-Base Reactions
Oxides
Hydrolysis
Oxidation-Reduction Reactions
Reaction Prediction--Overall
products of a chemical reaction
and answer a simple question
about it.
Strong and Weak Electrolytes (4.2)
Definitions: o Strong Electrolyte o Weak Electrolyte o Nonelectrolyte
Page 1 of 13

Look for the following:
1.
Two Ionic Compounds
This will be a double replacement reaction and either a ppt is formed or a gas. You MUST know your solubility rules.
2.
Oxides : know the oxide rules*
3.
Hydrolysis of a salt.
Water and a salt produces an acid/base combo.
4.
Acid/Base Reactionswater may be formed.
5.
Decomposition Reactions
6.
Words like Equilmolar are most likely acid/base reactions
7.
Words like Protic or
Diprotic lead you to acids
*Oxide Rules :
1.
nonmetal oxide and water → acid
2.
metal oxide and water → base
Page 2 of 13
Look for the following:
1.
Key words Acidified or basic solution
2.
Know basic oxidizing and reducing agents like

MnO
4

Cr
2
-
O
7
2-

H
2

SO
O
2
3
2-
3.
Combustion Reactions :
Reactions that involve air are combustion and the reactant is oxygen (O
2
).
4.
Single Replacement
Reactions: A pure metal
(NOT an ion) present is a single replacement reaction
5.
Any element like a halogen (Cl
2
, Br
2
) is S.R.
6.
Something is electrolyzed.

3.
nonmetal oxide + metal oxide → salt
Page 3 of 13

A.
Precipitation OR the formation of a ____________ _________________________ (4.4-6)
1.
Definition - the reaction of two ionic compounds to form an insoluble solid
2.
Writing Equations a.
Molecular Equation
Cd(NO
3
)
2
(aq) + Na
2
S (aq)
 b. Ionic Equations: All strong electrolytes are shown ionized
Cd 2+ (aq) + 2 NO
3
( aq) + 2 Na + ( aq) + S 2(aq)
 c. Net Ionic Equations:
Cd 2+ (aq) + S 2 -
(aq)

3. When a Weak Electrolyte is a Product
Molecular Equation
NaC
2
H
3
O
2
(aq) + HCl (aq)

Ionic Equation
Na + (aq) + C
2
H
3
O
2
(aq) + H + (aq) + Cl (aq)

Net Ionic Equation
C
2
H
3
O
2
(aq) + H + (aq)

4. What to do when a nonsoluble reactant is used:
Molecular Equation
Mg(OH)
2
(s) + 2 HCl (aq)

Ionic Equation
Net Ionic Equation
Page 4 of 13

YOU MUST LEARN THE SOLUBILITY RULES: THEY MUST BE MEMORIZED!!!
5 . Examples : a.
Solutions of strontium nitrate and sodium sulfate are mixed. b.
A solution of copper (II) chloride is added to a solution of sodium sulfide . c.
Solutions of sodium iodide and lead (II) nitrate are mixed d.
A solution of copper (II) sulfate is added to a solution of barium hydroxide . e.
Solutions of tri-potassium phosphate and zinc nitra te are mixed.
B. Formation of a Gas
1. There are four types of gases that can be formed. a. b. c. d.
2. Writing equations: a. Molecular:
2 HCl(aq) + Na
2
S(aq)
 b. Ionic c. Net Ionic
Page 5 of 13


3. Examples: a. Solid calcium carbonate is placed in hydrochloric acid. b. Solutions of sodium hydroxide and ammonium nitrate are mixed. c. Solid sodium sulfite is added to hydrobromic acid.
C. Acid – Base Reactions (Neutralization) (4.8)
1. Acid and a Base forms _____________ and a _______________.
a. Hydrochloric acid reacts with sodium hydroxide
2. Anhydrides: a. Acid anhydride
SO
3
(g) + H
2
O

N
2
O
5
(g) + H
2
O
 w / H
2
O
 b. Base anhydride
Na
2
O (s) + H
2
O

CaO (s) + H
2
O
 d. Anydrides with an acid or base produce ___________ and a ___________.
SO
2
(g) + NaOH (aq)
BaO (s) + HCl (aq)

 c.
Metallic oxides + nonmetallic oxides

Salt
CaO(s) + SO
2
(g)

Page 6 of 13

D.
Hydrolysis – salts reacting with _______________________.
Example : KF(s) + H
2
O

E. Some Decomposition Reactions
A. Base

metal oxide + water
Ca(OH)
2

CaO + H
2
O
B. Salt containing oxygen
 metal oxide + nonmetal oxide
CaCO
3

CaO + CO
2
C. Acid containing oxygen
H
2
CO
3

H
F. Non-Redox Rxn Practice
2

O + CO
2
water + nonmetal oxide
1.
Hydrogen sulfide is bubbled through a solution of silver nitrate. w / H
2
O
2.
Equal volumes of dilute equimolar solutions of sodium carbonate and hydrochloric acid are mixed.
3.
Dilute acetic acid solution is added to solid magnesium carbonate.
4.
Sulfur trioxide gas is added to excess water.
5.
Powdered magnesium oxide is added to a container of carbon dioxide gas.
6. Solid sodium acetate is added to water.
Page 7 of 13

II. Oxidation-Reduction Reactions
Oxidation Numbers
Rules
1.
Free elements have oxidation states of 0
2.
Ions keep their charges
3.
Oxygen in a compound is –2 unless as peroxide (then it is –1)
4.
Fluorine is –1
5.
Hydrogen is +1 unless as a hydride (then it is –1)
6.
Sum of oxidation states equals charge of substance. If it is a compound than it is zero.
Oxidation State Practice
Definitions :
Oxidation -
Reduction -
Oxidizing Agent -
Reducing Agent -
Examples:
2Pb + 3O
2

2PbO + 2SO
2
Cl
2
+ OH -

ClO + Cl + H
2
O
Page 8 of 13

Balancing Redox Reactions
Rules—Acid solution
1.
Write reactions as ½ reactions—One for the oxidation and one for the reduction
2.
Balance all elements except H and O
3.
Balance H with H +
4.
Balance O with H
2
O
5.
Balance Charge with e
-
6.
Multiply reactions by factors such that the e
-
7.
Add both ½ reactions
cancel
Examples
MnO
4
+_ Fe 2+

Fe 3+ + Mn 2+ (acid)
SO
4
2 + Cl -

Cl
2
+ SO
3
2-
Page 9 of 13

Rules—Base solution
1.
Write reactions as ½ reactions—One for the oxidation and one for the reduction
2.
Balance all elements except H and O
3.
Balance H with H
+
4.
Balance O with H
2
O
5.
Add OH to both sides to cancel the H +
6.
Cancel out any extra water and OH
-
7.
Balance Charge with e
-
8.
Multiply reactions by factors such that the e
-
Add both ½ reactions
cancel
Al + MnO
4
-

MnO
2
+ Al(OH)
4
-
Cl
2
+ MnO
4
-

MnO
2
+ ClO-
Page 10 of 13

I. Rules for Oxidation Numbers

The oxidation number of any free element (an element not combined chemically with a different element) is zero, regardless of how complex its molecules might be.

The oxidation number for any simple, monoatomic ion is equal to the charge on the ion.

The sum of all the oxidation numbers of the atoms in a molecule or ion must be equal the charge of the particle.

In its compound, fluorine has an oxidation of -1

In its compounds, hydrogen has an oxidation number of +, except hen combined with metals when it is -1

In its compounds, oxygen has an oxidation number of -2, except in forming peroxides hen it is -1.

All Group I metals in compounds have a charge of +1

All Group II metals in compounds have a charge of +2

All Halogens have a charge of -1 when combined with metals in binary compounds
II. Balancing Redox Equations
A. Oxidation States Method (we use the half-reaction method)
1.
Assign oxidation numbers to atoms in equation
2.
Identify the substance oxidized and determine the number of electrons lost
3.
Balance the electrons gained
4.
balance the non redox substances by inspection
B. Half-Reaction (Acidic Solution)
1.
Divide the skeleton equation into half-reactions
2.
Balance atoms other than H & O
3.
Balance oxygen by adding H
2
O to the side that needs O
4.
Balance hydrogen by adding H+ to the side that needs H
5.
Balance the charge by adding electrons
6.
Make the number of electrons gained equal to the number lost and then add the two halfreactions
7.
Cancel anything that is the same on both sides.
C. Half-Reactions Method (Basic Solution)
1.
1-7 Same as B above
2.
Add to both sides of the equation the same number of OH- as there are H+
3.
Combine H+ and OH- to form H
2
O
4.
Cancel any H
2
O that you can
Page 11 of 13

III. Types of Redox Reactions
A. Simple Redox
1. Hydrogen Displacement
Ca(s) + 2H
2
O(l)

Ca(OH)
2
(s) + H
2
2. Metal Displacement
Zn(s) + CuSO
4
(aq)

ZnSO
4
(aq) + Cu(s)
3. Halogen Displacement
Cl
2
(g) + KBr(aq)

2KCl(aq) + Br
2
(l)
4. Combustion
CH
4
(g) + 2O
2
(g)

CO
2
(g) + 2H
2
O(g)
B. Disportionation
This is where one substance both oxidizes and reduces
Cl
2
(g) + 2OH
-
(aq)

OCl
-
(aq) + Cl
-
(aq) + H
2
O(l)
C. Reactions involving oxoanions such as Cr
14H
+
(aq) + Cr
2
O
7
Prediction Practice with Redox
2-
+ 6 Fe
2+
2
O
7
2 -

2 Cr
3+
+ 7 H
2
O + 6 Fe
3+
1.
solid copper is added to a dilute nitric acid solution.
2.
a solution of potassium permanganate is mixed with an alkaline solution of sodium sulfite.
3.
ethanol is completely burned in air.
4.
sodium metal is added to water.
5.
hydrogen peroxide solution is added to a solution of iron(II) sulfate.
Page 12 of 13

IV. Redox Reaction Prediction
Important Oxidizers
MnO
4
-
(acid solution)
MnO
4
-
(basic solution)
MnO
2
(acid solution)
Cr
2
O
7
-2
(acid)
CrO
4
-2
HNO
HNO
H
2
SO
3
3
4
, conc
, dilute
, hot conc
Metallic Ions
Free Halogens (F
2
, Cl
2
, Br
2
, I
2
)
HClO
4
Na
2
O
2
H
2
O
2
Formed in reaction
Mn
+2
MnO
2
Mn +2
Cr
+3
Cr
+3
NO
NO
SO
2
2
Metallous Ions
Halide ions (F , Cl , Br , I )
Cl
-
OH
-
O
2
Important Reducers
Halide Ions
Free Metals
Metalous Ions
Nitrite Ions
Sulfite Ions
Free Halogens (dil, basic, sol)
Free Halogens (conc, basic sol)
C
2
O
4
2-
Formed in Reaction
Halogens
Metal Ions
Metallic ions
Nitrate Ions
SO
4
2-
Hypohalite ions
Halate ions
CO
2

Redox reactions involve the transfer of electrons. The oxidation numbers of at least two elements must change. Single replacement, some combination and some decomposition reactions are redox reactions.

To predict the products of a redox reaction, look at the reagents given to see if there is both an oxidizing agent and a reducing agent. When a problem mentions an acidic or basic solution, it is probably is redox.
Page 13 of 13