A.P. Chemistry Ch. 4 Notes (pg. # 1)
Reactions in Aqueous Solution
Types of Reactions in Aqueous Solution:
1.) Precipitation Reactions
2.) Acid-Base Reactions
3.) Oxidation-Reduction Reactions (R_______)
In an aqueous solution,
solvent =
solute= substance d________ in w________
ex. NaCl
Solute Concentration:
Concentration- Expression of the amount of s________ dissolved
in a certain amount of s_______.
Concentration Unit: M_________
Molarity = __________
M = _______
ex. 6.0 mol of NaOH dissolved in 1.00 L of solution:
M = ______ = ________
Units:
mol/L, molar, M
Ex. 84.2 g of NaCl is dissolved in enough water to make 2.51 L of
solution. What is the molarity of the solution?
Preparing a Solution:
1.) Dissolve the solute in a s______ amount of water.
2.) D______ the solution up to the desired volume.
Finding the number of moles of solute given the volume and
concentration:
Ex. How many moles of solute are there in 945 mL of a
3.71 M solution?
Dilution:
In a dilution, the number of m______ of s______ does not
change.
So:
st
1 solution:
2nd solution after dilution:
M1 = mol1
M2 =
L1
Mol1 = Mol2
So:
(M1)(L1) =
Ex. 0.025 L of 12 M HCl is diluted to 0.100 L. What is the new
concentration of the acid?
A.P. Chemistry Ch. 4 Notes (pg. # 2)
Precipitation Reactions:
-A reaction that forms an in_________ precipitate when two
solutions of i_____ compounds are mixed.
Reason Why a Precipitate Forms: Not all ionic compounds
are s
in water. (see figure 4.3 on pg. 80)
ex. NaCl and AgNO3 are soluble.
______ is insoluble
So: When solutions of AgNO3 and NaCl are combined, s_____
_____ precipitates.
NaCl (aq) + AgNO3 (aq) 
Net Ionic Equations: Shows only the r_______ ions. S______ ions
are not included.
(note: Ionic compounds in solution exist as their _____)
Ex.
NaCl 
AgNO3 
The reaction:
Na+ + Cl- + Ag+ + NO3-  AgCl (s) + __ + __
Net Ionic Equation:
____________________________
Ex. Write the net ionic equation for the reaction when solutions of
BaCl2 and Na2SO4 are combined.
First Step: Write the ions that are present in solution:
BaCl2 
Na2SO4 
Second Step: Determine the possible p________.
Third Step: Choose the reaction that produces the insoluble
compound.
Ex. Write the net ionic equation for the reaction when solutions of
CuSO4 and NaOH are combined.
Ex. Write the net ionic equation for the reaction when solutions of
CaCl2 and Na2CO3 are combined.
A.P. Chemistry Ch. 4 Notes (pg. # 3)
Stoichiometry of Precipitation Reactions:
Key: Use m________ to convert between m_______ and v________.
Ex. What volume of 0.200 M CuSO4 is required to react with
50.0 mL of 0.100 M NaOH?
Step 1: Write the B_________ n_________ i________equation for
the reaction.
Step 2: Use molarity to find moles of the given reacting ion.
Step 3: Use the balanced equation to find m_______ of the desired
answer.
Step 4: Use m_______ to find the volume.
(answer: 0.0125 L of CuSO4)
Ex. What volume of 0.500 M CaCl2 is required to react with
25.0 mL of 0.300 M AgNO3?
Ex. What volume of 0.715 M NaOH is required to react with
2.00 L of 1.00 M MgCl2?
A.P. Chemistry Ch. 4 Notes (pg. # 4)
Acids and Bases:
Acid- A substance that produces _____ in solution.
Base- A substance that produces _____ in solution.
(The above definitions are according to Arrhenius)
Strong Acids- Acids that c________ d_______ in water.
Ex. HCl  ___ + ____
A__ of the HCl molecules dissociate
Common Strong Acids:
HCl, HNO3, H2SO4
Ex. HNO3 
Weak Acids- Only p________ dissociate in water.
Ex. HC2H3O2  ___ + _______
(Note: The d_____ arrow indicates the e______
of the weak acid with its ions)
Another common weak acid:
HF 
Strong Bases- C________ dissociate in water.
Ex. NaOH 
Common examples: KOH, NaOH, LiOH
Weak BasesEx. NH3 (aq) + H2O (l) 
- Weak bases produce OH- by removing an ____ from the
____ molecule.
- The reaction does _____ go to completion, thus the
double arrows.
Acid-Base Reactions; N
1.)
Strong Acid + Strong Base
Ex. HCl + NaOH
HCl (aq) + NaOH (aq) 
Net Ionic equation:
2.)
Weak Acid + Strong Base
Ex. HF + NaOH
HF (aq) + NaOH (aq) 
Net Ionic Equation:
3.)
Strong Acid + Weak Base
Ex. HCl + NH3
HCl (aq) + NH3 (aq) 
Net Ionic Equation:
Reactions:
A.P. Chemistry Notes: Ch. 4 (pg. # 5)
Oxidation-Reduction Reactions: (Redox)
Ex.
Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s)
Oxidation: L_______ of one or more e________.
Reduction: G_______ of one or more e________.
Oxidation and Reduction always occur t__________.
Oxidizing Agent- The species that receives the electrons. (Oxidizing
agents undergo _________.)
Reducing Agent- the species that gives up the electrons. (Reducing
agents undergo __________.)
Ex. Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s)
______ is the reducing agent
______ is the oxidizing agent
Ex. HCl (g) + HNO3 (l)  NO2 (g) + ½ Cl2 (g) + H2O (l)
Which is the oxidizing agent, which is the reducing
agent?
Oxidation Number- A “pseu________” assigned to determine if
something is oxidized or reduced.
Oxidation- An increase in oxidation number
Reduction- A decrease in oxidation number
A_________ Rules for Assigning Oxidation Numbers:
1.) The oxidation number of an element in its elemental state is
____.
Ex. Na___ P____ Cl2 _____
2.) The oxidation number of a monatomic ion is equal to the charge
of that ion. Ex. Ca2+ ___
Br-___
3.) In compounds, the group 1 metals have an oxidation number of
+1. The group 2 metals have an oxidation number of _____.
Ex.
NaCl, Na = ____
CaCl2, Ca = ____
4.) In compounds, H is ordinarily +1, O is ordinarily -2, F is always 1
5.) The sum of the oxidation numbers of the elements in a
compound equals ____ in a neutral compound. In a polyatomic
ion, the sum of the oxidation numbers equals the ______ of
the ion.
Ex. NO3Ex. HCl
To Determine if a species is oxidized or reduced, compare the
oxidation number of the elements in the reactants to the products.
Ex. HCl (g) + HNO3 (l)  NO2 (g) + ½ Cl2 (g) + H2O (l)
Which is the oxidizing agent, which is the reducing
agent?
1.)
Determine the oxidation numbers of the elements in :
HCl
HNO3
 NO2
Cl2
H2O
A.P. Chemistry Ch. 4 Notes (pg. # 6)
Half Reactions:
Ex. Split the following redox reaction into an oxidation halfreaction and a reduction half-reaction:
Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s)
1st:
___________
2nd:
___________
Ex. Classify each of the following unbalanced half-reactions as
either oxidation or reduction:
TiO2 (s)  Ti3+ (aq)
NH4+ (aq)  N2 (g)
CH3OH (aq)  CH2O (aq)
Balancing Redox Reactions:
Ex.
Balance the following redox reaction:
ClO3- + I-  Cl- + I2
Steps to Balancing Redox Reactions:
1.) Split the reaction into two half reactions:
2.)
Balance all the elements other than O and H.
3.)
Balance Oxygen by adding H2O, and balance Hydrogen by
adding H+. (assuming it is an acidic solution)
4.)
Balance the charge by adding electrons.
5.)
Combine the half reactions. Multiply the half reactions by a
coefficient if necessary to allow the electrons to cancel out.
Ex. Fe2+ + Cr2O72-  Fe3+ + Cr3+