CHAPTER 5: REACTIONS IN AQUEOUS SOLUTION
5.0 OBJECTIVES
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Differentiate between nonelectrolytes and strong and weak electrolytes, theoretically and
operationally.
Know simple solubility rules and write dissociation equations.
Predict the products of precipitation reactions and write net ionic equations.
Recognize common acids and bases, their behavior in aqueous solutions, their products in
neutralization reactions, and the net ionic equation for the reaction between strong acids and
bases.
Predict the products and write reactions for the major gas forming reactions.
Recognize and balance redox reactions.
Calculate solution concentration in terms of molarity and use molarity in stoichiometry
problems.
Understand the process and uses of titration in the laboratory.
5.1 PROPERTIES OF COMPOUNDS IN AQUEOUS SOLUTION
1. Definitions – solution, aqueous, anion, cation, electrode
2. Arrhenius Theory of Electrolytic Dissolution
a. strong electrolyte
b. weak electrolyte
c. nonelectrolyte
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3. Solubility Rules
4. Dissolution equations Ex5.1 Write dissolution equations for: NaCl(s), Al(NO3)3(s),
Na2SO4(s), and MgCl2(s)
5.2 PRECIPITATION REACTIONS
1. Definitions – double replacement reaction, precipitate, full ionic equation, net ionic
equation, spectator ions
Double Replacement Rxns.
1.
2.
3.
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What is happening in the solutions?
Full (Complete) Ionic Equation
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
Net Ionic Equation
Spectator Ions
2. Ex5.2 Predict whether or not the following pairs of reactants will form precipitates.
Then write the double replacement reaction, the full ionic equation, and the net ionic
equation and identify spectator ions.
a. CuCl2 and (NH4)2SO4
b. Ba(NO3)2 and Na2CO3
c. MgCl2 and AgNO3
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3. Ex5.3 What two aqueous solutions of soluble compounds could be mixed to produce a
precipitate of CuCO3?
4. Ex5.4 You suspect that a particular unlabeled aqueous solution is one of the following:
sodium sulfate, ammonia (NH3), or barium nitrate. Explain how to use precipitation
reactions on small samples of the unknown solution to determine its identity.
5.3 ACIDS AND BASES
1. Arrhenius Definition of acids and bases
2. Strong acids – strong electrolytes
3. Weak acids – weak electrolytes
4. Strong bases – strong electrolytes
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5. Weak bases – weak electrolytes
6. Oxides of metals and nonmetals
5.4 REACTIONS BETWEEN ACIDS AND BASES
1. Neutralization reaction and net ionic equation
2. Ex5.5 Write the full neutralization reaction and net ionic equations for the interactions
between HNO3 and Ba(OH)2, and between H2SO4 and KOH.
3. Acid Base Titration
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5.5 GAS-FORMING REACTIONS
1. Definition
2. Metal and acid 
3. Metal carbonate and acid 
4. Metal sulfide and acid 
5. Metal sulfite and acid 
6. Ammonium salt and strong base
5.6 CLASSIFYING REACTIONS IN AQUEOUS SOLUTION
1. General classification
2. Redox reaction
3. Product favored reactions
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5.7 OXIDATION-REDUCTION REACTIONS
1. Examples and categories
2. Species which lose electrons are said to be
3. Species which gain electrons are said to be
4. Half-reaction examples
5. Rules for determining oxidation numbers
a. Neutral species
b. Monatomic ions, Groups 1A, 2A …
c. Oxygen and Hydrogen
d. Sum of oxidation numbers for neutral and charged species
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e. Ex5.6 Assign oxidation numbers for each element in the following species:
CaC2O4
Cr2O72-
N2O
N2O4
ClO1-
ClO41-
HAsO42-
K2O
HCO31-
MgSO4
HIO3
Na2MoO4
NO21-
NH41+
S2O32-
TeF82-
6. Rules for balancing redox reactions:
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Ex5.7 Balance in acidic solution:
MnO41- + Fe2+  Mn2+ + Fe3+
Cr2O72- + H2SO3  Cr3+ + SO42-
Ex5.8 Balance in Basic Solution
I1- + MnO41-  I2 + MnO2
S2O32- + I2  SO42- + I1-
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5.8 CONCENTRATION OF COMPOUNDS IN SOLUTION
1. Molarity – definition, equation, and notation (M and [ ])
2. Ex5.9 What is the molarity of a solution made by dissolving 20.0g of NaCl(s) in
enough water to make 1.00L of solution? In enough water to make 300.mL of solution?
3. Ex5.10 How many liters of 0.430M solution of Na2SO4 can be made starting with
10.0g of solid?
4. Ex5.11 How many grams of solid solute can be recrystallized from 10.0mL of 0.020M
CuCl2?
5. Dissolution equations and molarity
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6. Ex5.12 State the concentration of each ion in a. 0.250M Na2CO3 b. 0.023M
barium phosphate c. 0.380M glucose
7. Dilution problems and equation
8. Ex5.13 What is the new concentration of a solution made by adding 150.mL of water to
23.0mL of a 2.50M solution of NaCl?
9. Ex5.14 A student starts with100. mL of a 0.500M solution of HCl. How much water
needs to be added to obtain a solution of HCl with a concentration of 0.0250M. What is
the final volume of the solution?
5.9 CONCENTRATION SCALE FOR ACIDS AND BASES
1. pH scale – definitions and diagram
2. Ex5.15 State the pH of the following solutions with given hydrogen ion concentrations:
a. [H+] = 0.0230M
b. [H+] = 0.000560M
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c. [H+] = 3.00M
3. Ex5.16 What is the hydrogen ion concentration for each of the following solutions with
given pH?
a. pH = 4.30
b. pH = 11.2
c. pH = 6.30
5.10 STOICHIOMETRY OF REACTIONS IN AQUEOUS SOLUTION
Ex5.17 A precipitate forms when solutions of silver nitrate and scandium (III) chloride
are mixed. What volume of 0.0385M scandium (III) chloride is needed to react
completely with 22.00mL of 0.130M silver nitrate? What is the mass of the precipitate
formed?
Ex5.18 What is the mass percent of NaCl in a mixture of sodium chloride and sodium
nitrate if a 0.9056g sample of the mixture yields 0.9372g of AgCl(s) when allowed to
react with excess AgNO3(aq)?
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Ex5.19 How many milliliters of 0.0195M HCl are required to titrate 10.00mL of
0.0116M Ca(OH)2?
Ex5.20 A Vitamin C capsule (ascorbic acid) is analyzed by titrating is with 0.250M
sodium hydroxide. It is found that 10.30mL of the base is required to react with a
capsule with a mass of 0.518g. What is the percentage of vitamin C, C8H8O6, in the
capsule if the acid and base react in a 1 to 1 molar ratio?
Ex5.21 If 38.55mL of HCl are used to react with 2.150g of Na2CO3, what is the
molarity of the HCl solution?
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Ex5.22 To analyze an iron containing compound, you convert all the iron to Fe2+ in
aqueous solution and then titrate the solution with a known concentration of KMnO4
according to the following balanced equation:
MnO41-(aq) + 5Fe2+(aq) + 8H1+(aq)  Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)
A 0.598g sample of the iron containing compound requires 22.25mL of 0.0123M KMnO4
for titration to the equivalence point. What is the mass percent of iron in the sample?
Homework – Chapter 5
#1 – 21, 23, 25, 27, 29, 105, 107, 117 (solubility and ppt.)
#2 – 31, 33, 35, 37, 39, 41, 43 (dissolution rxns)
#3 – 45, 47, 49, 95, 97, 53, 55, 57 (ppt, ox#, ox or red)
#4 – Worksheet – Balancing Redox Equations
#5 – 59, 61, 63, 64, 67, 69, 71, 113 (Molarity)
#6 – 73, 75, 77, 109, 87, 89, 91, 93, 101(pH, M, stoich)
#7 – 79, 83, 99, 111, 115 (stoich)
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