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Chapter 4. Reactions in Aqueous Solution
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 4.4 A Precipitation Reaction 4.2 Precipitation Reactions
Table 4.1 Solubility Guidelines for Common Ionic 4.2 Precipitation Reactions
Problems
Animations: Section:
Electrolytes and Nonelectrolytes 4.1 General Properties of Aqueous Solutions
Dissolution of NaCl in Water 4.1 General Properties of Aqueous Solutions
Introduction to Aqueous Acids 4.3 Acids, Bases, and Neutralization Reactions
Acid-Base Titration 4.6 Solution Stoichiometry and Chemical Analysis
Movies: Section:
Strong and Weak Electrolytes 4.1 General Properties of Aqueous Solutions
Activities: Section:
Ionic Compounds 4.2 Precipitation Reactions
Writing a Net Ionic Equation 4.2 Precipitation Reactions
3-D Models: Section:
Water 4.1 General Properties of Aqueous Solutions
Hydrogen Chloride 4.1 General Properties of Aqueous Solutions
Reactions in Aqueous Solution
45
Ethanol 4.1 General Properties of Aqueous Solutions
Formic Acid 4.1 General Properties of Aqueous Solutions
Other Resources
Further Readings: Section:
Solubility Rules: Three Suggestions for Improved 4.2 Precipitation Reactions
Understanding
An Analogy for Solubility: Marbles and Magnets 4.2 Precipitation Reactions
Reinforcing Net Ionic Equation Writing 4.2 Precipitation Reactions
Live Demonstrations: Section:
Conductivity and Extent of Dissociation of Acids 4.1 General Properties of Aqueous Solutions
in Aqueous Solution
Name That Precipitate 4.2 Precipitation Reactions
Solubility of Some Silver Compounds 4.2 Precipitation Reactions
Alka Seltzer Poppers: an Interactive Exploration 4.3 Acids, Bases, and Neutralization Reactions
Food Is Usually Acidic, Cleaners Are Usually Basic 4.3 Acids, Bases, and Neutralization Reactions
A Hand-Held Reaction: Production of Ammonia Gas 4.3 Acids, Bases, and Neutralization Reactions
Chapter 4
46
Floating Pennies 4.4 Oxidation-Reduction Reactions
A Cool Drink! An Introduction to Concentrations 4.5 Concentrations of Solutions
Chapter 4. Reactions in Aqueous Solution
Common Student Misconceptions
• Molarity is moles of solute per liter of solution, not per liter of solvent.
• Students sometimes use moles instead of molarity in MinitialVinitial = MfinalVfinal.
• Students often disregard rules for significant figures when calculating or using molarities.
• Students sometimes think that water is a good conductor.
• Students sometimes have a problem with the arbitrary difference between strong and weak
electrolytes.
Teaching Tips
• Weaknesses in recollection of ionic nomenclature and the structure of common ions often make it
difficult for students to write molecular, complete ionic, and net ionic equations for metathesis
reactions. A brief review of ionic nomenclature is often useful prior to covering metathesis reactions.
Reactions in Aqueous Solution
47
Lecture Outline
4.1 General Properties of Aqueous Solutions
1
• A solution is a homogeneous mixture of two or more substances.
Electrolytic Properties
2
• All aqueous solutions can be classified in terms of whether or not they conduct electricity.
Ionic Compounds in Water
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• When an ionic compound dissolves in water, the ions are said to dissociate.
• This means that in solution, the solid no longer exists as a well-ordered arrangement of ions in
contact with one another.
• Instead, each ion is surrounded by several water molecules; it is called an aqueous ion denoted
Molecular Compounds in Water
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• When a molecular compound (e.g. CH3OH ) dissolves in water, the solution usually consists of intact
molecules dispersed homogeneously in the solution.
• Therefore, there is nothing in the solution to transport electric charge and the solution does not
conduct electricity and most molecular compounds are nonelectrolytes.
• There are some important exceptions.
• For example, HCl(g) in water ionizes to form H+(aq) and Cl–(aq).
Strong and Weak Electrolytes
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• Compounds whose aqueous solutions conduct electricity well are called strong electrolytes.
• These substances exist in solution mostly as ions.
• Example: NaCl
Chapter 4
48
NaCl(aq) → Na+(aq) + Cl–(aq)
• The single arrow indicates that the Na+ and Cl– ions have no tendency to recombine to form
NaCl.
• In general, soluble ionic compounds are strong electrolytes.
• Other strong electrolytes include strong acids and soluble strong bases.
• Compounds whose aqueous solutions conduct electricity poorly are called weak electrolytes.
• These substances exist as a mixture of ions and un-ionized molecules in solution.
• This balance produces a state of chemical equilibrium.
FORWARD REFERENCES:
• Double arrows () will be used in the chapter on chemical equilibria (Chapter 15) and
beyond.
4.2 Precipitation Reactions
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• Reactions that result in the formation of an insoluble product are known as precipitation reactions.
• A precipitate is an insoluble solid formed by a reaction in solution.
• Example: Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3(aq)
Solubility Guidelines for Ionic Compounds
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• The solubility of a substance at a particular temperature is the amount of that substance that can be
dissolved in a given quantity of solvent at that temperature.
• A substance with a solubility of less than 0.01 mol/L is regarded as being insoluble.
• Experimental observations have led to empirical guidelines for predicting solubility.
• Solubility guidelines for common ionic compounds in water:
Reactions in Aqueous Solution
49
• Exceptions are the compounds of Sr2+, Ba2+, Hg22+, and Pb2+.
• Compounds containing S2– are insoluble.
Exchange (Metathesis) Reactions
• Exchange reactions, or metathesis reactions, involve swapping ions in solution:
AX + BY → AY + BX.
• Many precipitation and acid-base reactions exhibit this pattern.
Ionic Equations
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• The complete ionic equation lists all strong soluble electrolytes in the reaction as ions:
Pb2+(aq) + 2NO3–(aq) + 2K+(aq) + 2I–(aq) → PbI2(s) + 2K+(aq) + 2NO3–(aq)
• Only strong electrolytes dissolved in aqueous solution are written in ionic form.
• Weak electrolytes and nonelectrolytes are written in their complete chemical form.
• The net ionic equation lists only those ions which are not common on both sides of the reaction:
Pb2+(aq) + 2I–(aq) → PbI2(s)
FORWARD REFERENCES:
• Net ionic equations will be frequently used in chapters dealing with acid-base reactions
(Chapters 16 and 17) as well as in electrochemistry (Chapter 20, Appendix E)
• Equilibria involving insoluble or poorly soluble compounds and their ions will be discussed
in more detail in Chapter 17 (section 17.4).
4.3 Acids, Bases, and Neutralization Reactions
Acids
• Acids are substances that are able to ionize in aqueous solution to form H+.
• Ionization occurs when a neutral substance forms ions in solution.
Chapter 4
50
• Acids that ionize to form one H+ ion are called monoprotic acids.
• Common monoprotic acids include HCl, HNO3 and HC2H3O2.
• Acids that ionize to form two H+ ions are called diprotic acids.
• A common diprotic acid is H2SO4.
Bases
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Strong and Weak Acids and Bases
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• Strong acids and strong bases are strong electrolytes.
• They are completely ionized in solution.
• Strong bases include: Group 1A metal hydroxides, Ca(OH)2, Ba(OH)2, and Sr(OH)2.
• Strong acids include: HCl, HBr, HI, HClO3, HClO4, H2SO4, and HNO3.
• We write the ionization of HCl as: HCl → H+ + Cl–
Identifying Strong and Weak Electrolytes
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• Compounds can be classified as strong electrolytes, weak electrolytes, or nonelectrolytes by looking
at their solubility.
• Strong electrolytes:
• Soluble ionic compounds are strong electrolytes.
• Molecular compounds that are strong acids are strong electrolytes.
• Weak electrolytes:
• Weak acids and bases are weak electrolytes.
• Nonelectrolytes:
Reactions in Aqueous Solution
51
• All other compounds, including water.
Neutralization Reactions and Salts
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• A neutralization reaction occurs when an acid and a base react:
• HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)
• (acid) + (base) (water) + (salt)
• In general, an acid and a base react to form a salt.
Neutralization Reactions with Gas Formation
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• There are many bases besides OH– that react with H+ to form molecular compounds.
• Reaction of sulfides with acid gives rise to H2S(g).
• Sodium sulfide (Na2S) reacts with HCl to form H2S(g):
• Molecular equation:
Na2S(aq) + 2HCl(aq) → H2S(g) + 2NaCl(aq)
• Net ionic equation: 2H+(aq) + S2–(aq) → H2S(g)
FORWARD REFERENCES:
• Strong acids and bases will be revisited in Chapter 16.
• Strong acids and bases will be used as titrants in acid-base titrations (Chapter 17)
• Equilibria involving weak acids and bases will be further discussed in Chapters 16 and 17.
• Environmental impact of weak acid equilibria will be discussed on Chapter 18.
Chapter 4
52
4.4 Oxidation-Reduction Reactions
Oxidation and Reduction
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• Oxidation-reduction, or redox, reactions involve the transfer of electrons between reactants.
• When a substances loses electrons, it undergoes oxidation:
Ca(s) + 2H+(aq) → Ca2+(aq) + H2(g)
Oxidation Numbers
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• Electrons are not explicitly shown in chemical equations.
• Oxidation numbers (or oxidation states) help up keep track of electrons during chemical reactions.
• Oxidation numbers are assigned to atoms using specific rules.
• For an atom in its elemental form, the oxidation number is always zero.
• For any monatomic ion, the oxidation number equals the charge on the ion; positive for metals
and negative for nonmetals.
• The oxidation number of oxygen is usually –2.
Oxidation of Metals by Acids and Salts
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• The reaction of a metal with either an acid or a metal salt is called a displacement reaction.
• The general pattern is: A + BX → AX + B
42
“Reduction of CuO” Movie from Instructor’s Resource CD/DVD
43
“Oxidation-Reduction Reactions: Part I” Animation from Instructor’s Resource CD/DVD
44
“Oxidation-Reduction Reactions: Part II” Animation from Instructor’s Resource CD/DVD
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“Oxidation and Reduction” from Further Readings
Reactions in Aqueous Solution
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• Example: It is common for metals to produce hydrogen gas when they react with acids. Consider
the reaction between Mg and HCl:
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
The Activity Series
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• We can list metals in order of decreasing ease of oxidation.
• This list is an activity series.
• The metals at the top of the activity series are called active metals.
• The metals at the bottom of the activity series are called noble metals.
FORWARD REFERENCES:
• Oxidation numbers will be frequently used in electrochemistry (Chapter 20, Appendix E).
• Balancing of redox reactions will be covered in Chapter 20.
4.5 Concentrations of Solutions
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• The term concentration is used to indicate the amount of solute dissolved in a given quantity of
solvent or solution.
Molarity
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• Solutions can be prepared with different concentrations by adding different amounts of solute to
solvent.
Chapter 4
54
Expressing the Concentration of an Electrolyte
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• When an ionic compound dissolves, the relative concentrations of the ions in the solution depend on
Interconverting Molarity, Moles, and Volume
• The definition of molarity contains three quantities: molarity, moles of solute, and liters of solution.
• If we know any two of these, we can calculate the third.
• Dimensional analysis can be helpful in these calculations.
Dilution
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• A solution in concentrated form (stock solution) is mixed with solvent to obtain a solution of lower
solute concentration.
FORWARD REFERENCES:
• Molarity will be used throughout the course as the most common form of concentration.
• The concept of molarity is not limited to solutions; one can calculate molarity for gases and
use them in Kc expressions in Chapter 15 and beyond.
4.6 Solution Stoichiometry and Chemical Analysis
68
• In approaching stoichiometry problems:
• recognize that there are two different types of units:
• laboratory units (the macroscopic units that we measure in lab) and
• chemical units (the microscopic units that relate to moles).
Reactions in Aqueous Solution
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• Always convert the laboratory units into chemical units first.
• Convert grams to moles using molar mass.
• Convert volume or molarity into moles using M = mol/L.
Titrations
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• A common way to determine the concentration of a solution is via titration.
• We determine the concentration of one substance by allowing it to undergo a specific chemical
reaction, of known stoichiometry, with a substance with known concentration (standard solution).
• Monoprotic acids and bases react with each other in a stoichiometric ratio of 1:1.
• Example: Suppose we know the molarity of an NaOH solution and we want to find the molarity of an
HCl solution.
• What do we know?
• molarity of NaOH, volume of HCl
• What do we want?
• molarity of HCl
• What do we do?
FORWARD REFERENCES:
• Acid-base titrations will be discussed in detail in Chapter 17.
69
“Acid-Base Titration” Animation from Instructor’s Resource CD/DVD
70
“Colorful Acid-Base Indicators” from Live Demonstrations
Chapter 4
56
Further Readings:
1. Bob Blake, “Solubility Rules: Three Suggestions for Improved Understanding,” J. Chem. Educ., Vol.
80, 2003, 1348–1349.
7. Albert Kowalak, “When Is a Strong Electrolyte Strong?” J. Chem. Educ., Vol. 65, 1988, 607.
8. John J. Fortman, “Pictorial Analogies X: Solutions of Electrolytes,” J. Chem. Educ., Vol. 71, 1994,
27–28.
9. Gian Calzaferri, “Oxidation Numbers,” J. Chem. Educ., Vol. 76, 1999, 362–363.
10. R. Lipkin, “What Makes Gold Such a Noble Metal?” Science News, July 22, 1995, 62.
11. Arthur M. Last, “A Cyclist’s Guide to Ionic Concentration,” J. Chem. Educ., Vol. 75, 1998, 1433.
12. Lloyd J. McElroy, “Teaching Dilutions,” J. Chem. Educ., Vol. 73, 1996, 765–766.
Reactions in Aqueous Solution
57
Live Demonstrations:
1. Bassam Z. Shakhashiri, “Conductivity and Extent of Dissociation of Acids in Aqueous Solution,”
Chemical Demonstrations: A Handbook for Teachers of Chemistry, Volume 3 (Madison: The University
of Wisconsin Press, 1989), pp. 140–145. Universal indicator and a conductivity probe are used to explore
the relative acidity and conductivity of a series of aqueous acids.
5. Bassam Z. Shakhashiri, “Food is Usually Acidic, Cleaners Are Usually Basic,” Chemical
Demonstrations: A Handbook for Teachers of Chemistry, Volume 3 (Madison: The University of Wisconsin
Press, 1989), pp. 65–69. The pH of a variety of household chemicals is determined using indicators and pH
meters.
6. Lee R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “A Hand-Held Reaction: Production of
Ammonia Gas,” Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 Washington: American
Chemical Society, 1988), p. 38. An example of a reaction involving two solids (NH4Cl and Ca(OH)2) is
demonstrated.
10. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “Milk of Magnesia versus Acid,”
Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical
Society, 1988), p. 173. An antacid, milk of magnesia, is mixed with acid in this demonstration.
Chapter 4
58
14. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “Making Hydrogen Gas from an Acid and
a Base,” Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American
Chemical Society, 1988), pp. 33–34. Hydrogen gas is collected as a product of the reaction of aluminum
with either HCl or NaOH.
15. Bassam Z. Shakhashiri, “An Activity Series: Zinc, Copper, and Silver Half Cells,” Chemical
Demonstrations: A Handbook for Teachers of Chemistry, Volume 4 (Madison: The University of Wisconsin
Press, 1992), pp. 101–106.
18. Bassam Z. Shakhashiri, “Colorful Acid-Base Indicators,” Chemical Demonstrations: A Handbook for
Teachers of Chemistry, Volume 3 (Madison: The University of Wisconsin Press, 1989), pp. 33–40.
19. Bassam Z. Shakhashiri, “Rainbow Colors with Mixed Acid-Base Indicators,” Chemical
Demonstrations: A Handbook for Teachers of Chemistry, Volume 3 (Madison: The University of Wisconsin
Press, 1989), pp. 41–46.
