Chapter 4
Reactions in Aqueous Solutions
This chapter introduces to the student the concepts of reactions in aqueous solutions. Because water is
used so universally as a solvent, it is important for the student to understand basic concepts of reactions that occur in
it. Upon completion of Chapter 4, the student should be able to:
1. Distinguish between solute, solvent, and solution.
2. Classify common compounds as strong electrolytes, weak electrolytes, or nonelectrolytes (strong or weak acids
or strong or weak bases).
3. Suggest why water is often called a universal solvent utilizing the terms polar solvent, dissociation, ionization,
and hydration.
4. Describe precipitation reactions using the terms solubility and precipitate.
5. Classify common ionic compounds as soluble or insoluble.
6. Predict the resulting products and write the molecular equation, ionic equation, and net ionic equation and
identify spectator ions given the reactants of a chemical reaction.
7. Distinguish between Arrhenius acids and bases and Brφnsted acids and bases.
8. Compare and contrast the properties of acids and bases.
9. List common examples of monoprotic, diprotic and triprotic acids.
10. Justify how some ions can act as an acid or as a base (amphoteric).
11. Explain, by using a chemical equation, how ammonia (NH3) is classified as a Brφnsted base.
12. Predict the products formed by acid-base neutralization reactions.
13. Discuss what factor results in an oxidation-reduction reaction.
14. Identify oxidation half-reactions, reduction half-reactions, oxidizing agents and reducing agents.
15. Assign oxidation numbers to elements in compounds and ions.
16. Categorize redox reactions in terms of combination reactions, decomposition reactions, displacement reactions,
and disproportionation reactions.
17. Predict the results of a chemical reaction involving metals given the activity series (electrochemical series).
18. Predict the results of a chemical reaction involving halogens given the halogen activity series.
19. Compute the molarity of a solution given the mass (number of moles) of solute and the volume of solution.
20. Describe the method for preparing a specific molar solution given the volume of solution required and the solute
to be used.
21. Relate in detail how to prepare a specific dilute solution given a known stock solution using dilution techniques.
22. Predict the mass of a precipitate formed using gravimetric analysis methods.
23. Deduce the mass percent of specific ions present in an original solution given the results of a gravimetric
analysis.
24. Use the terms titration, standard solution, equivalence point, and indicator to describe quantitative studies of
acid-base neutralization reactions.
25. Determine the concentration of an unknown acid (base) given the results of an acid-base titration.
26. Predict the amount (mass, moles, or volume of solution) of an acid (base) required to neutralize a base (acid).
27. Predict the volume of an oxidizing (reducing) agent solution required to oxidize (reduce) a specific volume of
reducing (oxidizing) agent solution provided that the net ionic equation is given.
Section 4.1
General Properties of Aqueous Solutions
Solutions are composed of the solvent (major component) and one or more solutes (minor components). If
a solution readily conducts electricity, the solute is classified as a strong electrolyte, if the solution weakly conducts
electricity, the solute is a weak electrolyte, and if the solution does not conduct electricity, the solute is a
nonelectrolyte. It should be noted that conductivity is concentrated dependent. It is possible for a concentrated
solution of a weak electrolyte to have a higher conductivity than a dilute solution of a strong electrolyte. Your author
defines dissociation as the breaking up of compounds into cations and anions. He reserves the term ionization to
describe the separation of acids or bases into their respective ions.
The reason that so many materials dissolve in water is because water is polar. It has, within the molecule, a
region that is positive (the H atoms) and a region that is negative (the O atom). The reason for this polarity will be
described in more detail later; however, in summary, it is because of the difference in electronegativity. It is not
because the oxygen atom has two lone pairs of electrons on it. The key idea for water acting as such a good solvent
is that “like dissolves like”. Polar water molecules will dissolve ionic materials while nonpolar solvents will not
dissolve ionic materials.
Acetic acid, CH3COOH, is used as an example of a weak electrolyte.
CH3COOH(aq) U CH3COOu(aq) + H+(aq)
Students may have a difficult time understanding the concept that the concentration of CH3COOH remains constant
when this reaction is at equilibrium, yet CH3COOH molecules are constantly ionizing to form CH3COO- and H+ ions
while CH3COO- and H+ ions are continually combining to form CH3COOH . One “thought” experiment that could be
described is to start with an equal mixture of CD3COOD (molar mass equals 64 g/mol) and CH3COOH (molar mass
equals 60 g/mol). When this mixture is dissolved in water and equilibrium is reached, the resultant will be a mixture
of not only CH3COOH and CD3COOD, but also CH3COOD (molar mass of 61 g/mol) and CD3COOH (molar mass of
63 g/mol). This “thought” experiment shows how when either CH3COOH or CD3COOD ionize, then the resultant H+ or
D+ is free to recombine with either CH3COO- or CD3COO-.
Section 4.2
Precipitation Reactions
The concept of what is soluble and what is insoluble is often misunderstood by students. Your author
suggests that a substance is soluble if “a fair amount of it visibly dissolves when added to water”. An interesting
question to pose is “Given that the solubility of malathion (C10H19O6PS6) is 145 mg/L, is malathion soluble or
insoluble”? Many students will see that a solubility is given and, therefore, assume that malathion must be soluble.
More experienced chemists would assume that 145 mg/L is low enough to classify malathion as only very slightly
soluble or insoluble.
A useful hint to assist students in identifying net ionic equations is to indicate that they involve either the
decomposition or formation of a solid, liquid, or gas. For example, the net ionic equation for the reaction of aqueous
barium chloride and aqueous sodium sulfate is
Ba2+(aq) + SO 24 − (aq)
BaSO4(s)
The result is the formation of a solid. The reaction of hydrochloric acid with sodium hydroxide results in the net ionic
equation of:
H+(aq) + OH-(aq)
H2O(ℓ)
A liquid is formed. It should be noted that all reactions involving strong acids with strong bases result in the same net
ionic equation because by definition strong acids and strong bases will result in the formation of soluble salts when
neutralization takes place.
A more complex example of finding the net ionic equation is the reaction of hydrochloric acid with aqueous
sodium bicarbonate. The net ionic equation is
H+(aq) + HCO3- (aq)
H2O(ℓ) + CO2(g)
This is a direct result of knowing that H2O(ℓ) and CO2(g) must each appear in the net ionic equation.
An example of a net ionic equation that involves the decomposition of a solid is the reaction of hydrochloric
acid with solid magnesium. The net ionic equation is
2H+(aq) + Mg(s)
Mg2+(aq) + H2(g)
It is also noted that this net ionic equation forms a gas.
Students should be cautioned that for ALL balanced equations to be correct, not only must there be a mass
balance, but also a charge balance. That is, the number of each atom of each element must be the same on both
the left and right sides and the sum of the charges on the left must equal the sum of the charges on the right.
Section 4.3
Acid-Base Reactions
Students are generally familiar with common acid-base reactions. Your author points out that both acids
and bases, when in solution, will conduct electricity. This is certainly true; however, the converse is not necessarily
true. That is to say that not all solutions that conduct electricity must be either acidic or basic.
Sulfuric acid is an unusual acid in that the first step in its ionization process is complete (making it a strong
electrolyte or strong acid) while the second step is not complete suggesting that it is a weak acid. The whole concept
of strong and weak acids is often confused by students. Most students have probably worked with hydrochloric acid
(a strong acid) and have seen that it dissolves zinc metal for example. They probably have heard that sulfuric acid
(another example of a strong acid) is in car batteries and have been warned of its dangers. Since acetic acid (a
weak acid) is found in vinegar and phosphoric acid (another weak acid) is found in soda pop, some students may
assume that weak acids are not hazardous. This, of course, is certainly NOT true. It is important that students
understand that the concentration of the acid is what tends to make it dangerous and not the fact that it is a weak or
strong acid. Certainly concentrated acetic acid and concentrated phosphoric acid ARE dangerous.
Section 4.4
Oxidation-Reduction Reactions
Oxidation-Reduction reactions are better known as redox reactions. Redox reactions must involve the
transfer of electrons which require that oxidation numbers must change. A “fun” way to assist students with redox
reactions is to recall that
LEO (the lion) says GER.
This suggests that the Loss of Electrons is Oxidation and Gain of Electrons is Reduction. Also note that “secret
agents” cause things to happen; therefore, oxidizing agents cause oxidation to take place (they are thus reduced)
and reducing agents cause reduction to occur (they are oxidized). Since both the mass and charge must balance for
an equation to be correct, the number of electrons lost by a reducing agent must equal the number of electrons
gained by an oxidizing agent.
When determining oxidation numbers (oxidation states), it is important to recall that the sum of the oxidation
numbers of a species (atom, molecule or ion) must equal the charge on that species. For example, the sum of the
oxidation states of sulfur and two oxygens in SO2 must sum to zero since SO2 is neutral. The sum of the oxidation
states of sulfur plus four oxygens in SO42- must sum to negative two since the sulfate ion has a negative two charge.
The oxidation number of Fe3+ must be +3 since the charge on Fe3+ is +3.
The metal activity series (Table 4.15) along with the halogen displacement series provides information
which can be used to predict if reactions will occur. The ability to predict what will happen is an example of the power
of science.
Section 4.5
Concentration of Solutions
Chapter 12 will deal with several other types of concentration units but the unit that most students have
used is likely molarity. Molarity is independent of how much solution we have, thus molarity is an intensive property.
The student should be instructed to read [K+] as the concentration of potassium ions in moles per liter (Molarity).
Certainly for dilutions the moles of solute before dilution must equal the moles of solute after dilution and, therefore,
equation 4.2 is correct, however, the student should be warned that the equation is for dilution only.
Mi Vi = Mf Vf
Your author states that quantitative analysis is the determination of the amount or concentrations of a
substance in a sample. This is in contrast to qualitative analysis which is the study of what species are present.
Some of your students may have performed qualitative analysis in high school but likely they have not done
quantitative analysis before.
Section 4.6
Gravimetric Analysis
It is possible to use the precipitation reactions to determine the amount of specific substances in a solution.
This is known as gravimetric analysis. Your author correctly points out that for a gravimetric analysis to be accurate,
the solid formed must be very insoluble (nearly 100% yield) otherwise there will be an inherent error in the procedure.
The concept of gravimetric analysis reaffirms the solubility rules stated earlier and gives some practical examples for
their use.
Section 4.7
Acid-Base Titration
The acid that is often used as a standard for accurate titrations is potassium hydrogen phthalate (KHP,
molar mass equals 204 g/mol). Students have been known to think that KHP corresponds to potassium hydrogen
phosphorous! That may sound unbelievable, but in fact it has actually happened.
Section 4.7
Redox Titrations
Your author has included redox titrations to complete the list of types of titrations that are commonly done.
Your students may be concerned about knowing how to balance redox reactions which will be covered in a later
chapter.