Acid–Base Properties
of Salt Solutions
Salts are solids (at SATP) composed of cations and anions arranged in a crystalline lattice. When they dissolve in water, they dissociate into individual hydrated ions that may
or may not affect the pH of the solution.
Sodium chloride, NaCl(s), like all sodium salts, is a strong electrolyte. When dissolved
in water it dissociates quantitatively into aqueous sodium and chloride ions.
8.3
INVESTIGATION 8.3.1
The pH of Salt Solutions (p. 627)
Do soluble salts affect the pH of
aqueous solutions?
NaCl(s) → Na
(aq) Cl(aq)
When the pH of a sodium chloride solution is tested with a pH meter or acid–base
indicator, it shows a pH of 7 at SATP: neutral. Sodium chloride does not produce hydronium or hydroxide ions in solution, so it is classed as a neutral electrolyte.
Sodium carbonate, Na2CO3(s), is another highly soluble ionic compound. It dissociates quantitatively in water to produce aqueous sodium and carbonate ions.
2
Na2CO3(s) → 2 Na
(aq) CO3(aq)
When tested with a pH meter or acid–base indicator, the sodium carbonate solution
proves to have a pH greater than 7: basic. However, sodium carbonate cannot contribute
hydronium or hydroxide ions to the solution directly. So, why is a sodium carbonate
solution basic while a sodium chloride solution is neutral? The reason must lie in the action
of aqueous carbonate ions, not sodium ions, because both salts release sodium ions in
solution.
The basic character of carbonate solutions can be explained by the Brønsted–Lowry
theory. As in the case of aqueous ammonia, a Brønsted–Lowry base acts as a proton
acceptor, and may remove a proton from water to form hydroxide ions in solution. The
following equation describes how the carbonate ion acts as a Brønsted–Lowry base.
H
2 H O
CO3(aq)
2 (l)
base
HCO OH(aq)
3(aq)
acid
Empirical studies have shown that the carbonate ion is a relatively weak base with a
base ionization constant, Kb, of 2.1 104.
Ammonium chloride, NH4Cl(s), is a soluble salt that forms an acidic solution when dissolved in water. The dissociation of the ions in aqueous solution is described by the
equation
NH4Cl(s) → NH4
(aq) Cl(aq)
The pH of an NaCl(aq)solution is neutral, so pH is unaffected by the presence of Cl(aq)
ions. Is it possible that the acidity of ammonium chloride is caused by the NH4(aq) ion?
We can hypothesize that a Brønsted-Lowry proton-transfer reaction occurs between
ammonium ions and water molecules:
NH4
(aq) H2O(l) e
H3O
(aq) NH3(aq)
Laboratory tests show that the ammonium ion is indeed a weak acid with an acid
ionization constant, Ka, of 5.8 1010.
In general, since salts contain two different ions, the pH of an aqueous salt solution
may be affected by the anion, the cation, or both.
NEL
Acid–Base Equilibrium 581
hydrolysis a reaction of an ion with
water to produce an acidic or basic
solution (hydronium or hydroxide
ions)
LEARNING
The reaction of an ion with water to produce an acidic or basic solution is called
hydrolysis (from the Greek hydro, meaning “water,” and lysis, meaning “splitting”).
In the case of sodium chloride, neither sodium nor chloride ions hydrolyze, so neither
ion affects the acid–base properties of an aqueous solution; the solution remains neutral. Conversely, when ammonium carbonate dissolves, carbonate ions hydrolyze to produce hydroxide ions (which may produce a basic solution) and ammonium ions hydrolyze
to produce hydronium ions (which may produce an acidic solution). You must always
consider both ions when assessing the effect of a salt on the pH of an aqueous solution.
So, can we accurately predict if a salt will produce an acidic or a basic solution? In the
following analysis, we will develop models that will help you make such predictions.
Salts That Form Neutral Solutions
TIP
Do not confuse the definition of
hydrolysis as it applies to
acid–base chemistry with the
meaning of the term presented
in Chapter 1. In acid–base
chemistry it means a reaction of
an ion with water to produce an
acidic or basic solution. In
Chapter 1 you used it to mean
the breaking of a covalent bond
using the elements of water.
In general, salts that consist of the cations of strong bases (like Na
(aq) and K (aq)) and
the anions of strong acids (like Cl(aq) and NO3(aq)) have no effect on the pH of an aqueous
solution. Examples include NaCl(aq), KCl(aq), NaI(aq), and NaNO3(aq) (Table 1).
Table 1 Composition of “Neutral” Salts
Salt
Cation from strong base
Anion from strong acid
NaCl
NaOH
HCl
KCl
KOH
HCl
NaI
NaOH
HI
NaNO3
NaOH
HNO3
Salts That Form Acidic Solutions
Table 3 Ka for Some Metal Ions
at SATP
Metal cation*
Ka
Zr 4
(aq)
Sn2
(aq)
Fe3
(aq)
Cr3
(aq)
Al3
(aq)
Be2
(aq)
Fe2
(aq)
Pb2
(aq)
Cu2
(aq)
2.1
Table 2 Composition and Hydrolysis of “Acidic” Salts
Salt
2.0 102
1.5 103
1.0 104
9.8 106
3.2 107
1.8 107
1.6 108
1.0 108
* The aqueous metal ion is a
hydrated complex ion (e.g.,
Cu(H2O)42
(aq)). Aqueous ions
of transition metals are usually written in a simplified
form, without showing the
number of water molecules
present in the actual hydrated
complex ion, as shown in
Table 3.
582 Chapter 8
Cations such as the ammonium ion, NH4
(aq), and hydrazinium ion, N2H5 (aq), act as
Brønsted–Lowry acids. They hydrolyze to form hydronium ions in solution. Consequently,
solutions of NH4Cl(aq) and N2H5Cl(aq) are acidic.
In general, cations that are conjugate acids of weak molecular bases act as weak acids
that have a tendency to lower the pH of a solution (Table 2).
Cation (acid) of weak base
Hydrolysis reaction
NH4Cl
NH4
(aq)
NH3
NH4
(aq) H2O(l) e H3O (aq) NH3(aq)
N2H5Br
N2H5
(aq)
N2H4
N2H5
(aq) H2O(l) e H3O(aq) N2H4(aq)
Certain metal cations also hydrolyze and act as acids. An aluminum nitrate solution,
Al(NO3)3(aq), is acidic. However, other solutions, such as sodium nitrate, NaNO3(aq),
and calcium nitrate, Ca(NO3)2(aq), are neutral. Why are some metal salts acidic and
some neutral? We could propose a hypothesis that the aluminum ion is responsible for
the acidity of the solution. Further tests with aluminum solutions indeed establish that
Al3 behaves as an acid in water (increases [H
(aq)]). The research indicates that Group
2
1 and 2 metal ions (except for Be ) do not produce acidic solutions, but that highly
charged small ions do form acidic solutions. (See Table 3 and Appendix C9.)
3
2
Ions such as Al3
(aq), Fe (aq), and Sn (aq) have large (positive) charge densities — a large
amount of charge in a small volume. These cations produce hydronium ions indirectly
by a slightly different reaction than the ammonium ion example above. When a highly
charged metal ion such as Al3 dissolves in water, it becomes hydrated with six water molecules (waters of hydration) according to the following equation:
3
Al3
(aq) 6 H2O(l) e Al(H2O)6(aq)
NEL
Section 8.3
This means that six water molecules bond to the ion with relatively weak electrostatic
3 ion increases the polarity of the
forces of attraction. The high-charge density of the Al(aq)
—OH bonds in the H2O molecules. That is, the H2O molecules in Al(H2O)63
(aq) are
more likely to transfer a proton to the solvent (H2O(l)) than are H2O(l) molecules in
pure water. Although we might imagine that several (or all) of the water molecules in
Al(H2O)63
(aq) could transfer a proton, experiments show that only one H2O will. Thus,
Al(H2O)63
(aq) behaves as a weak monoprotic acid in aqueous solution, according to the
following equation:
H(aq)
3 H O
Al(H2O)6(aq)
2 (l)
acid
base
Al(H O) (OH) 2
H3O(aq)
2 5
(aq)
conjugate
base
conjugate
acid
–
Al3+
+
Al3+
Al(H2O)63+(aq)
2+
Al(H2O)5(OH)(aq)
+
+
H+(aq)
Notice that this weak acid and its acid ionization constant appear in Appendix C9.
No cation with low charge density acts as an acid in this way. This includes all of the
singly charged ions of Group 1 metals, Li, Na, K, Rb, and Cs. Since they cannot
hydrolyze, they do not affect the pH of aqueous solutions. With the notable exception
of Be2, none of the Group 2 cations acts as an acid either.
Salts That Form Basic Solutions
A sodium chloride solution, NaCl(aq), is neutral, but a sodium acetate solution,
NaC2H3O2(aq), is basic. Why? Both solutions contain aqueous sodium ions that do not
hydrolyze in solution, so the change in pH must be due to the effect of the acetate ion,
. It is reasonable to suspect hydrolysis to be the cause. In this case, hydrolysis
C2H3O2(aq)
caused by the acetate ion produces hydroxide ions and increases the pH of the solution.
The following two equations show the dissociation of sodium acetate followed by the
hydrolysis of the resulting acetate ion:
NaC2H3O2(s) → Na
(aq) C2H3O2(aq)
C2H3O2
(aq) H2O(l) e
NEL
HC2H3O2(aq) OH
(aq)
(dissociation)
(hydrolysis)
Acid–Base Equilibrium 583
cannot hydrolyze. So, why is it that the
Earlier, we explained that Group 1 anions like Cl(aq)
acetate ion is able to hydrolyze, but the chloride ion cannot? Earlier in this chapter, you
learned that the stronger the acid, the weaker its conjugate base. Extremely strong acids
and NO .
such as HCl(aq) and HNO3(aq) possess extremely weak conjugate bases, Cl(aq)
3(aq)
These bases attract protons (H ions) so poorly that they are unable to hydrolyze to any significant degree. If they did attract protons strongly, HCl and HNO3 would not be strong acids
and would not ionize so readily in aqueous solution. In general, the conjugate bases of
, Br , I , NO–
strong acids (e.g., Cl(aq)
(aq) (aq)
3(aq)) do not affect the pH of an aqueous solution.
Conversely, the acetate ion, C2H3O2(aq), is the conjugate base of acetic acid, HC2H3O2(aq).
Acetic acid is a weak acid because the acetate ion portion of the molecule possesses a
relatively high affinity for the ionizable proton (H ion). This relatively strong attrac ions in aqueous
tion for protons makes the acetate ion able to hydrolyze to form OH(aq)
solution. A solution of sodium acetate is basic because the sodium ions cannot act as acids,
but the acetate ions do act as weak bases according to the hydrolysis equation above.
Remember that, in general, the conjugate base of a weak acid is also weak. The Ka of
acetic acid is 1.8 105 (weak), and the Kb of its conjugate base, the acetate ion, is
5.6 1010 (weak). However, when the acetate ion is introduced as the anion of a salt
containing a nonhydrolyzing cation like Na
(aq) or K(aq), the acetate ion forms the major
acid–base equilibrium in solution. (The only other reaction is the autoionization of
water.) This results in a net production of OH
(aq) and a basic solution.
When determining whether an anion will affect the pH of a solution, it helps to
remember that
•
the anion of a strong acid is too weak a base to affect the pH of an aqueous solu
tion (e.g., Cl (aq) and NO3(aq)), and
•
the anion of a weak acid, although it is also weak, is a strong enough base (when
acting alone) to affect the pH of an aqueous solution. It will tend to increase the
pH of the solution and make it more basic.
In general, a salt made of a nonhydrolyzing cation (like Na
(aq) or K(aq)), and an anion that
is the conjugate base of a weak acid, will form a basic aqueous solution (pH 7).
Salts That Act as Acids and Bases
Some salts contain the cation of a weak base and the anion of a weak acid — both ions
can hydrolyze. To obtain a precise estimate of the pH of such a solution, we would have
to solve two hydrolysis equilibria simultaneously. This is a difficult mathematical problem.
However, we can predict whether the solution will be acidic, basic, or (approximately)
neutral without performing any calculations.
Consider the ions produced when ammonium cyanide, NH4CN(s), dissolves in water.
NH4CN(s) → NH4
(aq) CN(aq)
(dissociation)
The cation, NH4
(aq), is the conjugate acid of the weak molecular base, NH3(aq), so it
will hydrolyze to produce hydronium ions in solution. Its tendency is to lower the pH.
The anion, CN
(aq), is the conjugate base of the weak acid, HCN(aq), so will hydrolyze to
produce hydroxide ions in solution. Its tendency is to increase the pH. What is the net
effect? That depends on the relative strengths of the two ions, as judged by the Ka value
for the acid and the Kb value for the base. If these values are equal, then the salt will
have no net effect on the pH of the solution. If the Ka of the acid is larger than the Kb of
the base, then the solution will be more acidic. Conversely, a higher Kb would result in
a basic solution. In the case of ammonium cyanide,
584 Chapter 8
NEL
Section 8.3
NH4
(aq) H2O(l) → NH3(aq) H3O (aq)
Ka 5.8 1010 mol/L
CN
(aq)
Kb 1.6 105 mol/L
H2O(l) → HCN(aq) OH
(aq)
Since the Kb of the cyanide ion (1.6 105 mol/L) is much larger than the Ka of the
ammonium ion (5.8 1010 mol/L), an aqueous solution of ammonium cyanide will
be basic.
SUMMARY
Predicting the Acid–Base
Behaviour of a Salt
The following steps will help you predict whether the ions of a salt have an effect on the
pH of an aqueous solution.
Step 1 Determine if the cation is the conjugate acid of a weak base or a cation
with high charge density. If so, it will make the solution more acidic. If
not, it will not affect the pH of the solution.
Step 2 Determine if the anion is the conjugate base of a weak acid. If so, it will
make the solution more basic. If not, it will not affect the pH of the
solution.
Step 3 If the salt has a cation and an anion that can both hydrolyze, compare
the Ka and Kb values of the cation and anion. If Ka Kb, then the solution will become more acidic. If Ka Kb, the solution will become
more basic. If Ka Kb, the solution will be neutral.
Predicting the Acidic or Basic Nature of Solutions
SAMPLE problem
(a) Predict whether a 0.10 mol/L solution of NaNO2(aq) will be acidic, basic, or
neutral.
(b) Calculate the pH of a 0.10 mol/L solution of NaNO2(aq) .
(a) First, write the dissociation equation for NaNO2(s).
NaNO2(s) → Na
(aq) NO2(aq)
Next, examine the cation. Na
(aq) is a Group 1 metal cation so will have no effect on pH.
Now, examine the anion, NO2
(aq). It is the conjugate base of the weak acid, HNO2(aq),
so will act as a base and increase the pH.
A 0.1 mol/L solution of NaNO2(aq) will be basic.
(b) Since Na
(aq) cannot hydrolyze, the pH of a 0.10 mol/L NaNO2(aq) solution is determined solely by the reaction of NO2
(aq) with water.
Write the chemical equation for the hydrolysis reaction between the base, NO2
(aq),
and water, then write the corresponding Kb expression.
NO2
(aq) H2O(l) e OH(aq) HNO2(aq)
[OH
(aq)][HNO2(aq)]
Kb
[NO2
(aq)]
NEL
Acid–Base Equilibrium 585
Calculate the value of Kb for NO2
(aq) from the value of Ka for HNO2(aq) (which you can find
in Appendix C9).
KaKb Kw
Kb
Kw
Ka
1.0 1014
7.2 104
Kb
1.4 1011
Now, construct an ICE table to show the changes in concentration that occur as the
reaction reaches equilibrium. Note that, in the reaction, one mole of HNO2(aq) and one
mole of OH
(aq) are produced for every mole of NO2(aq) that reacts. Let x represent the
changes in concentration that occur as the reaction establishes equilibrium.
Table 4 ICE Table for the Hydrolysis of NO2
(aq)
NO2
(aq) H2O(l) e OH(aq) HNO2(aq)
Initial concentration (mol/L)
0.10
0
0
Changes in concentration (mol/L)
x
x
x
Equilibrium concentration (mol/L)
0.10 x
x
x
[OH
(aq)][HNO2(aq)]
Kb
[NO2
(aq)]
[OH
(aq)] [HNO2(aq)] x
[NO2
(aq)] 0.10 x 0.10
(Since the value of Kb is so small, we will make the assumption that (0.10 x ) 0.10.
This simplification is warranted by the hundred rule. Check it yourself.)
x2
1.4 1011
0.10
x 1.4 1012
x 1.2 106
As usual, we now justify our simplifying assumption, using the 5% rule.
1.2 106
100% 1.2 103 %
0.10
Since 1.2 103 % 5%, our assumption is justified. Therefore,
x 1.2 106
[OH
(aq)]
1.2 106 mol/L
Now, we can calculate pOH and pH.
pOH log [OH
(aq)]
log(1.2 106 mol/L)
pOH 5.92
pH pOH 14.00
14.00 pOH
14.00 5.92
pH 8.08
The pH of a 0.10 mol/L NaNO2(aq) solution is 8.08.
586 Chapter 8
NEL
Section 8.3
Example
(a) Predict whether a 0.20 mol/L solution of ammonium chloride, NH4Cl(aq), will be acidic,
basic, or neutral.
(b) Calculate the pH of a 0.20 mol/L solution of NH4Cl(aq).
Solution
(a) First, write the dissociation equation for NH4Cl(aq).
NH4Cl(aq) → NH4
(aq) Cl(aq)
Since Cl
(aq) is the conjugate base of a strong acid, it will not affect the pH of the solution. NH4
(aq) is the conjugate acid of the weak base, NH3(aq), so it will hydrolyze
according to the following equation:
NH4
(aq) H2O(l) e H3O (aq) NH3(aq)
A solution of NH4Cl(aq) will be acidic.
(b) NH4
(aq) H2O(l) e H3O (aq) NH3(aq)
Ka 5.8 1010 (from Appendix C9)
[H3O
(aq)][NH3(aq)]
Ka
[NH4
(aq)]
Table 5 ICE Table for the Hydrolysis of NH4
(aq)
NH4
(aq) H2O(l) e H3O(aq) NH3(aq)
Initial concentration (mol/L)
0.20
0
0
Changes in concentration (mol/L)
x
x
x
x
x
Equilibrium concentration (mol/L) 0.20 x
[H3O(aq)][NH3(aq)]
Ka
[NH4(aq)]
x2
5.8 1010
0.20 x
Predicting whether 0.20 x 0.20...
[HA]initial
0.20
Ka
5.8 1010
3.4 108
Since 3.4 108 100, we assume that 0.20 x 0.20.
x2
5.8 1010
0.20
x 2 2.9 109
x 5.4 105
Since x [H3O
(aq)],
5
[H3O
(aq)] 5.4 10
pH log [H3O
(aq)]
log(5.4 105)
pH 4.
The pH of a 0.20 mol/L NH4Cl(aq) solution is 4.
NEL
Acid–Base Equilibrium 587
Practice
Understanding Concepts
Answers
3. 4.88
4. 4.92
1. Predict whether the following solutions are acidic, basic, or neutral. Provide expla-
nations to support your predictions.
(a) ammonium phosphate, (NH4)3PO4(aq) (fertilizer)
(b) ammonium sulfate, (NH4)2SO4(aq) (fertilizer)
(c) magnesium oxide, MgO(aq) (milk of magnesia)
2. Predict whether a solution of sodium sulfite, Na2SO3(aq) (photographic developer)
will be acidic, basic, or neutral.
3. Calculate the pH of a 0.30 mol/L ammonium nitrate (fertilizer) solution.
4. Calculate the pH of 0.25 mol/L NH4Br(aq).
5. Predict whether a solution of NH4C2H3O2(aq) is acidic, basic, or neutral. Explain.
Making Connections
6. What kind of fertilizers would be appropriate for acid-loving plants like evergreens?
Hydrolysis of Amphoteric Ions
Remember that all polyatomic ions whose chemical formulas begin with hydrogen, H (e.g.,
HCO3–(aq) and HSO4–(aq)) are amphoteric. As mentioned in Section 8.1, the term
amphiprotic may also be used to describe such entities because they can either donate
or accept a hydrogen ion (proton).
NaHCO3(s)dissolves in water, forming a conducting solution containing sodium and
hydrogen carbonate ions.
NaHCO3(s) → Na
(aq) HCO3 (aq)
– ion may hydrolyze as an acid or a base, according
Because it is amphoteric, the HCO3(aq)
to the following equilibrium equations:
2
+
HCO3
(aq) H2O(l) e H3O (aq) CO3(aq)
(acid hydrolysis)
HCO3
(aq) H2O(l) e OH (aq) H2CO3(aq)
(base hydrolysis)
Testing a sodium hydrogen carbonate solution with litmus paper reveals that it is
basic. Therefore, we assume that the base hydrolysis equilibrium dominates in solution.
Similarly, the acidic character of a sodium hydrogen sulfate, NaHSO4(s) solution is
explained by the dissociation and subsequent acid hydrolysis of the hydrogen sulfate ion.
NaHSO4(s) → Na
(aq) HSO4 (aq)
(dissociation)
2
+
HSO4
(aq) H2O(l) e H3O (aq) SO4(aq)
(acid hydrolysis )
In both cases, the hydrolyzing ion may act as an acid or a base, but one equilibrium
predominates and gives rise to the overall acidic or basic character of the aqueous solution. How do we predict which one will win out? The following discussion describes
the theory.
Picture a laboratory setting in which a student tests the pH of a sodium dihydrogen
borate, NaH2BO3(aq), solution. The solution has a pH greater than 7. What reaction
occurs to make the solution basic?
588 Chapter 8
NEL
Section 8.3
The following equations describe the dissociation of NaH2BO3(s) in water, and the
hydrolysis of the H2BO3
(aq) ions in solution.
NaH2BO3(s) → Na
(aq) H2BO3 (aq)
H2BO3
(aq)
H2BO3
(aq)
H2O(l) e
H2O(l) e
(dissociation)
2
H3O
(aq) HBO3(aq)
OH(aq) H3BO3(aq)
(acid hydrolysis)
(base hydrolysis)
Since the solution is basic when tested with an acid–base indicator, H2BO3
(aq) must
hydrolyze primarily as a base. Is there a way to predict the acid–base character of a solution without testing it directly in the laboratory? Remember that when a salt containing
both a hydrolyzing anion and a hydrolyzing cation, such as NH4CN(aq), is dissolved in
water, we predict the acid–base characteristics of the solution by comparing Ka and Kb
values. The following sample problem will provide you with a similar model for
predicting hydrolysis in solutions of amphoteric ions.
Predicting the Acidity of Amphoteric Ions
SAMPLE problem
Predict whether an aqueous solution of baking soda, NaHCO3(s), is acidic, basic,
or neutral.
First, write the dissociation equation.
NaHCO3(s) → Na
(aq) HCO3 (aq)
(dissociation)
Next, find the acid ion dissociation constant (from a reference table), as if the hydrogen
carbonate ion were to act as an acid:
2
HCO3
(aq) H2O(l) e CO3 (aq) H3O(aq)
Ka 4.7 1011 (hydrolysis as an acid)
Next, find the base ion dissociation constant, as if the hydrogen carbonate ion were to
act as a base (use either a table of K b values, or calculate from K a):
HCO3
(aq) H2O(l) e H2CO3(aq) OH (aq)
Kb 2.7 10–8 (hydrolysis as a base)
According to the relative values of Ka and Kb, the baking soda solution should be basic
because the Kb is larger than the Ka. The hydrogen carbonate ion is a stronger base than it
is an acid.
Example
Predict whether a solution of NaHBO3(aq) is acidic, basic, or neutral.
Solution
NaH2BO3(s) → Na
(aq) H2BO3(aq)
(dissociation)
H O e HBO 2 H O
H2BO3(aq)
2 (l)
3(aq)
3 (aq)
Ka 5.8 1010 (hydrolysis as an acid)
H2BO3
(aq) H2O(l) e H3BO3(aq) OH(aq)
Kb 1.7 105 (hydrolysis as a base)
Since Kb Ka, a solution of NaH2BO3(aq) is basic.
Practice
Understanding Concepts
7. Make a list of all amphoteric ions from your acid–base table, Appendix C9.
8. Predict whether the following solutions are acidic, basic, or neutral.
(a) NaHSO4(aq)
NEL
(b) Na2HPO4(aq)
Acid–Base Equilibrium 589
Hydrolysis of Metal and Nonmetal Oxides
When calcium oxide, CaO(s), is dissolved in water it produces a basic solution. However,
an aqueous carbon dioxide solution, CO2(aq), is acidic. Can we explain this evidence
using the hydrolysis theory? First, we must realize that calcium oxide and carbon dioxide
both have low solubility in water. Therefore, we show the pure substance reacting with
water rather than dissolving in water.
CaO(s) H2O(l) → Ca2
(aq) 2 OH(aq)
As you know, most metal oxides have low solubility in water, but the accepted theory
is that the solid state oxide ions are converted quantitatively into aqueous hydroxide
ions by the reaction with water to form a basic solution, according to the following equation.
O2
(s) H2O(l) → 2 OH(aq)
The O2
(s) ions do not exist in aqueous solution; they are quantitatively converted to
OH (aq) ions from the solid state (in CaO(s)). However, the equation may be used to
explain the basic nature of the solutions of metal oxides.
Now, let us try to explain the acidic character of carbon dioxide in water. A possible
explanation is the two-step process presented below.
CO2(g) H2O(l) e H
2CO3(aq)
H
2CO3(aq) H2O(l) e H3O (aq) HCO3(aq)
CO2(g) 2 H2O(l) e H3O
(aq) HCO3(aq)
(net equation)
Chemists have done numerous tests on metal oxides and nonmetal oxides to determine
the acidic and basic character of the solutions formed. Their evidence led to the following generalizations.
• Metal oxides react with water to produce basic solutions.
• Nonmetal oxides react with water to produce acidic solutions.
SAMPLE problem
Predicting the Acidity of Solutions of Oxides
Predict, using the hydrolysis concept, whether solutions of the following oxides
will be acidic, neutral, or basic. Write an appropriate hydrolysis equation in each
case.
(a) magnesium oxide, MgO(s)
(b) sulfur dioxide, SO2(g)
(a) Magnesium oxide, MgO(s), is a metal oxide, so will react with water to form a basic
solution, according to the following hydrolysis equation:
MgO(s) H2O(l) e Mg2
(aq) 2 OH(aq)
The hydroxide ions are produced according to the following equation:
2 H O
O(s)
2 (l) e 2 OH (aq)
(b) Sulfur dioxide, SO2(g), is a nonmetal oxide, so will react with water to form an acidic
solution, according to the following hydrolysis equation:
SO2(g) 2 H2O(l) e H3O
(aq) HSO 3(aq)
590 Chapter 8
NEL
Section 8.3
Example
Use hydrolysis concepts to predict whether an aqueous solution of copper(II) oxide,
CuO(s), is acidic, basic, or neutral.
Solution
Copper (II) oxide, CuO(s), is a metal oxide, so will react with water to form a basic
solution, according to the following equation:
CuO(s) H2O(l) e Cu2
(aq) 2 OH(aq)
2 H O
O(s)
2 (l) e 2 OH (aq)
The acidic properties of nonmetal oxides are responsible for many natural processes
such as the the weathering of minerals, the absorption of nutrients by the roots of plants,
and the chemistry of tooth decay. For example, limestone, CaCO3(s), dissolves in water
that is made acidic by the dissolution of CO2(g) from the atmosphere, according to the
following two-step process:
CO2(g) H2O(l) e H+(aq) HCO3
(aq)
(acidification of rain or ground water)
H
(aq)
(dissolution of CaCO3(s) in acidic solution)
CaCO3(s) e
Ca2+
(aq)
HCO3
(aq)
Limestone caves and the stalagmites and stalactites they contain are formed by the
action of acidic ground water on limestone deposits in Earth’s crust.
A cave is formed when the above reactions proceed to the right and calcium carbonate
in underground limestone deposits is dissolved. Stalactites are formed in the cave when
the aqueous solution on the ceiling evaporates. As it evaporates, carbon dioxide escapes.
This shifts the equilibrium to the left, causing solid calcium carbonate to precipitate.
This precipitation causes stalactites to form on the ceilings and stalagmites to form
where drops of the solution drip onto the cave floor. Stalactite and stalagmite formation
is very slow; the growth rate is, on average, about one millimetre per century.
Tooth decay is also caused by the dissolution of minerals in acidic solutions. Tooth
enamel is composed of the mineral hydroxyapatite, Ca5(PO4)3OH (Ksp 6.8 1037).
Acids (in saliva, from fruits and fruit juices, or formed when sugars are metabolized by
bacteria in the mouth) react with hydroxyapatite, leading to erosion of tooth enamel and
eventually tooth decay (cavities). Fluoride salts (as a source of fluoride ions) are often
added to toothpaste and to drinking water in treatment plants in some communities to
help prevent tooth decay. The fluoride ions react with the Ca5(PO4)3OH in tooth enamel
to form more decay-resistant fluorapatite Ca5(PO4)3F (Ksp 1.0 1060).
DID YOU
KNOW
?
Growing Up or Growing Down?
When viewing photos of stalactites
and stalagmites in a cave, you may
not be able to tell which way is
up. Look at the tips of the
formations. Stalactites almost
always have pointed tips, whereas
stalagmites are usually rounded or
flat. Is the photo right-side up or
upside down?
Practice
Understanding Concepts
9. For each of the following solutions of compounds, write an ionization or dissociation
equation where appropriate, and then write a net equation showing reactions with
water to produce either hydronium or hydroxide ions (consistent with the evidence).
(a) Na2O(s) in solution turns red litmus blue.
(b) SO3(g) in solution turns blue litmus red.
NEL
Acid–Base Equilibrium 591
Making Connections
10. Limestone caves are very popular tourist attractions. Some of the most popular caves
attract up to 500 000 people per year. Many speleologists (cave experts) believe that
cave formations such as stalactites and stalagmites are in danger of dissolving away
on account of the large numbers of visitors.
(a) Explain how the presence of large numbers of humans may affect the chemical
stability of stalagmites and stalactites in caves.
(b) Suggest possible solutions that do not include barring people from the caves.
11. Many communities are against the addition of fluoride to municipal drinking water.
Conduct library and Internet research to learn more about this issue. List arguments
for and against this practice and write a brief position paper.
GO
SUMMARY
www.science.nelson.com
Acid–Base Characteristics of Salts
Table 6
Type of Salt
Examples
Description
pH
cation of strong base and
anion of strong acid
NaCl(aq), KNO3(aq)
NaI(aq)
does not hydrolyze
as an acid or as a base
neutral
cation of strong base and
anion of weak acid
NaC2H3O2(aq)
KF(aq)
anion hydrolyzes as a base;
cation does not hydrolyze
basic
cation is conjugate acid
of a weak base;
anion of a strong acid
NH4NO3(aq)
cation hydrolyzes as an acid;
anion does not hydrolyze
acidic
cation is conjugate acid
of a weak base; anion is
conjugate base of a weak
acid
NH4C2H3O2(aq)
NH4F(aq)
cation hydrolyzes as an acid;
anion hydrolyzes as a base
acidic if Ka Kb
basic if Ka Kb
neutral if Ka Kb
cation is highly
charged metal ion;
anion of a strong acid
AlCl3(aq)
FeI3(aq)
hydrated cation hydrolyzes
as an acid;
anion does not hydrolyze
acidic
metal oxides
CuO(s)
solid state oxide ion reacts with
water to form OH(aq)
basic
nonmetal oxides
CO2(g)
compound reacts with water
to form H3O
(aq)
acidic
NH4Cl(aq)
The Lewis Model of Acids and Bases
Lewis acid an electron-pair
acceptor
Lewis base an electron-pair donor
592 Chapter 8
The Arrhenius and Brønsted–Lowry models successfully explain much of the behaviour of acids and bases. Nevertheless, both of these models contain limitations. Remember
that the Arrhenius model could not adequately explain the basic properties of an aqueous
ammonia solution. In the early 1920s, G. N. Lewis expanded the Brønsted–Lowry model
to encompass a number of substances that would not normally be classified as
Brønsted–Lowry acids or bases. According to the Lewis model, a Lewis acid is an electron-pair acceptor and a Lewis base is an electron-pair donor. In order to act as a Lewis
base, a substance must possess a non-bonded pair of electrons in one of its orbitals.
Conversely, in order to act as a Lewis acid, a substance must possess an empty valence
orbital that may accept (share) a pair of non-bonding valence electrons from a Lewis base.
The following structural formula equation illustrates the reaction between a Lewis acid,
H
(aq), and a Lewis base, H2O(l), to form H3O (aq).
NEL
Section 8.3
+
H
H+ ion
(Lewis acid)
H
H
O
H
+
O
H
H
water
(Lewis base)
hydronium ion
In the above reaction, the H
(aq) ion (proton) acts as the Lewis acid (electron pair
acceptor) and the water molecule acts as a Lewis base (electron-pair donor).
In this case, the Lewis base is also a Brønsted–Lowry base, and the Lewis acid is also
a Brønsted–Lowry acid. However, in the following example, this is not the case.
F
F
H
B
F
N
H
F
B N
F
H
boron trifluoride
(Lewis acid)
H
F
ammonia
(Lewis base)
H
H
boron trifluoride
ammonia complex
Note that the Brønsted–Lowry model also accounts for ammonia as a base, but does
not characterize boron trifluoride as an acid.
The Lewis acid–base theory explains the reaction between BF3 (boron trifluoride)
and NH3 (ammonia). Boron trifluoride is a trigonal planar molecule with sp2 hybrid
orbitals. This arrangement leaves an empty 2p orbital on the boron atom that is able to
accommodate the pair of nonbonding electrons in the sp3 hybrid orbital of NH3. A
covalent bond forms between the boron and the nitrogen, forming the compound
BF3NH3.
The Lewis acid–base theory can also explain why small, highly positive ions such as
3
Al form complex ions in water:
3
Al3
(aq) + 6 H2O(l) e Al(H2O)6(aq)
This is a Lewis acid–base reaction. The water molecules each possess nonbonding
3 ion possesses empty 3s, 3p, and
pairs of electrons and so act as Lewis bases, and the Al(aq)
3d orbitals that may accommodate electron pairs. The electron configuration of the
Al3 ion can be represented as Al3: [Ne] 3s 0 3p 0. The complex Al(H2O)63
(aq) is formed
when an Al3 ion, acting as a Lewis acid, bonds with six water molecules, each acting as
a Lewis base.
The Lewis model is a more general model of acids and bases that not only encompasses the Brønsted–Lowry model, but extends it. Brønsted–Lowry acids and bases are
acids and bases in the Lewis model, but the reverse is not always true.
Example
Identify the Lewis acid and the Lewis base in the following reaction.
SO3(aq) H2O(l) e H2SO4(aq)
O
S
O
O
sulfur trioxide
(Lewis acid)
NEL
H
O
O
H
water
(Lewis base)
O
H
S
O
H
O
sulfuric acid
Acid–Base Equilibrium 593
Solution
SO3(aq) is the Lewis acid and the water is the Lewis base.
Practice
Understanding Concepts
12. Identify the Lewis acid and the Lewis base in each of the following reactions.
(a) H
(aq) OH(aq) → H2O(l)
(b) H(aq) NH3(aq) e NH4
(aq)
Section 8.3 Questions
Understanding Concepts
1. Predict whether the following solutions are acidic, basic, or
neutral. Provide explanations to support your predictions.
(a) table salt, NaCl(aq) (saline or brine) solution
(b) aluminum chloride, AlCl3(aq) (antiperspirant)
(c) Na2CO3(aq) (washing soda)
2. Predict whether a solution of ammonium carbonate (a com-
ponent of baking powder) will be acidic, basic, or neutral.
3. What is the strongest possible acid in an aqueous solution,
and what is the strongest possible base in an aqueous
solution?
4. Will an aqueous solution of BeCl2(aq) turn litmus red or
blue? Explain.
5. Predict whether the following solutions are acidic, basic, or
neutral.
(a) a carbonated beverage containing CO2(aq) (pop)
(b) strontium oxide, SrO(s)
Applying Inquiry Skills
6. Analyze the Evidence (Table 7) to determine the kind of
solutions (acidic, basic, or neutral) formed when Period 3
oxides are placed in water.
Experimental Design
acid) and a strong base (sodium hydroxide) are added to
determine if a neutralization reaction occurs.
Evidence
Table 7 Litmus and Neutralization Tests on Oxides
Oxide
Litmus test
HCl(aq) test
NaOH(aq) test
Na2O(s)
red to blue
neutralizes
no reaction
MgO(s)
red to blue
neutralizes
no reaction
As2O3(s)
(insoluble)
neutralizes
neutralizes
SiO2(s)
(insoluble)
no reaction
neutralizes
P2O3(s)
blue to red
no reaction
neutralizes
SO3(g)
blue to red
no reaction
neutralizes
Cl2O(g)
blue to red
no reaction
neutralizes
Extension
7. Nitrogen oxides such as nitrogen dioxide, NO2(g), are pro-
duced in automobile engines and released into the atmosphere via exhaust fumes. In the atmosphere, nitrogen
dioxide will dissolve in droplets of rain. Use hydrolysis concepts to predict what will happen to the pH of rain when
NO2(g) dissolves.
Oxides of elements in Period 3 are tested in water, using
litmus paper. To all the oxides, a strong acid (hydrochloric
594 Chapter 8
NEL