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5 Acid–Base Properties of Elements and Their Oxides
760 Chapter 16 | Acids and Bases, A Molecular Look
Acidity of Hydrated Metal Ions
n The force of attraction between
an ion and water molecules is often
strong enough to persist when the
solvent water evaporates. Solid
hydrates crystallize from solution, for
example, CuSO4 ⋅ 5H2O (page 14).
When an ionic compound dissolves in water, molecules of the solute gather around the
ions and we say the ions are hydrated. Within a hydrated cation, the metal ion behaves as
a Lewis acid, binding to the partial negative charges on the oxygens of the surrounding
water molecules, which serve as Lewis bases. Hydrated metal ions themselves tend to be
Brønsted acids because of the equilibrium shown below. For simplicity, the equation represents a metal ion as a monohydrate—namely, M(H2O)n+ with a net positive charge of
n+ (n being 1, 2, or 3, depending on M ).
M(H2O)n+ + H2O �
� MOH(n-1)+ + H3O+
Figure 16.3 | Polarization of a
water molecule by a metal
cation. The positive charge on
the metal ion pulls electron density
away from the H atoms of the
water molecule. This increases the
partial positive charge on the H
atoms and makes them easier to
transfer to another H2O molecule
In other words, hydrated metal ions tend to be proton donors in water. Let’s see why.
The positive charge on the metal ion attracts the water molecule and draws electron
density from the O!H bonds, causing them to become more polar (Figure 16.3). This
increases the partial positive charge on H and weakens the O!H bond with respect to the
transfer of H+ to another nearby water molecule in the formation of a hydronium ion.
The entire process can be illustrated as follows.
Electron density is reduced in the O 9 H bonds of water (indicated by
the curved arrows) by the positive charge of the metal ion, thereby
increasing the partial positive charge on the hydrogens. This promotes
the transfer of H+ to a water molecule.
The degree to which metal ions produce acidic solutions depends primarily on two
things: the amount of charge on the metal ion and the metal ion’s size. As the metal ion’s
charge increases, the polarizing effect is increased, which thereby favors the release of H+.
This means that highly charged metal ions ought to produce solutions that are more acidic
than ions of lower charge, and that is generally the case.
The size of the cation also affects its acidity because when the cation is small, the positive charge is highly concentrated. A highly concentrated positive charge is better able to
pull electrons from an O ! H bond than a positive charge that is more spread out. Therefore,
for a given positive charge, the smaller the cation, the more acidic are its solutions.
Both size and amount of charge can be considered simultaneously by referring to a
metal ion’s positive charge density, the ratio of the positive charge to the volume of the
cation (its ionic volume).
Charge density =
Acidity of metal cations
The higher the positive charge density, the more effective the metal ion is at drawing electron density from the O ! H bond and the more acidic is the hydrated cation.
Very small cations with large positive charges have large positive charge densities and
tend to be quite acidic. An example is the hydrated aluminum ion, Al3+. The hexahydrate,
Al(H2O)63+, is one of several of this cation’s hydrated forms that are present in an aqueous
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16.5 | Acid–Base Properties of Elements and Their Oxides
Figure 16.4 | Hydrated aluminum ion is a weak
proton donor in water. (a) The highly charged Al3+
ion (blue) binds six water molecules in an aqueous
solution to give an octahedral structure with the
formula Al(H2O)63+. (b) When it donates a proton,
the ion loses one unit of positive charge as a water
molecule is transformed into a hydroxide ion,
shown at the top of the structure. The resulting ion
now has the formula Al(H2O)5(OH)2+.
solution of an aluminum salt (see Figure 16.4). The ion is acidic in water because of the
Al(H2O)63+(aq) + H2O �
� Al(H2O)5(OH)2+(aq) + H3O+(aq)
The equilibrium, while not actually strongly favoring the products, does produce enough
hydronium ion so that a 0.1 M solution of AlCl3 in water has about the same concentration of hydronium ions as a 0.1 M solution of acetic acid, roughly 1 × 10-3 M.
Periodic Trends in the Acidity of Metal Ions
Within the periodic table, atomic size increases down a group and decreases from left to
right in a period. Cation sizes follow these same trends, so within a given group, the cation
of the metal at the top of the group has the smallest volume and the largest charge density.
Therefore, hydrated metal ions at the top of a group in the periodic table are the most
acidic within the group.
The cations of the Group 1A metals (Li+, Na+, K+, Rb+, and Cs+), with charges of just
1+, have little tendency to increase the H3O+ concentration in an aqueous solution.
Within Group 2A, the Be2+ cation is very small and has sufficient charge density to
cause the hydrated ion to be a weak acid. The other cations of Group 2A (Mg2+, Ca2+,
Sr2+, Ba2+), have charge densities that become progressively smaller as we go down the
group. Although their hydrated ions all generate some hydronium ion in water, the
amount is negligible.
Some transition metal ions are also acidic, especially those with charges of 3+. For
example, solutions containing salts of Fe3+ and Cr3+ tend to be acidic because their ions
in solution exist as Fe(H2O)63+ and Cr(H2O)63+, respectively, and undergo the same ionization reaction as does the Al(H2O)63+ ion discussed previously.
Influence of Oxidation Number
on the Acidity of Metal Oxides
Not all metal oxides are basic. As the oxidation number (or charge) on a metal ion increases,
the metal ion becomes more acidic ; it becomes a better electron pair acceptor. For metal
hydrates, we’ve seen that this causes electron density to be pulled from the OH bonds of
water molecules, causing the hydrate itself to become a weak proton donor. The increasing
acidity of metal ions with increasing charge also affects the basicity of their oxides.
When the positive charge on a metal is small, the oxide tends to be basic, as we’ve seen
for oxides such as Na2O and CaO. With ions having a 3+ charge, the oxides are less basic
and begin to take on acidic properties as well; they become amphoteric. (Recall from
Section 16.1 that a substance is amphoteric if it is capable of reacting as either an acid or a
base.) Aluminum oxide is an example; it can react with both acids and bases. It has basic
properties when it dissolves in acid.
Al2O3(s) + 6H+(aq) → 2Al3+(aq) + 3H2O
As noted earlier, the hydrated aluminum ion has six water molecules surrounding it, so in
an acidic solution the aluminum exists primarily as Al(H2O)63+.
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762 Chapter 16 | Acids and Bases, A Molecular Look
Aluminum oxide exhibits acidic properties when it dissolves in a base. One way to
write the equation for the reaction is
Al2O3(s) + 2OH-(aq) → 2AlO2-(aq) + H2O
Actually, in basic solution the formula of the aluminum-containing species is more complex than this and is better approximated by Al(H2O)2(OH)4-. Note that the difference
between the two formulas is just the number of water molecules involved in the formation
of the ion.
Al(H2O)2(OH)4- is equivalent to AlO2- + 4H2O
However we write the formula for the aluminum-containing ion, it is an anion, not a cation.
When the metal is in a very high oxidation state, the oxide becomes acidic.
Chromium(VI) oxide, CrO3, is an example. When dissolved in water, the resulting solution is quite acidic and is called chromic acid. One of the principal species in the solution
is H2CrO4, which is a strong acid that is more than 95% ionized. The acid forms salts
containing the chromate ion, CrO42-.
16.6 | Advanced Ceramics
and Acid–Base Chemistry
Ceramic materials have a long history, dating to prehistoric times. Examples of pottery
about 13,000 years old have been found in several parts of the world. Today, manufactured ceramics include common inorganic building materials such as brick, cement, and
glass. We find ceramics around the home as porcelain dinnerware, tiles, sinks, toilets, and
artistic pottery and figurines. These materials are made from inorganic minerals such as
clay, silica (sand), and other silicates (compounds containing anions composed of silicon
and oxygen), which are taken from the earth’s crust.
In recent times, an entirely new set of materials, generally referred to as advanced ceramics,
have been prepared by chemists in laboratories and have high-tech applications. We find
such ceramics in places we wouldn’t expect, such as in cell phones and inside diesel engines.
Several methods are used to form ceramics from their raw materials. In many cases, the
components of the ceramic are first pulverized to give a very fine powder. The powder is
then mixed with water to form a slurry that can be poured into a mold where it sets up
into a solid form that has little structural strength. (Alternatively, the fine powder is mixed
with a binder and pressed into the desired shape.) The newly formed object is next placed
in a kiln where it is heated to a high temperature, often over 1000 °C. At these high temperatures the fine particles stick together by a process called sintering, which produces the
finished ceramic. Although sintering is sometimes accompanied by partial melting of the
fine particles, often the particles remain entirely solid during the process.
There are some problems with making ceramics by the traditional methods just
described. Because it is difficult to produce uniform and very small particle sizes by grinding, ceramics made from materials prepared this way often contain small cracks and voids,
which adversely affect physical properties such as strength. In addition, the chemical composition of the ceramic cannot be easily and reproducibly controlled by mixing various
The Sol-Gel Process
For some types of ceramics, the problems of particle size and uniformity can be avoided
by using a method called the sol-gel process. The chemistry involved is based on acid–base
reactions similar to those discussed earlier in this chapter. The starting materials are metal
salts or compounds in which a metal or metalloid (e.g., Si) is bonded to some number of
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16.6 | Advanced Ceramics and Acid–Base Chemistry
Chemistry Outside the ClassrOOm
Applications of Advanced
There are some properties that all ceramics have in common, and
there are others that can be tailored by controlling the ceramic’s
composition and method of preparation. For example, virtually
all ceramic materials have extremely high melting points and are
much harder than steel. Some, such as boron nitride, BN, are
nearly as hard as diamond, the hardest substance known. Silicon
carbide, SiC, is a hard and relatively inexpensive ceramic to make
in bulk and has long been used as an abrasive in sandpaper and
The high strength and relatively low density of ceramic materials make them useful for applications in space. They are used
to line the surfaces of rocket engines because they transmit heat
very slowly and are able to protect the metals used to construct
the engines from melting. The exterior of the space shuttle
(Figure 1) is coated with protective ceramic tiles that serve as
thermal “heat” shields during re-entry from space. The tiles are
made of a low-density porous silica ceramic derived from common sand, SiO2. They are able to withstand very high temperatures and have a very low thermal conductivity, preventing heat
transfer to the body of the shuttle.
High-tech ceramics are relatively new materials, and new applications for them are continually being found. It is interesting to
note how varied their uses are. Here are a few examples.
Thin ceramic films, deposited by the sol-gel process or other
methods, are found on optical surfaces where they serve as antireflective coatings and as filters for lighting. Tools such as drill
bits are given thin coatings of titanium nitride (TiN) to make them
more wear resistant (Figure 2).
Figure 1 Space shuttle heat shield. Individual heat-shielding ceramic
tiles are visible on the surface of the space shuttle Challenger. The
photo was taken shortly after the spacecraft was completed. In 1986,
the Challenger was lost in a tragic accident shortly after lift-off. Damage
to the tiles during liftoff is believed to be responsible for the loss of the
Columbia and its crew in February 2003 when the spacecraft broke apart
during reentry into the earth’s atmosphere. (Roger Ressmeyer/© Corbis)
Zirconia (ZrO2) is used to
make ceramic golf spikes for
golf shoes, as well as portions
of hip-joint replacement parts
for medicine and knives that
stay sharp much longer than
steel. All these applications
are possible because of the
hardness of the ZrO2 ceramic.
Boron nitride (BN) powder
is composed of flat, plate-like
crystals that can easily slide
over one another, and one
of its uses is in cosmetics,
where it gives a silky texture
to the product (Figure 3).
Another boron-containing cerFigure 2 Titanium nitride. A drill
amic, boron carbide, is used bit with a thin golden coating of
along with Kevlar polymer to titanium nitride, TiN, will retain its
make bulletproof vests. When sharpness longer than steel drill
a bullet strikes the vest, it bits. (Andy Washnik)
is shattered by the ceramic,
which absorbs most of the kinetic energy, with residual energy
being absorbed by the Kevlar backing.
Silicon nitride, Si3N4, is used to make engine components for
diesel engines because it is extremely hard and wear resistant,
has a high stiffness with low density, and can withstand extremely
high temperatures and harsh chemical environments.
Piezoelectric ceramics have the property that they produce an
electric potential when their shape is deformed. Conversely, they
deform when an electric potential is applied to them. A company
makes “smart skis” that incorporate piezoelectric devices that
use both properties. Vibrations in the skis are detected by the
potential developed when they deform. A potential is then applied
to cancel the vibration. This “smart materials” technology was
first developed to dampen vibrations in optical components of
the “Star Wars” weapons system designed to protect the United
States from missile attack.
Figure 3 Ceramics in cosmetics. A variety of cosmetics use
boron nitride, BN, as an ingredient. (Andy Washnik)
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764 Chapter 16 | Acids and Bases, A Molecular Look
alkoxide groups. An alkoxide is an anion formed by removing a hydrogen ion from an
alcohol.6 For instance, removal of H+ from ethanol yields the ethoxide ion.
→ C 2 H 5O −
The ethoxide ion forms salts with metal ions that generally are soluble in ethanol, which
should not be surprising considering the similarities between the solute and solvent.
Alcohols are extremely weak acids with very little tendency to lose H+ ions, so alkoxide
ions are very strong bases with a very strong affinity for H+. When placed in water they
react immediately and completely by removing H+ from H2O to give the alcohol and
C 2H5O − + H 2O
→ C 2H5OH + OH −
n Typical metals used in the sol-gel
process are Si, Ti, Zr, Al, Sn, and Ce,
all with high oxidation states.
This reaction forms the basis for the start of a sequence of reactions that ultimately yields the
ceramic material. Let’s use zirconium(IV) ethoxide, Zr(C2H5O)4, to illustrate the process.
The sequence of reactions begins with the gradual addition of water to an alcohol solution of Zr(C2H5O)4, which causes ethoxide ions to react to form ethanol and be replaced
by hydroxide ions. This reaction is called hydrolysis because it involves a reaction with
water. The first step in the process is
Zr(C2H5O)4 + H2O → Zr(C2H5O)3OH + C2H5OH
The large positive charge on the zirconium ion binds the OH- tightly and polarizes the
O ! H bond, causing it to weaken. When two Zr(C2H5O)3OH units encounter each
other they undergo an acid–base reaction. A proton is transferred from the OH attached
to one zirconium ion to the OH attached to another zirconium ion. The result is the
simultaneous formation of a water molecule and an oxygen bridge between the two zirconium ions.
HO 9Zr 9 OC2H5
C2H5O 9 Zr 9 OH
C2H5O 9 Zr 9 O 9 Zr 9 OC2H5
As more and more water is added, this process continues, with ethoxide ions changing to
ethanol molecules and more and more oxygen bridges being formed between zirconium
ions. The result is a network of zirconium atoms bridged by oxygens, which produces
extremely small insoluble particles that are essentially oxides with many residual hydroxide
ions. These particles are suspended in the alcohol solvent and have a gel-like quality.7
An alkyl group is a hydrocarbon fragment derived from an alkane by removal of a hydrogen atom. For
example, the methyl group, CH3!, is derived from methane, CH4, and is an alkyl group. If we represent an
alcohol by the general formula R ! OH, where R represents an alkyl group, then the general formula of an
alkoxide ion is R ! O-.
The particle size is larger than one would find in a true solution, but smaller than in a typical precipitate.
This type of mixture is called a colloid, and a colloid composed of tiny solid particles suspended in a liquid
medium is called a sol. That’s where part of the name sol-gel comes from.
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Figure 16.5 | Sol-gel technologies and their products.
Once the sol-gel suspension is formed, it can be used in a number of ways, as indicated
in Figure 16.5. It can be deposited on a surface by dip coating to yield very thin ceramic
coatings. It can be cast into a mold to produce a semisolid gelatin-like material (wet gel).
The wet gel can be dried by evaporation of the solvent to give a porous solid called a
xerogel. Further heating causes the porous structure of the xerogel to collapse and form a
dense ceramic or glass with a uniform structure. If the solvent is removed from the wet
gel under supercritical conditions (at a temperature above the critical temperature of the
solvent), a very porous and extremely low density solid called an aerogel is formed (see the
photo on page 61). By adjusting the viscosity of the gel suspension, ceramic fibers can be
formed. And by precipitation, ultrafine and uniform ceramic powders are formed. What
is amazing is that all these different forms can be made from the same material depending
on how the suspension is handled.
n Xerogels have very large internal
surface areas because of their fine
porous structure. Some have been
found to be useful catalysts.
Brønsted–Lowry Acids and Bases. A Brønsted–Lowry acid
(often simply referred to as a Brønsted acid) is a proton donor;
a Brønsted–Lowry base (or simply Brønsted base) is a proton acceptor. According to the Brønsted–Lowry approach, an acid–base
reaction is a proton transfer event. In an equilibrium involving a
Brønsted acid and base, there are two conjugate acid–base pairs.
The members of any given acid–base pair differ from each other
by only one H+, with the acid having one more H+ than the base.
A substance that can be either an acid or a base, depending on the
nature of the other reactant, is amphoteric or, with emphasis on
proton–transfer reactions, amphiprotic.
Lewis Acids and Bases. A Lewis acid accepts a pair of electrons from a Lewis base in the formation of a coordinate covalent
bond. Lewis bases often have filled valence shells and must have
at least one unshared electron pair. Lewis acids have an incomplete valence shell that can accept an electron pair, have double
bonds that allow electron pairs to be moved to make room for an
incoming electron pair from a Lewis base, or have valence shells
that can accept more than an octet of electrons.
Relative Acidities and the Periodic Table. Binary acids
contain only hydrogen and another nonmetal. Their strengths
increase from top to bottom within a group and left to right
across a period. Oxoacids, which contain oxygen atoms in addition to hydrogen and another element, increase in strength as
the number of oxygen atoms on the same central atom increases.
Delocalization of negative charge enhances the stability of
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766 Chapter 16 | Acids and Bases, A Molecular Look
oxoacid anions, making them weaker bases and, as a result, their
conjugate acids are correspondingly stronger. Oxoacids having
the same number of oxygens generally increase in strength as the
central atom moves from bottom to top within a group and from
left to right across a period.
Acid–Base Properties of the Elements and Their
Oxides. Oxides of metals are basic anhydrides when the charge
on the ion is small. Those of the Groups 1A and 2A metals neutralize acids and tend to react with water to form soluble metal
hydroxides. The hydrates of metal ions tend to be proton donors
when the positive charge density of the metal ion is itself sufficiently high, as it is when the ion has a 3+ charge. The small
beryllium ion, Be2+, forms a weakly acidic hydrated ion. Metal
oxides become more acidic as the oxidation number of the metal
becomes larger. Aluminum oxide is amphoteric, dissolving in
both acids and bases. Chromium(VI) oxide is acidic, forming
chromic acid when it dissolves in water.
Advanced Ceramics. Common ceramic materials are made
from inorganic minerals obtained from the earth’s crust. They
are ground to fine powders, formed into shapes, and fired at high
temperatures where sintering occurs, causing the fine particles
to stick together to give a rigid solid. The sol-gel process uses
the reaction of water with metal alkoxides to form extremely
small, gel-like particles suspended in alcohol that can be used in
a variety of ways to give coatings, xerogels, aerogels, fibers, and
ultrafine powders for making ceramics.
T O O L S
Tools for Problem Solving
The following tools related to acid–base reactions were introduced in this
Brønsted–Lowry definitions (page 742)
The Brønsted–Lowry acid–base definitions, acids are proton donors and bases are proton acceptors, are useful to keep in mind
whenever working with acids and bases, especially in aqueous solutions.
Conjugate acid–base pairs (page 743)
Every conjugate acid–base pair consists of an acid plus a base with less proton (H+). You should be able to write the formula of a
conjugate acid, given the formula of the base, and vice versa.
Conjugate acid–base equilibria (page 744)
The conjugate acid is always on the opposite side of an equation from its conjugate base, and the Brønsted–Lowry acid–base
equilibrium is typically written with two sets of conjugate acid–base pairs.
Acid–base strength and the position of equilibrium (page 748)
If you know the relative strengths of the acids or bases in an equilibrium, you can predict the position of equilibrium. You can also
use the position of equilibrium to establish the relative strengths of the acids or bases in an equilibrium.
Reciprocal relationship in acid–base strengths (page 748)
The stronger an acid, the weaker is its conjugate base. This tool can help establish the position of equilibrium when dealing with
more than one acid–base pair.
Periodic trends in strengths of binary acids (page 750)
The acidities of H!X bonds increases from left to right across a period and from top to bottom in a group. You can use the trends to
predict the relative acidities of X!H bonds, both for the binary hydrides themselves and for molecules that contain X!H bonds.
Trends in the strengths of oxoacids (pages 750 and 753)
You use the trends to predict the relative acidities of oxoacids according to the nature of the central nonmetal as well as the
number of oxygens attached to a given nonmetal. The principles involved also let you compare acidities of compounds containing
different electronegative elements or electron-attracting groups bonded to an atom bonded to an !OH group.
Lewis acid–base definitions (page 755)
Keep the definitions in mind whenever analyzing an acid–base reaction in terms of Lewis acids and bases.
Acidity of metal cations (page 760)
Metal ions become better at polarizing H2O molecules as their charge increases and their size decreases. Use this to compare the
relative abilities of different metal ions to produce acidic solutions in water.
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= WileyPLUS, an online teaching and learning solution. Note to instructors: Many of the end-of-chapter problems are available for assignment via the WileyPLUS system. www.wileyplus.com.
= An Interactive Learningware solution is available for this problem.
= An Office
Hour video is available for this problem. Review Problems are presented in pairs separated by blue rules. Answers to problems whose numbers
appear in blue are given in Appendix B. More challenging problems are marked with an asterisk .
Brønsted–Lowry Acids and Bases
16.14 Acetic acid, HC2H3O2, is a weaker acid than nitrous acid,
16.1 How is a Brønsted acid defined? How is a Brønsted base
16.2 How are the formulas of the members of a conjugate acid–
base pair related to each other? Within the pair, how can
you tell which is the acid?
16.3 Is H2SO4 the conjugate acid of SO42-? Explain your answer.
16.4 What is meant by the term amphoteric ? Give two chemical
equations that illustrate the amphoteric nature of water.
HNO2. How do the strengths of the bases C2H3O2- and
16.15 Nitric acid, HNO3, is a very strong acid. It is 100% ion-
ized in water. In the reaction below, would the position
of equilibrium lie to the left or to the right?
NO3-(aq) + H2O
HNO3(aq) + OH-(aq)
16.16 HClO4 is a stronger proton donor than HNO3, but in
water both acids appear to be of equal strength; they are
both 100% ionized. Why is this so? What solvent property would be necessary in order to distinguish between
the acidities of these two Brønsted–Lowry acids?
16.5 Define the term amphiprotic.
Trends in Acid–Base Strengths
16.6 Within the periodic table, how do the strengths of the
16.17 Formic acid, HCHO2, and acetic acid, HC2H3O2, are
binary acids vary from left to right across a period? How
do they vary from top to bottom within a group?
classified as weak acids, but in water HCHO2 is more
fully ionized than HC2H3O2. However, if we use liquid
ammonia as a solvent for these acids, they both appear
to be of equal strengths; both are 100% ionized in liquid
ammonia. Explain why this is so.
16.7 Astatine, atomic number 85, is radioactive and does not
occur in appreciable amounts in nature. On the basis of
what you have learned in this chapter, answer the following.
(a) How would the acid strength of HAt compare with
that of HI?
(b) How would the acid strength of HAtO3 compare with
that of HBrO3?
16.8 Explain why nitric acid is a stronger acid than nitrous acid.
HO ! NO2
HO ! NO
16.10 Which is the stronger Brønsted–Lowry base, CH3CH2O
or CH3CH2S-? What is the basis for your selection?
16.11 Explain why HClO4 is a stronger acid than H2SeO4.
16.12 The position of equilibrium in the equation below lies far
to the left. Identify the conjugate acid–base pairs. Which
of the two acids is stronger?
H3O+(aq) + OCl-(aq)
16.13 Consider the following: CO3
is a weaker base than hydroxide ion, and HCO3- is a stronger acid than water. In
the equation below, would the position of equilibrium lie
to the left or to the right? Justify your answer.
CO3 (aq) + H2O
base? Explain your choice.
16.9 Explain why H2S is a stronger acid than H2O.
stronger Brønsted–Lowry acid? Why?
16.19 Which of the molecules below has the stronger conjugate
HOCl(aq) + H2O
16.18 Which of the molecules below is expected to be the
HCO3 (aq) + OH (aq)
Cl 9 C 9 C
Lewis Acids and Bases
16.20 Define Lewis acid and Lewis base.
16.21 Explain why the addition of a proton to a water molecule
to give H3O+ is a Lewis acid–base reaction.
16.22 Methylamine has the formula CH3NH2 and the structure
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768 Chapter 16 | Acids and Bases, A Molecular Look
Use Lewis structures to illustrate the reaction of methylamine with boron trifluoride.
16.23 Use Lewis structures to show the Lewis acid–base reac-
tion between SO3 and H2O to give H2SO4. Identify the
Lewis acid and the Lewis base in the reaction.
16.24 Explain why the oxide ion, O , can function as a Lewis
base but not as a Lewis acid.
16.25 The molecule SbF5 is able to function as a Lewis acid.
Explain why it is able to be a Lewis acid.
better written as B(OH)3. It functions not as a Brønsted–
Lowry acid, but as a Lewis acid. Using Lewis structures,
show how B(OH)3 can bind to a water molecule and
cause the resulting product to be a weak Brønsted acid.
16.30 Many chromium salts crystallize as hydrates containing
the ion Cr(H2O)63+. Solutions of these salts tend to be
acidic. Explain why.
16.31 Which ion is expected to give the more acidic solution,
Fe2+ or Fe3+? Why?
16.26 In the reaction of calcium with oxygen to form calcium
16.32 Ions of the alkali metals have little effect on the acidity of
oxide, each calcium gives a pair of electrons to an oxygen atom. Why isn’t this viewed as a Lewis acid–base
16.33 What acid is formed when the following oxides react
Acid–Base Properties of the Elements and Their Oxides
16.27 Suppose that a new element was discovered. Based on the
discussions in this chapter, what properties (both physical
and chemical) might be used to classify the element as a
metal or a nonmetal?
16.28 If the oxide of an element dissolves in water to give an
acidic solution, is the element more likely to be a metal
or a nonmetal?
16.29 Boric acid is very poisonous and is used in ant bait (to kill
ant colonies) and to poison cockroaches. It is a weak acid
with a formula often written as H3BO3, although it is
a solution. Why?
16.34 Consider the following oxides: CrO, Cr2O3, CrO3.
(a) Which is most acidic?
(b) Which is most basic?
(c) Which is most likely to be amphoteric?
16.35 Write equations for the reaction of Al2O3 with
(a) a strong acid, and
(b) a strong base.
Brønsted Acids and Bases
Trends in Acid–Base Strengths
16.36 Write the formula for the conjugate acid of each of the
16.40 Choose the stronger acid: (a) HBr or HCl; (b) H2O or
HF; (c) H2S or HBr. Give your reasons.
16.37 Write the formula for the conjugate base of each of the
16.42 Choose the stronger acid and give your reason:
(a) HOCl or HClO2;
16.41 Choose the stronger acid: (a) H2S or H2Se; (b) H2Te or
HI; (c) PH3 or NH3. Give your reasons.
16.38 Identify the conjugate acid–base pairs in the following
(a) HNO3 + N2H4
(b) NH3 + N2H5+
(c) H2PO4- + CO32(d) HIO3 + HC2O4-
NO3- + N2H5+
NH4+ + N2H4
HPO42- + HCO3IO3- + H2C2O4
(b) H2SeO4 or H2SeO3.
16.43 Choose the stronger acid: (a) HIO3 or HIO4; (b) H3AsO4
or H3AsO3. Give your reasons.
16.44 Choose the stronger acid: (a) HClO3 or HIO3; (b) HIO2
or HClO3; (c) H2SeO3 or HBrO4. Give your reasons.
16.45 Choose the stronger acid: (a) H3AsO4 or H3PO4;
(b) H2CO3 or HNO3; (c) H2SeO4 or HClO4. Give your
16.39 Identify the conjugate acid–base pairs in the following
(a) HSO4- + SO32HSO3- + SO42(b) S2- + H2O
HS- + OH+
(c) CN + H3O
HCN + H2O
(d) H2Se + H2O
HSe- + H3O+
Acidity of Metal Oxides and Hydrated Metal Ions
16.46 The ion Cr(H2O)63+ is weakly acidic. Write an equation
showing its behavior as a Brønsted–Lowry acid in water.
16.47 The compound Mg(OH)2 is basic, but Si(OH)4 is an
acid (silicic acid). Why?
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Lewis Acids and Bases
Use Lewis structures to show how the reaction
2AlCl3 → Al2Cl6 is a Lewis acid–base reaction.
16.48 Use Lewis symbols to diagram the reaction
NH2- + H+ → NH3
Identify the Lewis acid and Lewis base in the reaction.
16.49 Use Lewis symbols to diagram the reaction
BF3 + F- → BF4Identify the Lewis acid and Lewis base in the reaction.
16.50 Beryllium chloride, BeCl2, exists in the solid as a polymer
16.52 Use Lewis structures to diagram the reaction
CO2 + H2O → H2CO3
Identify the Lewis acid and Lewis base in this
16.53 Use Lewis structures to diagram the reaction
CO2 + O2- → CO32-
composed of long chains of BeCl2 units arranged as follows.
The formula of the chain can be represented as (BeCl2)n,
where n is a large number. Use Lewis structures to show
how the reaction nBeCl2 → (BeCl2)n is a Lewis acid–
16.51 Aluminum chloride, AlCl3, forms molecules with itself
with the formula Al2Cl6. Its structure is
16.54 Use Lewis structures to show how the following reaction
can be viewed as the displacement of one Lewis base by
another Lewis base from a Lewis acid. Identify the two
Lewis bases and the Lewis acid.
NH2- + H2O → NH3 + OH16.55 Use Lewis structures to show how the following reaction
involves the transfer of a Lewis base from one Lewis acid
to another. Identify the two Lewis acids and the Lewis
CO32- + SO2 → CO2 + SO32-
16.60 Which of the following compounds is the stronger base?
16.61 Which of the two molecules below is the stronger
Brønsted–Lowry acid? Why?
H 9 O 9 C 9 C 9 C 9 C 9 O 9H
(CH3)2NH? What is the formula of its conjugate base?
16.57 Using liquid ammonia as a solvent, sodium amide reacts
with ammonium chloride in an acid–base neutralization
reaction. Assuming that these compounds are completely
dissociated in liquid ammonia, write molecular, ionic,
and net ionic equations for the reaction. Which substance is the acid and which is the base?
16.58 In liquid SO2 as a solvent, SOCl2 reacts with Na2SO3 in a
reaction that can be classified as neutralization (acid plus
base yields solvent plus a salt). Write an equation for the
reaction. Which solute is the acid and which is the base?
Describe what is happening in terms of the Lewis definition of acids and bases.
16.59 The following space-filling model depicts the structure of
a compound called ethanamide.
16.56 What is the formula of the conjugate acid of dimethylamine,
Identify the Lewis acid and Lewis base in this
H 9 O9 C 9 C 9 O 9H
16.62 Write equations that illustrate the amphiprotic nature of
the bicarbonate ion.
16.63 Hydrogen peroxide is a stronger Brønsted–Lowry acid
(a) Explain why this is so.
How would you expect its base strength to compare with
that of ammonia? Justify your answer.
(b) Is an aqueous solution of hydrogen peroxide acidic or
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770 Chapter 16 | Acids and Bases, A Molecular Look
16.64 Sodium hydroxide, NaOH, is basic. Aluminum hydrox-
16.67 In the reaction in the preceding exercise, the position of
ide, Al(H2O)3(OH)3, is amphoteric. The compound
O3ClOH (usually written HClO4) is acidic. Considering
that each compound contains one or more OH groups,
why are their acid–base properties so different?
equilibrium lies to left. Identify the stronger acid in each
of the conjugate pairs in the reaction.
16.65 Hydrazine, N2H4, is a weaker Brønsted–Lowry base than
ammonia. In the following reaction, would the position
of equilibrium lie to the left or to the right? Justify your
N2H5+ + NH3
N2H4 + NH4+
16.68 How would you expect the degree of ionization of
HClO3 to compare in the solvents H2O(l ) and HF(l )?
The reactions are
HClO3 + H2O
H3O+ + ClO3-
HClO3 + HF
H2F+ + ClO3-
Justify your answer.
16.66 Identify the two Brønsted–Lowry acids and two bases in
NH2OH + CH3NH3+
NH3OH+ + CH3NH2
16.69 On the basis of the VSEPR theory, sketch the structure
16.70 A mixture is prepared containing 0.10 M of each of the
of the Al(H2O)63+ ion. Sketch the two possible structures
of the ion formed when a proton is removed from four of
the six water molecules. Which is the more likely of the
following: arsenic acid, sodium arsenate, arsenous acid,
and sodium arsenite. What are the formulas and chemical
structures of the predominant arsenic-containing species
present in the solution at equilibrium? Write the equation for the equilibrium.
in Critical Thinking
16.71 Are all Arrhenius acids Brønsted–Lowry acids? Are they
all Lewis acids? Give examples if they are not. Give a reasoned explanation if they are.
!OH group on almost every carbon atom; are they acids
or bases? Phenol is a benzene ring with an !OH group.
Is phenol an acid or base?
16.72 How could you determine whether HBr is a stronger acid
16.74 Acid rain, acid mine runoff, and acid leaching of metals
16.73 Alcohols are organic compounds that have an !OH
group. Are alcohols acids or bases? Sugars have an
from soils are important environmental considerations.
What do these topics refer to and how do they affect you
as a person?
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