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5 Acid–Base Properties of Elements and Their Oxides

5 Acid–Base Properties of Elements and Their Oxides

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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+



δ+

Mn+

δ+



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

nearby.



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.

δ+



H



H O

Mn+ O



H

H



+



H



δ+



H



(n −1)+



M9O



+



H9O

H



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.

H2O



n+



n+



M(H2O)n+







MOH(n –1)+



+

H3O+



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



jespe_c16_740-770hr.indd 760



ionic charge

ionic volume



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



11/16/10 12:17 PM



16.5 | Acid–Base Properties of Elements and Their Oxides

3+



2+



Al(H2O)5(OH)2+



Al(H2O)63+

(b)



(a)



761



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

equilibrium

���

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

powdered components.



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

Ceramic Materials



16.1



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

grinding wheels.

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)



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763



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)



12/2/10 3:43 PM



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.

+



−H

C 2H5OH 

→ C 2 H 5O −



ethanol



ethoxide ion







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

hydroxide ion.

100%

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.



C2H5O



OC2H5



H2O



OC2H5

9 9



9



HO 9Zr 9 OC2H5



C2H5O



9 9



9



C2H5O 9 Zr 9 OH



OC2H5



9



9



C2H5O



C2H5O 9 Zr 9 O 9 Zr 9 OC2H5

C2H5O



OC2H5

+ H2O



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

6



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-.

7

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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Summary



Xerogel film

Metal

alkoxide

solution



765



Dense film



Coating



Heat



ng



ati



Co



Wet gel

Xerogel



Hydrolysis

polymerization



Evaporation



Gelling



Dense

ceramics

Heat



Aerogel



Uniform particles



Sol



Extraction

of solvent



Precipitating



Furnace



Spinning



Ceramic fibers



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.



| Summary

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



jespe_c16_740-770hr.indd 765



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



11/16/10 12:17 PM



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



chapter.



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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11/16/10 12:17 PM



Review Questions



767



= 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 .



| Review



Questions



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



defined?

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

NO2- compare?



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



nitric acid



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)

2-



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



jespe_c16_740-770hr.indd 767



F



O



H9C9C

H



Cl

-



-



H



O



F9C9C

O9H



F



O9H



base? Explain your choice.



16.9 Explain why H2S is a stronger acid than H2O.



2-



stronger Brønsted–Lowry acid? Why?



16.19 Which of the molecules below has the stronger conjugate



nitrous acid



HOCl(aq) + H2O



16.18 Which of the molecules below is expected to be the



-



HCO3 (aq) + OH (aq)



Cl



O



Cl 9 C 9 C

Cl



O



H9C9C

O9H



Cl



O9H



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



H



H



H9C9N

H



H



11/16/10 12:17 PM



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.

2-



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

reaction?



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



| Review



a solution. Why?

with water?

(a) SO3

(b) CO2

(c) P4O10

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.



Problems



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.



following.

- 



(a) F



-



(c) C5H5N



(e) HCrO4



(b) N2H4



(d) O2



16.37 Write the formula for the conjugate base of each of the



-



(b) HSO3



16.42 Choose the stronger acid and give your reason:



(a) HOCl or HClO2;



following.

(a) NH2OH



16.41 Choose the stronger acid: (a) H2S or H2Se; (b) H2Te or



HI; (c) PH3 or NH3. Give your reasons.



2-



(c) HCN



(e) HNO2



(d) H5IO6



16.38 Identify the conjugate acid–base pairs in the following



reactions.

(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

reasons.



16.39 Identify the conjugate acid–base pairs in the following



reactions.

(a) HSO4- + SO32HSO3- + SO42(b) S2- + H2O

HS- + OH+

(c) CN + H3O

HCN + H2O

(d) H2Se + H2O

HSe- + H3O+



jespe_c16_740-770hr.indd 768



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?



11/16/10 12:17 PM



Additional Exercises



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

reaction.

16.53 Use Lewis structures to diagram the reaction



CO2 + O2- → CO32-



composed of long chains of BeCl2 units arranged as follows.

Cl

Cl



Be

Cl



Cl

Be



Cl



Cl

Be



Cl



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–

base reaction.

16.51 Aluminum chloride, AlCl3, forms molecules with itself



with the formula Al2Cl6. Its structure is

Cl



Cl

Al



Cl



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

base.

CO32- + SO2 → CO2 + SO32-



Al

Cl



Cl



Exercises

16.60 Which of the following compounds is the stronger base?



Explain.



or



16.61 Which of the two molecules below is the stronger



Brønsted–Lowry acid? Why?

H



H



O



H 9 O 9 C 9 C 9 C 9 C 9 O 9H

H



O



O



"



O



"



(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,



9 9



| Additional



Cl



Identify the Lewis acid and Lewis base in this

reaction.



9 9



Cl



Be



Cl



"



Be



Cl



769



H 9 O9 C 9 C 9 O 9H



H



16.62 Write equations that illustrate the amphiprotic nature of



the bicarbonate ion.

16.63 Hydrogen peroxide is a stronger Brønsted–Lowry acid



than water.

(a) Explain why this is so.

How would you expect its base strength to compare with

that of ammonia? Justify your answer.



jespe_c16_740-770hr.indd 769



(b) Is an aqueous solution of hydrogen peroxide acidic or

basic?



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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

answer.

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



the reaction

NH2OH + CH3NH3+



| Multi-Concept



NH3OH+ + CH3NH2



Problems



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

two structures?



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.



| Exercises



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



than HI?

16.73 Alcohols are organic compounds that have an !OH



group. Are alcohols acids or bases? Sugars have an



jespe_c16_740-770hr.indd 770



from soils are important environmental considerations.

What do these topics refer to and how do they affect you

as a person?



11/16/10 12:17 PM



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