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2 Strengths of Brønsted–Lowry Acids and Bases

2 Strengths of Brønsted–Lowry Acids and Bases

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16.2 | Strengths of Brønsted–Lowry Acids and Bases



747



to select some reference base, and because we are interested most in reactions in aqueous media, that base is usually water (although other reference bases could also be

chosen).

In Chapter 5 we discussed strong and weak acids from the Arrhenius point of view, and

much of what we said there applies when the same acids are studied using the Brønsted–

Lowry concept. Thus, acids such as HCl and HNO3 react completely with water to give

H3O+ because they are strong proton donors. Hence, we classify them as strong Brønsted–

Lowry acids. On the other hand, acids such as HNO2 (nitrous acid) and HC2H3O2 (acetic

acid) are much weaker proton donors. Their reactions with water are far from complete,

and we classify them as weak acids.

In a similar manner, the relative strengths of Brønsted bases are assigned according to

their abilities to accept and bind protons. Once again, to compare strengths, we have to

choose a standard acid. Because water is amphiprotic, it can serve as the standard acid as

well. Substances that are powerful proton acceptors, such as the oxide ion, react completely and are considered to be strong Brønsted–Lowry bases.

100%

O2 − + H 2O 

→ 2OH −



Weaker proton acceptors, such as ammonia, undergo incomplete reactions with water; we

classify them as weak bases.



Hydronium Ion and Hydroxide Ion in Water

Both HCl and HNO3 are very powerful proton donors. When placed in water they react

completely, losing their protons to water molecules to yield H3O+ ions. Representing

them by the general formula HA, we have

100%

HA + H 2O 

→ H3O+ + A −



acid



base



acid



base



Because both reactions go to completion, we really can’t tell which of the two, HCl or

HNO3, is actually the better proton donor (stronger acid). This would require a reference

base less willing than water to accept protons. In water, both HCl and HNO3 converted

quantitatively to another acid, H3O+. The conclusion, therefore, is that H3O+ is the strongest acid we will ever find in an aqueous solution, because stronger acids react completely

with water to give H3O+.

A similar conclusion is reached regarding hydroxide ion. We noted that the strong

Brønsted base O2- reacts completely with water to give OH-. Another very powerful

proton acceptor is the amide ion, NH2-, which also reacts completely with water.

100%



Amide

ion



Water





Ammonia Hydroxide

ion



Of course, we can also express this reaction in the form of a regular chemical equation,

%

NH 2− + H 2O 100



→ NH3 + OH−



base



acid



acid



base



Using water as the reference acid, we can’t tell which is the better proton acceptor, O2or NH2-, because both react completely, being replaced by another base, OH-. Therefore,

we can say that OH- is the strongest base we will ever find in an aqueous solution, because

stronger bases react completely with water to give OH-.



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748 Chapter 16 | Acids and Bases, A Molecular Look



Comparing Acid–Base Strengths of Conjugate Pairs

The chemical equation for the ionization of an acid in water actually shows two Brønsted–

Lowry acids. One is the acid itself and the other is hydronium ion. In almost every case,

one of the acids is stronger than the other, and the position of equilibrium tells us which

of the two acids is stronger. Let’s see how this works using the acetic acid equilibrium, in

which the position of equilibrium lies to the left.

���

HC 2H3O2 (aq ) + H 2O �

� H3O+ (aq ) + C 2H3O2− (aq )

acid



base



acid



base



The two Brønsted–Lowry acids in this equilibrium are HC2H3O2 and H3O+, and it’s helpful to think of them as competing with each other in donating protons to acceptors. The fact

that nearly all potential protons stay on the HC2H3O2 molecules, and only a relative few

spend their time on the H3O+ ions, means that the hydronium ion is a better proton donor

than the acetic acid molecule. Thus, the hydronium ion is a stronger Brønsted–Lowry acid

than acetic acid, and we inferred this relative acidity from the position of equilibrium.

The acetic acid equilibrium also has two bases, C2H3O2- and H2O. Both compete for

any available protons. But at equilibrium, most of the protons originally carried by acetic

acid are still found on HC2H3O2 molecules; relatively few are joined to H2O in the form

of H3O+ ions. This means that acetate ions must be more effective than water molecules at

obtaining and holding protons from proton donors. This is the same as saying that the acetate

ion is a stronger base than the water molecule, a fact we are able to infer, once again, from

the position of the acetic acid equilibrium.

The preceding analysis of the acetic acid equilibrium brings out an important point.

Notice that both the weaker of the two acids and the weaker of the two bases are found on

the same side of the equation, which is the side favored by the position of equilibrium.

The position of an acid–base equilibrium favors the weaker acid and base.

Acid–base strength and the

position of equilibrium



HC2H3O2(aq) + H2O

weaker acid



weaker base



H3O+(aq) + C2H3O2−(aq)

stronger acid



stronger base



Position of equilibrium lies to the left,

in favor of the weaker acid and base.



Reciprocal Relationship in Acid–Base Strength

One aid in predicting the relative strengths of conjugate acids and bases is the existence of

a reciprocal relationship.

The stronger a Brønsted acid is, the weaker is its conjugate base.

Reciprocal relationship



To illustrate, recall that HCl( g ) is a very strong Brønsted acid; it’s 100% ionized in a dilute

aqueous solution.

100%

HCl( g ) + H 2O 

→ H3O+ (aq ) + Cl − (aq )



As we explained in Chapter 5, we don’t write double equilibrium arrows for the ionization of a strong acid. Not doing so with HCl is another way of saying that the chloride

ion, the conjugate base of HCl, must be a very weak Brønsted–Lowry base. Even in the

presence of H3O+, a very strong proton donor, chloride ions aren’t able to win protons. So

HCl, the strong acid, has a particularly weak conjugate base, Cl-.

There’s a matching reciprocal relationship.

The weaker a Brønsted acid is, the stronger is its conjugate base.



Consider, for example, the conjugate pair, OH- and O2-. The hydroxide ion is the conjugate acid and the oxide ion is the conjugate base. But the hydroxide ion must be an

extremely weak Brønsted–Lowry acid; in fact, we’ve known it so far only as a base. Given



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16.2 | Strengths of Brønsted–Lowry Acids and Bases



749



the extraordinary weakness of OH- as an acid, its conjugate base, the oxide ion, must be

an exceptionally strong base. And as you’ve already learned, oxide ion is such a strong base

that its reaction with water is 100% complete. That’s why we don’t write double equilibrium arrows in the equation for the reaction.

100%



→ OH− + OH−

O2− + H2O 

base



acid



acid



base



2−



The very strong base O has



a very weak conjugate acid, OH .



An amphoteric substance will act as a base if mixed with an acid, but it will act as an

acid if mixed with a base. You might ask: If two amphoteric substances are mixed together,

which will act as the acid and which as the base? The obvious answer is that the stronger

acid will act as the acid and the other will be the base. In the reaction below, we find that

the position of equilibrium lies to the left (with the reactants).

H2S(aq) + HCO3−(aq)

acid



base



HS−(aq) + H2CO3(aq)

base



acid



Position of equilibrium lies to the left.

HCO3− must be a weaker base than HS−.



This can be interpreted as meaning that H2CO3(aq) is a stronger acid than H2S(aq),

which in turn makes HCO3-(aq) a weaker base than HS-(aq). Therefore, if we mix a solution containing HS-(aq) with a solution containing HCO3-(aq), the reaction will be

���

HS-(aq) + HCO3- (aq) �

� H2S (aq) + CO32-(aq)

base



acid



acid



base



because the stronger base, HS , will remove the H from the weaker base, HCO3-, causing the latter to actually behave as an acid.

-



+



Example 16.3



Using Reciprocal Relationships to Predict Equilibrium Positions

In the reaction below, will the position of equilibrium lie to the left or to the right, given

the fact that acetic acid is known to be a stronger acid than the hydrogen sulfite ion?

���

HSO3-(aq) + C2H3O2-(aq) �

� HC2H3O2(aq) + SO32-(aq)

n Analysis:



In the equation above, the two acids are HC2H3O2 and HSO3-. To find

the answer to the question posed in the problem, we have to ask: How are the relative

strengths of the acids related to where the position of equilibrium will be? The tools we

select must relate to this issue.



n Assembling the Tools: The position of an acid–base equilibrium favors the weaker acid and base. That’s one of the tools we need to solve this problem. The other is the reciprocal relationship between the strengths of the members of a conjugate acid–base pair.

n Solution:



We’ll write the equilibrium equation, using the given fact about the relative

strengths of acetic acid and the hydrogen sulfite ion to start writing labels.

���

HSO3-(aq) + C2H3O2-(aq) �

� HC2H3O2(aq) + SO32-(aq)

weaker acid



stronger acid



Now we’ll use our tool about the reciprocal relationships to label the two bases. The stronger acid must have the weaker conjugate base; the weaker acid must have the stronger

conjugate base.

���

HSO3-(aq) + C2H3O2-(aq) �

� HC2H3O2(aq) + SO32-(aq)

weaker acid



jespe_c16_740-770hr.indd 749



weaker base



stronger acid



stronger base



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750 Chapter 16 | Acids and Bases, A Molecular Look



Finally, because the position of equilibrium favors the weaker acid and base, the position

of equilibrium lies to the left.

n Is



the Answer Reasonable? There are two things we can check. First, both of the

weaker conjugates should be on the same side of the equation. They are, so that suggests

we’ve made the correct assignments. Second, the reaction will proceed farther in the direction of the weaker acid and base, so that places the position of equilibrium on the left,

which agrees with our answer.



Practice Exercises



16.9 | Given that HSO4- is a stronger acid than HPO42-, what is the chemical reaction

if solutions containing these ions are mixed together? (Hint: One of these must act as an

acid and the other as a base.)

16.10 | Given that HSO4- is a stronger acid than HPO42-, determine whether the substances on the left of the arrows or those on the right are favored in the following

equilibrium.

���

HSO4-(aq) + PO43-(aq) �

� SO42-(aq) + HPO42-(aq)



Table 16.1



Acidic Binary Compounds

of Hydrogen and Nonmetalsa



Group 6A



(H2O)



16.3 | Periodic Trends



in the Strengths of Acids

In most cases, acids are formed by nonmetals, either in the form of solutions of their compounds with hydrogen [such as HCl(aq)] or as solutions of their oxides, which react with water to form oxoacids (such as

H2SO4). The strengths of these acids vary in a systematic way with the

location of the nonmetal in the periodic table.



H2S Hydrosulfuric acid

H2Se Hydroselenic acid

H2Te Hydrotelluric acid

Group 7A



Trends in the Strengths of Binary Acids



HF Hydrofluoric acid

*HCl Hydrochloric acid

*HBr Hydrobromic acid

*HI Hydroiodic acid

a



The names are for the aqueous solutions of these compounds. Strong acids are indicated with asterisks.



Many (but not all) of the binary compounds between hydrogen and

nonmetals, which we may represent by HX and H2X, are acidic and are

called binary acids. Table 16.1 lists those that are acids in water. The strong

acids are marked by asterisks.

The relative strengths of binary acids correlate with the periodic table

in two ways.



The strengths of the binary acids increase from left to right within the same period.

The strengths of binary acids increase from top to bottom within the same group.

Periodic trends in strengths

of binary acids



Increases



Binary acid strength

Increases



jespe_c16_740-770hr.indd 750



Two factors can account for these variations. One is the electronegativity of the nonmetal X,

and the other is the strength of the H!X bond.

Variations in electronegativity of the atom X alter the polarity of the H!X bond, so as

X becomes more electronegative, the partial positive charge (d+) on H becomes greater.

This makes it easier for the hydrogen to separate as H+, so the molecule becomes a better

proton donor.

A major influence on acidity is the strength of the H!X bond. Breaking this bond is

essential for the hydrogen to separate as an H+ ion, so anything that contributes to variations in bond strength will also affect variations in acid strength.

Going from left to right across a period, atomic size varies little, so the strengths of the

H!X bonds are nearly the same. The major influence is the variation in the electronegativity of X, which increases from left to right. For example, as we go from left to right

within Period 3, from S to Cl, the electronegativity increases and we find that HCl is a

stronger acid than H2S. A similar increase in electronegativity occurs going left to right in

Period 2, from O to F, so HF is a stronger acid than H2O.



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16.3 | Periodic Trends in the Strengths of Acids



751



In going down a group, there are two opposing factors. The electronegativity decreases,

so the H!X bonds become less polar. At the same time, the atoms X become much larger,

and this leads to a large drop in the H!X bond strength.4 The variation in polarity tends

to make the acids weaker, but the variation in bond strength tends to make them stronger.

Changes in bond strength seem to win out, because the acids Hn X become stronger as we

go down a group. For example, among the binary acids of the halogens the following order

of relative acidity is observed.5

HF < HCl < HBr < HI

Thus, HF is the weakest acid in the series, and HI is the strongest. The identical trend occurs

in the series of the binary acids of Group 6A elements, having formulas of the general type

H2X. Notice that these trends in acidity are the opposite of what we would expect on the

basis of trends in electronegativities, which tell us that the H!F bond is more polar than the

H!I bond and that the O!H bond is more polar than the H!S bond.



16.11 | Order the following groups of acids from the weakest to the strongest: (a) HI, HF,

HBr; (b) HCl, PH3, H2S; (c) H2Te, H2O, H2Se; (d) AsH3, HBr, H2Se; (e) HI, PH3, H2Se.

(Hint: These are all binary acids.)



Practice Exercises



16.12 | Using only the periodic table, choose the stronger acid of each pair: (a) H2Se or

HBr, (b) H2Se or H2Te, (c) H2O or H2S.



Trends in the Strengths of Oxoacids

In Chapter 5 you learned that acids composed of hydrogen, oxygen, and some other

element are called oxoacids (see Table 16.2). Those that are strong acids in water are identified in the table by asterisks.

Table 16.2

Group 4A



Some Oxoacids of Nonmetals and Metalloidsa

Group 5A



Group 6A



H2CO3 Carbonic acid *HNO3 Nitric acid

HNO2 Nitrous acid

H3PO4 Phosphoric acid

*H2SO4 Sulfuric acid

H3PO3 Phosphorous acidb H2SO3 Sulfurous acidc



H3AsO4 Arsenic acid

H3AsO3 Arsenous acid



*H2SeO4 Selenic acid

H2SeO3 Selenous acid

Te(OH)6 Telluric acide

H2TeO3 Tellurous acid



Group 7A



HFO Hypofluorous acid

*HClO4 Perchloric acid

*HClO3 Chloric acid

HClO2 Chlorous acid

HClO Hypochlorous acid

*HBrO4 Perbromic acidd

*HBrO3 Bromic acid

HIO4 Periodic acid (H5IO6)f

HIO3 Iodic acid



a



Strong acids are indicated with asterisks.

Phosphorous acid, despite its formula, is only a diprotic acid.

c

Hypothetical. An aqueous solution actually contains just dissolved sulfur dioxide, SO2(aq).

d

Pure perbromic acid is unstable; a dihydrate is known.

e

Te(OH)6 is a diprotic acid.

f

H5IO6 is formed from HIO4 + 2H2O.

b



4



Small atoms tend to form stronger bonds than large atoms, so as we go down a group there is a rapid

decrease in the H!X bond strength as the atoms X become larger.

5

We compare acid strengths by the acid’s ability to donate a proton to a particular base. As we noted earlier,

for strong acids such as HCl, HBr, and HI, water is too strong a proton acceptor to permit us to see

differences among their proton-donating abilities. All three of these acids are completely ionized in water and

so appear to be of equal strength, a phenomenon called the leveling effect (the differences are obscured or

leveled out). To compare the acidities of these acids, a solvent that is a weaker proton acceptor than water

(liquid HF or HC2H3O2, for example) has to be used.



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752 Chapter 16 | Acids and Bases, A Molecular Look



A feature common to the structures of all oxoacids is the presence of O!H groups

bonded to some central atom. For example, the structures of two oxoacids of the

Group 6A elements are

O



O



H9O9S9O9H



H 9 O 9 Se 9 O 9 H



O

H2SO4



O

H2SeO4



sulfuric acid



selenic acid



When an oxoacid ionizes, the hydrogen that’s lost as an H+ comes from the same kind

of bond in every instance—specifically, an O!H bond. The “acidity” of such a hydrogen,

meaning the ease with which it’s released as H+, is determined by how the group of atoms

attached to the oxygen affects the polarity of the O!H bond. If this group of atoms

makes the O!H bond more polar, it will cause the H to come off more easily as H+ and

thereby increase the acidity of the molecule.

δ−



δ+



G9O9H



If the group of atoms, G, attached to the O 9 H group is able to draw

electron density from the O atom, the O will pull electron density

from the O 9 H bond, thereby making the bond more polar.



It turns out that there are two principal factors that determine how the polarity of the

O!H bond is affected. One is the electronegativity of the central atom in the oxoacid and

the other is the number of oxygens attached to the central atom.

Effect of the Electronegativity of the Central Atom

To study the effects of the electronegativity of the central atom, we must compare oxoacids having the same number of oxygens. When we do this, we find that as the electronegativity of the central atom increases, the oxoacid becomes a better proton donor (i.e.,

a stronger acid). The following diagram illustrates the effect.

δ−



δ+



9X9O9H



As the electronegativity of X increases,

electron density is drawn away from O,

which draws electron density away from

the O 9 H bond. This makes the bond more polar

and makes the molecule a better proton donor.



Because electronegativity increases from bottom to top within a group and from left to

right within a period, we can make the following generalization.

Trends in the strengths

of oxoacids



When the central atoms of oxoacids hold the same number of oxygen atoms, the acid

strength increases from bottom to top within a group and from left to right within a period.



In Group 6A, for example, H2SO4 is a stronger acid than H2SeO4 because sulfur is more

electronegative than selenium. Similarly, among the halogens, acid strength increases for

acids with the formula HXO4 as follows:

Increases



Oxoacid strength

Increases



HIO4 < HBrO4 < HClO4

Going from left to right within Period 3, we can compare the acids H3PO4, H2SO4, and

HClO4, where we find the following order of acidities.

H3PO4 < H2SO4 < HClO4



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16.3 | Periodic Trends in the Strengths of Acids



16.13 | Which is the stronger acid: (a) HClO3 or HBrO3, (b) H3PO4 or H2SO4? (Hint:

Note that each pair has the same number of oxygen atoms.)



753



Practice Exercises



16.14 | In each pair indicate the weaker acid: (a) H3PO4 or H3AsO4, (b) HIO4 or H2TeO4.

Effect of the Number of Oxygens Bound to the Central Atom

Comparing oxoacids with the same central atom, we find that as the number of lone oxygens increases, the oxoacid becomes a better proton donor. (A lone oxygen is one that is

bonded only to the central atom and not to a hydrogen.) Thus, comparing HNO3 with

HNO2, we find that HNO3 is the stronger acid. To understand why, let’s look at their

molecular structures.

δ−



O



δ−



O"N9O9H <



N9O9H

O

δ−



nitrous acid



nitric acid



In an oxoacid, lone oxygens pull electron density away from the central atom, which

increases the central atom’s ability to draw electron density away from the O!H bond.

It’s as though the lone oxygens make the central atom more electronegative. Therefore,

the more lone oxygens that are attached to a central atom, the more polar will be the

O!H bonds of the acid and the stronger will be the acid. Thus, the two lone oxygens

in HNO3 produce a greater effect than the one lone oxygen in HNO2, so HNO3 is the

stronger acid.

Similar effects are seen among other oxoacids, as well. For the oxoacids of chlorine, for

instance, we find the following trend:

HClO < HClO2 < HClO3 < HClO4

O



O



O " Cl 9 O 9 H



O " Cl 9 O 9 H



HClO



HClO2



HClO3



electronegative element and has

a strong tendency to pull electron

density away from any atom to which

it is attached.



n Usually, the formula for



Comparing their structures, we have

Cl 9 O 9 H



n Oxygen is the second most



O " Cl 9 O 9 H



hypochlorous acid is written HOCl to

reflect its molecular structure. We’ve

written it HClO here to make it easier

to follow the trend in acid strengths

among the oxoacids of chlorine.



O

HClO4



This leads to another generalization.

For a given central atom, the acid strength of an oxoacid increases with the number of

oxygens held by the central atom.



The ability of lone oxygens to affect acid strength extends to organic compounds as

well. For example, compare the molecules below.

H



H



H 9 C 99 C 9 O 9 H

H



H

ethanol



H



Trends in the strengths of

oxoacids



O



H 9 C 99 C 9 O 9 H

H

acetic acid



In water, ethanol (ethyl alcohol) is not acidic at all. Replacing the two hydrogens on the

carbon adjacent to the OH group with an oxygen, however, yields acetic acid. The greater

ability of oxygen to pull electron density from the carbon produces a greater polarity of the

O ! H bond, which is one factor that causes acetic acid to be a better proton donor than

ethanol.



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754 Chapter 16 | Acids and Bases, A Molecular Look



The Reciprocal Relationship—A Molecular Interpretation

On page 748 we noted the reciprocal relationship between the strength of an acid and that

of its conjugate base. For oxoacids, the lone oxygens play a part in determining the basicity

of the anion formed in the ionization reaction. Consider the acids HClO3 and HClO4.

O



O



H 9 O 9 Cl " O



H 9 O 9 Cl " O

O

HClO4



HClO3



From our earlier discussion, we expect HClO4 to be a stronger acid than HClO3, which it

is. Ionizations of their protons yield the anions ClO3- and ClO4-.





O



O



O 9 Cl " O







O 9 Cl " O

O



ClO3−



n The concept of delocalization of

electrons to provide stability was

presented in Section 10.8.



Practice Exercises



ClO4−



In an oxoanion (an anion formed from an oxoacid) the lone oxygens carry most of the negative charge, which is delocalized over the lone oxygens. In the ClO3- ion, the single negative charge is spread over three oxygens, so each carries a charge of about 31 -. This is not

a formal charge obtained by the rules on page 389. We’re saying that the 1- charge on the

ion is delocalized over three oxygens to give each an actual charge of 31 -. By the same

reasoning, in the ClO4- ion each oxygen carries a charge of about 14 -. The smaller negative charge on the oxygens in ClO4- makes this ion less able than ClO3- to attract H+ ions

from H3O+, so ClO4- is a weaker base than ClO3-. Thus, the anion of the stronger acid

is the weaker base.



16.15 | In each pair, select the stronger acid: (a) HIO3 or HIO4, (b) H2TeO3 or H2TeO4,

(c) H3AsO3 or H3AsO4. (Hint: This problem focuses on the effect of oxygen atoms in

acids.)

16.16 | In each pair, select the weaker acid: (a) H2SO4 or HClO4, (b) H3AsO4 or H2SO4.



Strengths of Organic Acids

On page 753 we noted that replacing two H atoms in an alcohol by a lone oxygen causes

the resulting molecule to be acidic. The acidity of the molecule can be increased further

if other electronegative groups such as halogens are bonded to carbon atoms near the

!CO2H group of the acid. Such a group, labeled X below, pulls electron density from

the carbon atom of the carboxyl group, further increasing the polarity of the O ! H bond.

This causes an increase in the acidity of the molecule.

"



O



X 999 C 9 O 9 H

As the ability of X to draw electrons

away from C increases, the polarity

of the O ! H bond increases and the

molecule becomes more acidic.



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16.4 | Lewis Definition of Acids and Bases



755



Thus, chloroacetic acid is a stronger acid than acetic acid, and dichloroacetic acid and

trichloroacetic acids are increasingly stronger acids.

CH3CO2H < CH2ClCO2H < CHCl2CO2H < CCl3CO2H

Addition of each chlorine effectively withdraws more electron density from the O!H

bond, resulting in a weaker bond and a stronger acid.



16.17 | How would you expect the acidities of the following molecules to compare?

(b) F 9 C 9 C 9 OH



H



H



O



"



O



9 9



Cl 9 C 9 C 9 OH



H



"



"



O



9 9



(a)



9 9



H



Practice Exercises



(c) Br 9 C 9 C 9 OH



H



H



16.4 | Lewis Definition



of Acids and Bases



In our preceding discussions, acids and bases have been characterized by their tendency to

lose or gain protons. However, there are many reactions not involving proton transfer that

have properties we associate with acid–base reactions. For example, if gaseous SO3 is

passed over solid CaO, a reaction occurs in which CaSO4 forms.

CaO(s) + SO3(g) → CaSO4(s)



(16.2)



If these reactants are dissolved in water first, they react to form Ca(OH)2 and H2SO4, and

when their solutions are mixed the following reaction takes place.

Ca(OH)2(aq) + H2SO4(aq) → CaSO4(s) + 2H2O



(16.3)



The same two initial reactants, CaO and SO3, form the same ultimate product, CaSO4. It

certainly seems that if Reaction 16.3 is an acid–base reaction, we should be able to consider Reaction 16.2 to be an acid–base reaction, too. But there are no protons being transferred, so our definitions require further generalizations. These were provided by G. N.

Lewis, after whom Lewis symbols are named.

Lewis Definitions of Acids and Bases



1. A Lewis acid is any ionic or molecular species that can accept a pair of electrons

in the formation of a coordinate covalent bond.

2. A Lewis base is any ionic or molecular species that can donate a pair of electrons

in the formation of a coordinate covalent bond.

3. Neutralization is the formation of a coordinate covalent bond between the donor

(base) and the acceptor (acid).



Lewis acid–base definitions

n Remember, a coordinate covalent



bond is just like any other covalent

bond once it has formed. By using

this term, we are following the

origin of the electron pair that forms

the bond.



Examples of Lewis Acid–Base Reactions

The reaction between BF3 and NH3 illustrates a Lewis acid–base neutralization. The reaction is exothermic because a bond is formed between N and B, with the nitrogen donating

an electron pair and the boron accepting it.



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756 Chapter 16 | Acids and Bases, A Molecular Look

New bond—

coordinate

covalent



H



H F



F



H9N + B9F

H



H9N B9F

H F



F



Lewis

base



Lewis

acid



n A similar reaction was described

in Chapter 9 between BCl3 and

ammonia. Compounds like BF3NH3,

which are formed by simply joining

two smaller molecules, are called

addition compounds.



NH3



BF3



NH3BF3



Note that as the bond forms from ammonia to boron

trifluoride, the geometry around the boron changes from

planar triangular to tetrahedral, which is what is

expected with four bonds to the boron.



The ammonia molecule thus acts as a Lewis base. The boron atom in BF3, having only six

electrons in its valence shell and needing two more to achieve an octet, accepts the pair of

electrons from the ammonia molecule. Hence, BF3 is functioning as a Lewis acid.

As this example illustrates, Lewis bases are substances that have completed valence

shells and unshared pairs of electrons (e.g., NH3, H2O, and O2-). A Lewis acid, on the other

hand, can be a substance with an incomplete valence shell, such as BF3 or H+.

A substance can also be a Lewis acid even when it has a central atom with a complete

valence shell. This works when the central atom has a double bond that, by the shifting of

an electron pair to an adjacent atom, can make room for an incoming pair of electrons

from a Lewis base. Carbon dioxide is an example. When carbon dioxide is bubbled into

aqueous sodium hydroxide, the gas is instantly trapped as the bicarbonate ion.

CO2(g) + OH-(aq) → HCO3-(aq)

Lewis acid–base theory represents the movement of electrons in this reaction as follows.

One of the electron pairs of the

double bond shifts to the oxygen atom

as a covalent bond forms from the O

atom of the OH ion to the C atom.



H9O



New covalent bond

(coordinate covalent)



O

C

O



Lewis base



H



O

O9C

O



Lewis acid



CO2



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16.4 | Lewis Definition of Acids and Bases



757



The donation of an electron pair from the oxygen of the OH- ion produces a bond, so the

OH- ion is the Lewis base. The carbon atom of the CO2 accepts the electron pair, so CO2

is the Lewis acid.

Lewis acids can also be substances that have valence shells capable of holding more

electrons. For example, consider the reaction of sulfur dioxide as a Lewis acid with oxide

ion as a Lewis base to make the sulfite ion. This reaction occurs when gaseous sulfur

dioxide, an acidic anhydride, mingles with solid calcium oxide, a basic anhydride, to give

calcium sulfite, CaSO3.

SO2( g ) + CaO(s) → CaSO3(s)

Let’s see how electrons relocate as the sulfite ion forms. We use one of the two resonance

structures of SO2 and one of the three such structures for the sulfite ion.

new covalent bond

(coordinate covalent)



O



2



O

O



S



2



O



S9O

O



sulfur

dioxide



oxide ion



Lewis acid



Lewis base



sulfite ion



O2−



SO2



SO32−



The two very electronegative oxygens attached to the sulfur in SO2 give the sulfur a substantial positive partial charge, which induces the formation of the coordinate covalent

bond from the oxide ion to the sulfur. In this case, sulfur can accommodate more than an

octet in its valence shell, so relocation of electron pairs is not necessary.

In the chapter opening photograph we described some of the chemistry involved in

steelmaking by the basic oxygen process. In this process, O2 is blown through the molten

steel mixture, causing carbon and other nonmetals to be oxidized. Silicon, for example, is

oxidized to SiO2 and must be removed from the steel. To accomplish this, calcium oxide

is added that reacts with the silicon dioxide to form calcium silicate.

CaO(s)

calcium oxide



+



SiO2(l ) →

silicon dioxide



CaSiO3(l )

calcium silicate



This is a reaction between the Lewis base O2- and the Lewis acid SiO2. At the molecular

level the reaction is complex because SiO2 is the empirical formula for a polymeric oxide

containing oxygen bridges between silicon atoms; the structure of the product is similarly

complicated. At the high temperature of the furnace, the product is a liquid called slag,

which is easily separated from the steel.

Table 16.3 summarizes the kinds of substances that behave as Lewis acids and bases.

Study the table and then work on the Practice Exercises below.



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