Chapter 16, Section 1

16.1 Acids and Bases: A Brief Review

From the earliest days of experimental chemistry, scientists have recognized acids and bases by their characteristic properties.
Acids have a sour taste (for example, citric acid in lemon juice) and cause certain dyes to change color (for example, litmus turns red on contact with acids).
Indeed, the word acid comes from the Latin word acidus, meaning sour or tart.
Bases, in contrast, have a bitter taste and feel slippery (soap is a good example).
The word base comes from an old English meaning of the word, which is "to bring low."
(We still use the word debase in this sense, meaning to lower the value of something.)
When bases are added to acids, they lower the amount of acid.
Indeed, when acids and bases are mixed in certain proportions, their characteristic properties disappear altogether.
(Section 4.3)

Historically, chemists have sought to relate the properties of acids and bases to their compositions and molecular structures.
By 1830 it was evident that all acids contain hydrogen but not all hydrogen-containing substances are acids.
In the 1880s the Swedish chemist Svante Arrhenius (1859-1927) linked acid behavior with the presence of  and base behavior with the presence of  in aqueous solution.
He defined acids as substances that produce  in water, and bases as substances that produce in water.
Indeed, the properties of aqueous solutions of acids, such as sour taste, are due to  whereas the properties of aqueous solutions of bases are due to  Over time the Arrhenius concept of acids and bases came to be stated in the following way: Acids are substances that, when dissolved in water, increase the concentration of  ions.
Likewise, bases are substances that, when dissolved in water, increase the concentration of  ions.

Hydrogen chloride is an Arrhenius acid.
Hydrogen chloride gas is highly soluble in water because of its chemical reaction with water, which produces hydrated  and

The aqueous solution of HCl is known as hydrochloric acid.
Concentrated hydrochloric acid is about  by mass and is 12 M in HCl.

Sodium hydroxide is an Arrhenius base.
Because NaOH is an ionic compound, it dissociates into  and  when it dissolves in water, thereby releasing into the solution.

Chapter 16, Section 2

16.2 Brønsted-Lowry Acids and Bases

The Arrhenius concept of acids and bases, while useful, has limitations.
For one thing, it is restricted to aqueous solutions.
In 1923 the Danish chemist Johannes Brønsted (1879-1947) and the English chemist Thomas Lowry (1874-1936) proposed a more general definition of acids and bases.
Their concept is based on the fact that acid-base reactions involve the transfer of  from one substance to another.

In Equation 16.1 hydrogen chloride is shown ionizing in water to form An  ion is simply a proton with no surrounding valence electron.
This small, positively charged particle interacts strongly with the nonbonding electron pairs of water molecules to form hydrated hydrogen ions.
For example, the interaction of a proton with one water molecule forms the hydronium ion,

The formation of hydronium ions is one of the complex features of the interaction of the  ion with liquid water.
In fact, the  ion can form hydrogen bonds to additional  molecules to generate larger clusters of hydrated hydrogen ions, such as  and  (Figure 16.1).

Figure 16.1 Lewis structures and molecular models for  and .
There is good experimental evidence for the existence of both these species.

Chemists use  and  interchangeably to represent the same thing -- namely the hydrated proton that is responsible for the characteristic properties of aqueous solutions of acids.
We often use the  ion for simplicity and convenience, as we did in Equation 16.1.
The  ion, however, more closely represents reality.

Proton-Transfer Reactions

When we closely examine the reaction that occurs when HCl dissolves in water, we find that the HCl molecule actually transfers an  ion (a proton) to a water molecule as depicted in Figure 16.2.
Thus, we can represent the reaction as occurring between an HCl molecule and a water molecule to form hydronium and chloride ions:

Figure 16.2 When a proton is transferred from HCl to  HCl acts as

the Brønsted-Lowry acid and  acts as the Brønsted-Lowry base.

Brønsted and Lowry proposed defining acids and bases in terms of their ability to transfer protons.
According to their definition, an acid is a substance (molecule or ion) that can donate a proton to another substance.
Likewise, a base is a substance that can accept a proton.
Thus, when HCl dissolves in water (Equation 16.3), HCl acts as a Brønsted-Lowry acid (it donates a proton to  and  acts as a Brønsted-Lowry base (it accepts a proton from HCl).

Because the emphasis in the Brønsted-Lowry concept is on proton transfer, the concept also applies to reactions that do not occur in aqueous solution.
In the reaction between HCl and  for example, a proton is transferred from the acid HCl to the base

This reaction can occur in the gas phase.
The hazy film that forms on the windows of general chemistry laboratories and on glassware in the lab is largely solid  formed by the gas-phase reaction of HCl and  (Figure 16.3).

Figure 16.3 The HCl(g) escaping from concentrated hydrochloric acid and the  escaping from aqueous ammonia (here labeled ammonium hydroxide) combine to form a white fog of

Let's consider another example that compares the relationship between the Arrhenius definitions and the Brønsted-Lowry definitions of acids and bases -- an aqueous solution of ammonia, in which the following equilibrium occurs:

Ammonia is an Arrhenius base because adding it to water leads to an increase in the concentration of  It is a Brønsted-Lowry base because it accepts a proton from  The  molecule in Equation 16.5 acts as a Brønsted-Lowry acid because it donates a proton to the  molecule.

An acid and a base always work together to transfer a proton.
In other words, a substance can function as an acid only if another substance simultaneously behaves as a base.
To be a Brønsted-Lowry acid, a molecule or ion must have a hydrogen atom that it can lose as an  ion.
To be a Brønsted-Lowry base, a molecule or ion must have a nonbonding pair of electrons that it can use to bind the  ion.

Some substances can act as an acid in one reaction and as a base in another.
For example,  is a Brønsted-Lowry base in its reaction with HCl (Equation 16.3) and a Brønsted-Lowry acid in its reaction with (Equation 16.5).
A substance that is capable of acting as either an acid or a base is called amphoteric.
An amphoteric substance acts as a base when combined with something more strongly acidic than itself, and as an acid when combined with something more strongly basic than itself.

Conjugate Acid-Base Pairs

In any acid-base equilibrium both the forward reaction (to the right) and the reverse reaction (to the left) involve proton transfers.
For example, consider the reaction of an acid, which we will denote HX, with water.

In the forward reaction HX donates a proton to  Therefore, HX is the Brønsted-Lowry acid, and  is the Brønsted-Lowry base.
In the reverse reaction the  ion donates a proton to the  ion, so  is the acid and  is the base.
When the acid HX donates a proton, it leaves behind a substance,  which can act as a base.
Likewise, when  acts as a base, it generates  which can act as an acid.

An acid and a base such as HX and  that differ only in the presence or absence of a proton are called a conjugate acid-base pair.* Every acid has a conjugate base, formed by removing a proton from the acid.
For example,  is the conjugate base of  and  is the conjugate base of HX.
Similarly, every base has associated with it a conjugate acid, formed by adding a proton to the base.
Thus,  is the conjugate acid of  and HX is the conjugate acid of

In any acid-base (proton-transfer) reaction we can identify two sets of conjugate acid-base pairs.
For example, consider the reaction between nitrous acid  and water:

Likewise, for the reaction between  and  (Equation 16.5), we have

(a) What is the conjugate base of each of the following acids:

(b) What is the conjugate acid of each of the following bases:

Analyze: We are asked to give the conjugate base for each of a series of species and to give the conjugate acid for each of another series of species.

Plan: The conjugate base of a substance is simply the parent substance minus one proton, and the conjugate acid of a substance is the parent substance plus one proton.

Solve: (a)  less one proton  is  The other conjugate bases are   and  (b) plus one proton  is HCN.
The other conjugate acids are   and

Notice that the hydrogen carbonate ion  is amphoteric: It can act as either an acid or a base.

Write the formula for the conjugate acid of each of the following:    CO.

The hydrogen sulfite ion  is amphoteric.
(a) Write an equation for the reaction of  with water, in which the ion acts as an acid.
(b) Write an equation for the reaction of  with water, in which the ion acts as a base.
In both cases identify the conjugate acid-base pairs.

Analyze and Plan: We are asked to write two equations representing reactions between  and water, one in which  should donate a proton to water, thereby acting as a Brønsted-Lowry acid, and one in which  should accept a proton from water, thereby acting as a base.
We are also asked to identify the conjugate pairs in each equation.

The conjugate pairs in this equation are  (acid) and  (conjugate base); and  (base) and  (conjugate acid).

The conjugate pairs in this equation are  (acid) and  (conjugate base); and  (base) and  (conjugate acid).

When lithium oxide  is dissolved in water, the solution turns basic from the reaction of the oxide ion  with water.
Write the reaction that occurs, and identify the conjugate acid-base pairs.

Answer:   is the conjugate acid of the base   is also the conjugate base of the acid

Relative Strengths of Acids and Bases

Some acids are better proton donors than others; likewise, some bases are better proton acceptors than others.
If we arrange acids in order of their ability to donate a proton, we find that the more readily a substance gives up a proton, the less readily its conjugate base accepts a proton.
Similarly, the more readily a base accepts a proton, the less readily its conjugate acid gives up a proton.
In other words, the stronger an acid, the weaker is its conjugate base; the stronger a base, the weaker is its conjugate acid.
Thus, if we know something about the strength of an acid (its ability to donate protons), we also know something about the strength of its conjugate base (its ability to accept protons).

The inverse relationship between the strengths of acids and the strengths of their conjugate bases is illustrated in Figure 16.4.
Here we have grouped acids and bases into three broad categories based on their behavior in water.

Figure 16.4 Relative strengths of some common conjugate acid-base pairs, which are listed opposite one another in the two columns.

We can think of proton-transfer reactions as being governed by the relative abilities of two bases to abstract protons.
For example, consider the proton transfer that occurs when an acid HX dissolves in water:

If  (the base in the forward reaction) is a stronger base than  (the conjugate base of HX), then  will abstract the proton from HX to produce and  As a result, the equilibrium will lie to the right.
This describes the behavior of a strong acid in water.
For example, when HCl dissolves in water, the solution consists almost entirely of  and  with a negligible concentration of HCl molecules.

is a stronger base than  (Figure 16.4), so  acquires the proton to become the hydronium ion.

When  is a stronger base than  the equilibrium will lie to the left.
This situation occurs when HX is a weak acid.
For example, an aqueous solution of acetic acid  consists mainly of  molecules with only a relatively few  and

is a stronger base than  (Figure 16.4) and therefore abstracts the proton from  From these examples we conclude that in every acid-base reaction the position of the equilibrium favors transfer of the proton to the stronger base.