Chapter 15

Properties and reactions of acids and

bases

15.1 properties of acids and bases

Acids and bases are important compounds in many aspects of our daily

lives. Acids and bases are significant in human biology, the environment,

agriculture and industry. Examples of commonly used acids include sulfuric acid

(battery acids), hydrochloric acid (spirits of salt; gastric juices) and ethanoic or

acetic acid (vinegar). Certain fruits such as lemons and grapefruit also exhibit

acidic properties due to the presence of acids such as citric acid.

Common household and industrial bases include sodium hydroxide (caustic soda),

sodium carbonate (soda ash), sodium carbonate-10-water (washing soda) and

ammonia. Bases are used in a variety of household cleaning agents such as oven,

drain and window cleaners. Bases such as magnesium hydroxide and aluminum

hydroxide are also used in various preparations as antacids for the relief of upset

stomachs.

In agriculture, soil acidity is an important factor in determining what crops

may be grown. Sometimes chemical such as lime (calcium oxide) are added to

make the soil more effective for growing particular crops.

Properties of acids

Aqueous solutions of acids such as hydrochloric, sulfuric and ethanoic

acids exhibit a range of common properties. These properties result from the

production of hydrogen ions (H+) or hydronium (H3O

+) in solutions. The common

properties of aqueous solutions of acids are shown in the table 15.1.

Table 15.1 Properties of aqueous acids solutions

Turn blue litmus red

Conduct an electric current

Taste sour

React with reactive metals such as K, Na, Ca, Mg, Zn, Al and Fe to

produce hydrogen gas

Acid + reactive metal hydrogen + salt

For example, 2 HCl(aq) + Mg(s) H2(g) + MgCl2(aq)

Or 2H+

(aq) + Mg(s) H2(g) + Mg2+

(aq)

React with carbonates and hydrogenates to form carbon dioxide

gas

Acid + carbonate carbon dioxide + water + salt

For example, H2SO4(aq) + Na2CO3(s) CO2(g) + H2O(l) +

Na2CO4(aq)

Or 2H+

(aq) + Na2CO3(s) CO2(g) + H2O(l) +

2Na+

(aq)

Acid + hydrogen carbonate carbon dioxide + water + salt

For example, HNO3(aq) + KHCO3(aq) CO2(g) + H2O(l) + KNO3

(aq)

Or H+

(aq) + HCO3-(aq) CO2(g) + H2O(l)

React with metal oxides to produce a salt and water

Acid + metal oxide salt + water

For example, 2HNO3(aq) + CuO(s) Cu(NO3)2(aq) + H2O(l)

Or 2H+

(aq) + CuO(s) Cu2+

(aq) + H2O(l)

React with metal hydroxides to produce a salt plus water

Acid + metal hydroxide salt + water

For example, 2HCl(aq) + Ba(OH)2(aq) BaCl2(aq) + 2H2O(l)

Or H+

(aq) + OH-(aq) H2O(l)

If the metal hydroxide is an hydroxide or the acid is added to a

solid rather than a solution, the equations are better written as

follows

2HCl(aq) + Mg(OH)2(s) MgCl2(aq) + 2H2O(l)

2H+

(aq) + Mg(OH)2(s) Mg2+

(aq) + 2H2O(l)

Properties of metallic hydroxide bases

Aqueous solutions of metallic hydroxide bases, such as sodium hydroxide,

also exhibit some characteristic properties. These properties result from the

formation of hydroxide ion (OH-) in solution. The properties commonly shown by

aqueous solutions of metallic hydroxides are illustrated in Table 15.2

Properties of aqueous solutions of hydroxide bases

Turn red litmus blue

Conduct an electric current

Taste bitter

React with acids to produce a salt and water as indicated in Table 15.1

React with amphoteric metals such as aluminum, chromium and zinc to produce

hydrogen gas; for example,

2Al(s) + 2NaOH(aq) + 6H2O(l) 2Na[Al(OH)4](aq) + 3H2(g)

Or

2Al(s) + 2OH-(aq) + 6H2O(l) 2[Al(OH)4]

-(aq) + 3H2(g)

Dissolved amphoteric metal hydroxides such as Al(OH)3, Cr(OH)3 and Zn(OH)2,

for example, Al(OH)3(s) + NaOH(aq) Na[Al(OH)4](aq)

or Al(OH)3(s) + OH-(aq) [Al(OH)4]

-(aq)

That properties above indicates that the metals Al, Cr, and Zn react with

solutions of metal hydroxides, such as sodium hydroxide, to produce hydrogen

gas. These metals are describes as amphoteric metals, as they also react with acid

solutions to form hydrogen. An amphoteric substance is one which reacts with

both acids and bases.

When these metals react with bases to produce hydrogen they also form

complex ions. These complex ion usually contain four hydroxide ions bonded to

the central metal ion. The formulas and names of these complex ions are provided

in Table 15.3

Table 15.3 Complex ion formed by amphoteric metals, their oxides and

hydroxides in basic solutions

Metal Formula of complex

ion

Name of complex

ion

Al [Al(OH)4]-

Aluminate ion

Cr [Cr(OH)4]-

Chromite ion

Zn [Zn(OH)4]2-

Zincat ion

Table 15.2 also indicates that some insoluble metal hydroxides, such as

[A(OH)3, Cr(OH)3 and Zn(OH)2, will dissolve in basic solution as well as being

soluble in acid solution. Because these hydroxides will dissolve in acids and bases

they are called amphoteric hydroxide. In acid solution they form the aquated

metal cations such as Al3+

, Cr3+

, and Zn2+

. In base solution they form the same

complex ions as shown in Table 15.3

Review exercise 15.1

1. Write ionic equations for the reactions between the following:

a. Zinc and hydrochloric acid

b. Calcium carbonate and hydrochloric acid

c. Potassium oxide and sulfuric acid

d. Barium hydroxide solution and nitric acid

2. Write equations for the reaction between the following:

a. The amphoteric metal, chromium and potassium hydroxide solution

b. The amphoteric hydroxide, chromium (III) hydroxide, and sodium

hydroxide solution

3. The amphoteric hydroxide Al(OH)3 is insoluble in water. Write equation for

the reaction of aluminium hydroxide with the following:

a. Sulfuric acid to form aluminium ions.

b. Sodium hydroxide solution to form aluminate ions

15.2 Theories of acids and bases

Arrhenius theory

Acids

The conductivities of acid solutions indicate that these solutions contain

ions. The fact that acids react with many metals to produce hydrogen gas further

suggests that acid solutions contain hydrogen ions. The theory that an acid is a

substance which produces hydrogen ions in water was first proposed by the

Swedish chemist, Svante Arrhenius. Arrhenius proposed that hydrogen ions were

produced by the ionization of an acid in water. For example, in hydrochloric acid

the hydrogen chloride molecules are ionized to form hydrogen ions and chloride

ions.

HCl(g) H+

(aq) + Cl-(aq)

Similarly, sulfuric acid and nitric acid ionize in aqueous solutions to form

hydrogen ions. Each of these acids is a strong acid because it is essentially

completely ionized in aqueous solution. In hydrochloric acid, for example, there

are very few unionized hydrogen chloride molecules. Virtually all the molecules

have ionized to form H+ and Cl

- ions. Because the equilibrium strongly favours

the formation of products, a single arrow ahowing only the forwar reaction is used

in the equation.

In many acids such as ethanoic (acetic) and citric acid only a small

proportion of the molecules are ionized. These acids are called weak acids. The

equation for the ionization of ethanoic acid is as follows.

CH3COOH(aq) H+

(aq) + CH3COO-(aq)

The double arrow indicates that equilibrium exists between unionized

molecules of ethanoic acid and ions solution. For example, in a 0.10 molL-1

ethanoic acid solution only about 1% of ethanoic acid molecules are ionized and

the equilibrium strongly favours the reactants.

The hydrogen ion in aqueous solution:

Hydrogen ions produced in aqueous solution are sometimes represented as H+

,

H3O+ or H

+(aq). These different representations are used partly because each can be

useful in certain circumstances and partly because of doubt surrounding the nature

of protons in aqueous solutions.

Protons, because they have no electrons, are very small. They have a radius of

about 10-13

cm compared with a radius of about 10-8

cm for other cations. Because

the proton‟s charge is located in such a small volume, the attraction between is in

fact more likely to exist as a hydronium (H3O)+ ion. The formation of a

hydronium ion can be represented as follows.

In the H3O+

ion, the H+

has formed the coordinate covalent bond with one

of the lone pairs of electrons on the oxygen atom of the water molecule.

As well as the H3O+ ion, there is some evidence for the formation of other

species such as H5O2+ and H9O4

+. these species represent groups of water

molecules containing one extra proton. Their structures are shown in Figure 15.2.

From these observations it is apparent that no single species adequately

represents a proton in solution. The various symbols H+

, H+

(aq), H3O+

(aq) may be

used, although in most situations H+

(aq) is preferred.

Bases

Arrhenius also proposed that a base is a substance which produces

hydroxide ions in aqueous solution. Sodium hydroxide (NaOH), an ionic solid, is

dissociated in water to form sodium ions and hydroxide ions.

NaOH (s) Na+

(aq) + OH-(aq)

Sodium hydroxide and potassium hydroxide are examples of strong bases

because they are essentially completely dissociated into ions in solution.

As with acids, there are numerous bases in which only a small proportion as weak

bases. Ammonia is an example of a weak base. Although it contains no hydroxide

ions of its own, ammonia produces hydroxide ions by reacting with water

according to the equation:

NH3(aq) + H2O(l) NH4+

(aq) + OH-(aq)

the equilibrium between unionized NH3 molecules and NH4+

and OH- ions

strongly favours the reactants. A solution of ammonia therefore consists mainly of

dissolved NH3 molecules and only relatively small quantities of NH4+

and OH-

are present at equilibrium.

Neutralization

In the Arrhenius model, hydrogen ions are responsible for the properties of acids

and hydroxide ions for the properties of bases. In the neutralization reaction

between acids and bases, the acidic properties of H+ and basic properties of OH

-

are „neutralised‟ when these ions combine to form water molecules. This can be

represented by the equation:

H+

(aq) + OH-(aq) H2O(l)

Bronsted-Lowry theory

The Arrhenius model of acids and bases is a very useful one but is restricted in

that it is limited to aqueous solutions. A more general model of acids and bases

was developed independently by Bronsted in Denmark and Lowry in England. In

the Bronsted-Lowry theory an acid-base reaction is one that involves the transfer

is an acid while the proton acceptor is the base.

In the Bronsted-Lowry model the ionization of HCl is represented by the

following equation:

HCl(g) + H2O(l) H3O+

(aq) + Cl-(aq)

In this process the HCl is donating a proton and is therefore acting as an acid. The

H2O, which accepts a proton, is classified as a base.

In the reaction of ammonia with water:

NH3(aq) + H2O(l) NH4+

(aq) + OH-(aq)

The NH3 accepts a proton and is acting as a base. The H2O, which donates a

proton, is acting as an acid.

In the Bronsted-Lowry theory many substances can react as acids or bases. In the

two examples above,water is acting as a base in its reaction with HCl and as an

acid in its reaction with NH3. This is further illustrated by the equation for the

autoionisation of water.

H2O(l) + H2O(l) H3O+

(aq) + OH-(aq)

In this reaction one water molecule is donating a proton to the other. The proton

donor molecule is the acid and the proton acceptor is the base.

In a similar way, the hydrogencarbonate ion can act as an acid or a base. The

hydrogencarbonate ion reacts with water in two different reactions, although in

both reactions the reactants are strongly favoured at equilibrium.

HCO3-(aq) + H2O(l) H2CO3(aq) + OH

-(aq)

HCO3-(aq) + H2O(l) H3O

+(aq) + CO3

2-(aq)

In the first of these reactions the hydrogencarbonate is reacting as a base, while in

the second it is reacting as an acid.

Conjugate acid-base pairs

In the Bronsted-Lowry theory a base, after it has received a proton, has the

potential to react as an acid. Similary, an acid which has donated a proton is a

potential base. In the reaction:

CH3COOH(aq) + OH-(aq) H2O(l) + CH3COO

-(aq)

acid base acid base

the CH3COOH is acting as an acid and the OH- as a base. The CH3COO

- ion can,

under certain conditions, accept a proton. For example, in the reaction between

hydrochloric acid and sodium ethanoate:

CH3COO-(aq) + H3O

+(aq) CH3COOH(aq) + H2O(l)

base acid acid base

the CH3COO- accepets a proton from the H3O

+ and is therefore acting as a base.

The ammonium ion and ammonia, NH4+/NH3, constitue another conjugate acid-

base pair. In general, conjugate acid-base pairs can be visualized in the following

way:

A table of conjugate acid-base pairs is shown in the table 15.4. this table is

arranged in order of decreasing acid strength. Thus HCl is a strong acid, HF is a

fairly weak acid and H2O is very weak.

Table 15.4 Conjugate acid-base pairs

Name of acid Acid H+ + Base Name of base

Aci

d s

tren

gth

in

crea

ses

Hydrochloric HCl H+ + Cl- Chloride ion

bas

e st

ren

gth

in

crea

ses

Sulfuric H2SO4 H+ + HSO4

- Hydrogensulfate ion

Nitric HNO3 H+ + NO3

- Nitrate ion

Hydronium ion H3O+ H

+ + H2O Water

Hydrogensulfate ion HSO4- H

+ + SO42-

Sulfate ion

Hydrofluoric HF H+ + F

- Fluoride ion

Ethanoic CH3COOH H+ + CH3COO

- Ethanoate ion

Hydrogen sulfide H2S H+ + HS

- Hydrogensulfide ion

Carbonic H2CO3 H+ + HCO3

- Hydrogencarbonate ion

Ammonium ion NH4+ H

+ + NH3 Ammonia

Hydrogencarbonate ion HCO3- H

+ + CO32-

Carbonate ion

Hydrogen sulfide ion HS- H

+ + S2-

Sulfide ion

Water H2O H+ + OH

- Hydroxide ion

Hydroxide ion OH- H

+ + O2-

Oxide ion

The stronger a particular acid is, the weaker will be its conjugate base. Thus

the substances on the right-hand of the table are arranged in order of increasing strength

as bases. The Cl- ion is the weakest base, HCO3

- is fairly weak, and OH- is strong base.

Predicting reactions between acids and bases Table 15.4 emphasises the Bronsted-Lowry concept that acid-base reactions are

essentially reactions involving the competition for protons. The table makes it possible to

predict the extent to which acid-base reactions will take place.

Figure 15.5

In the Bronsted-

Lowry theory acid-

base reactions are

those which involve

the transfer of a

proton from one

species to another

Generally, strong acids will react with strong base to form weaker conjugate acid and

weaker conjugate bases. Thus acids on the left of table will react with bases below them

to form the corresponding conjugate acids and bases. For example, the reactions:

HCl(aq) + F-(aq)→HF(aq) + Cl

-(aq)

HF(aq) + NH3(aq) → NH4+(aq) + F

-(aq)

Occur virtually to completion. At equilibrium the products are strongly favoured and the

equation is written as going completely from reactans to products. In the first reaction

HCl is astronger acid (proton donor) than HF. Alternatively, F- is stronger base (proton

acceptor) than Cl-. Similarly, HF is a stronger acid than NH4+.

Consider the following reactions.

NH4+(aq) + H2O(l) H3O

+(aq) + NH3(aq)

In the first reaction the H3O+ is a stronger acid (proton donor) than NH4

+ and NH3 is a

stronger base (proton acceptor) than H2O. This reaction will therefore have little tendency

to occur. An equilibrium will be established in which only very small amounts of the

product are formed. In the second reaction H3O+ is a stronger acid than HF. Again, this

reaction will have little tendency to go in the forward direction. In both these reactions

double arrows are used to indicate that the reactions occur to only a limited extent.

Review exercise 15.2

1. a. identify the conjugate acids of Cl-, CO3

2-, NH3, ClO4

-, SO3

2-

b. identify the conjugate bases of HF, HSO4-, NH4

+, HPO4

2-,H3O

+

2. write two equations for the hydrogenphosphate ion reacting with water as:

a. a base b. an acid

3. for the following reactions:

i. identify the conjugate acid-base pairs

ii. predict whether the reaction is likely to occur to alarge extent or to only a

small extent as written

a. HCO3- (aq) + F

-(aq) CO3

2-(aq) + HF (aq)

b. HSO4-(aq) + NH3(aq) SO42-

(aq) + NH4+ (aq)

c. HF(aq) + H2O(l) H3O+(aq) + F

-(aq)

4. If solid sodium oxide is dissolved in water predict the composition of the

resulting solution

Water is weak electrolyte and to a very small extent undergoes auto- or self-ionization.

This is represented by the following equations.

2H2O(l) H3O+(aq) + OH

-(aq)

or

H2O(l) H+(aq) + OH

-(aq)

The first equation represents a Bronsted-Lowry approach. One water molecule, the proton

donor, is acting as an acid while the other water molecule is acting as a base. The second

simplified equation indicates that water ionizes to produce some equated hydrogen and

hydroxide ions. Both equations indicate that equal amounts of acid and base are

produced.

In the ionization of water, equilibrium strongly favours the reactants. This means that

only small concentrations of hydrogen ions and hydroxide ions are formed and most of

the water remains as unionized water molecules.

The equilibrium constant for the ionization of water is given by:

Kw = [H+] x [OH

-]=1.0x10

-14

Where [H+] is the concentration of hydrogen ions in mol L

-1 and

[OH-] is the concentration of hydroxide ions in mol L

-1

Note that there is no term for the concentration of H2O as this is solvent and its

concentration is virtually the same in all dilute aqueous solution. Kw is called the

ionization constant, or dissociation constant, for water. As with any equilibrium constant

Kw depends on the temperature but has a value of 1.0 x 10-14

at 25ºC. This means that in

any aqueous solution at 25ºC the product of the hydrogen ion and hydrogen ion and

hydroxide ion concentrations is always 1.0 x 10-14

.

In pure water or any neutral solution:

[H+] = [OH

-]=1.0x10

-7 mol L

-1

Figure 15.6

In pure water [H+] = [OH

-]=1.0x10

-7 mol L

-1. When [H

+] increases, [OH

-] decreases. When [OH

-

]increases,[H+] decreases

In acidic solution the [H+] is greater than 1.0x10

-7 mol L

-1and [OH

-] is less than

1.0x10-7

mol L-1

. Conversely, in basic solutions the [H+] is less than 1.0x10

-7 mol L

-1

whereas the [OH-] is greater than 1.0x10

-7 mol L

-1 .

Example 15.1

Calculate the concentrations of H+, Cl- and OH- in a 1.0x10-2

mol L-1

HCl solution.

HCl is a strong acid which ionizes virtually completely in aqusous solution.

HCl(g) → H+(aq) + Cl-(aq)

[H+] = [Cl

-]=1.0x10

-2 mol L

-1

Use Kw to calculate [OH-]

Kw =[H+] [OH

-]=1.0x10

-14 mol L

-1

(1.0x10-2

) [OH-] =1.0x10

-14 mol L

-1

[OH-] = 1.0 x 10

-12 mol L

-1

[H+] ,[Cl

-]=1.0x10

-2 mol L

-1 and [OH-] = 1.0x10

-12 mol L

-1

Example 15.1

Calculate the concentrations of Na+, OH

-, and H

+ in a 4.0 x 10

-3 mol L

-1 NaOH solution.

NaOH ia a strong base which dissociates completely in aqueous solution.

NaOH(s)→Na+(aq) + OH-(aq)

[Na+]=[OH-]=4.0 x 10-3

mol L-1

Use Kw to calculate [H+]

Kw =[H+] [OH

-]=1.0x10

-14 mol L

-1

[H+](4.0 x 10

-3 mol L

-1) =1.0x10

-14 mol L

-1

[H+] = 2.5 x 10

-12 mol L

-1

[Na+] ,[OH

-]=4.0x10

-3 mol L

-1 and [H

+] = 2.5 x10

-12 mol L

-1

Review exercise 15.2

1. A sample of water from a swimming pool has [OH-] = 1.6 x 10-7

mol L-1

.

a. Calculate [H+]

b. Is the swimming pool water acidic or basic?

2. Calculate the concentrations of the following.

a. H+, NO3

- and OH

- in 5.0 x 10

-3 mol L

-1 HNO3

b. H+, Cl

- and OH

- in 1.5 mol L

-1 HCl

c. K+, OH

- and H

+ in 0.25 mol L

-1 KOH

d. Ba2+

, OH- and H

+ in 6.0 x 10

-2 mol L

-1 Ba(OH)2

15.4

The pH acidity scale

Although the hydrogen ion and hydroxide ion concentrations are valid ways of

expressing the acidity or basicity of aqueous solutions they are somewhat cumber-

some because they involve the use of indices. The common method of indicating

acidity, whether it is with reference to swimming pools, soils or hair shampoos, is

the pH scale. The pH of an aqueous solution is defined by:

pH=-log10[H+]

The pH of pure water or a neutral solution is 7.0 pH values less than 7 indicate a

solution is acidic, while values greater than 7 are characteristic of basic or alkaline

solutions.

Example 15.3

Calculate the pH of a 1.0 x 10-3

mol L-1

HCl solution.

HCl(g)→H+(aq) + Cl

-(aq)

[H+] = 1.0 x 10

-3 mol L

-1

pH = -log[H+]

= -log(1.0 x 10-3

)

pH = 3.00

Example 15.4

Calculate the pH of a 2.0 x 10-2

mol L-1

NaOH solution.

NaOH(s)→Na+(aq) + OH

-(aq)

[OH-] = 2.0 x 10

-2 mol L

-1

Kw =[H+] [OH

-]=1.0x10

-14 mol L

-1

[H+](2.0 x 10

-2) =1.0x10

-14 mol L

-1

[H+] = 5.0 x 10-13 mol L

-1

pH = -log[H+]

= -log(5.0 x 10-13

)

pH = 12.30

Example 15.5

Calculate the hydrogen ion and hydroxide ion concentrations in a sample of milk

with a pH of 6.40.

pH = -log[H+]

6.40 = -log [H+]

log [H+] = -6.40

[H+] = inverse log (-6.40)

[H+] = 3.98 x 10-7 mol L

-1

[OH-] = 2.5 x 10

-8 mol L

-1

[H+] = 4.0 x 10

-7 mol L

-1 and [OH-] = 2.5 x 10

-8 mol L

-1

Table 15.5 contains values of the hydrogen ion and hydroxide ion concentrations

for aqueous solutions with various pH values.

Table 15.5 The hydrogen ion and hydroxide ion concentrations and pH values of

aqueous solutions at 25ºC

pH [H+] [OH

-]

Incr

easi

ng a

cidit

y

0 100 10

-14

1 10-1 10

-13

2 10-2 10

-12

Acidic 3 10-3 10

-11

4 10-4

10-10

5 10-5 10

-9

6 10-6 10

-8

Neutral 7 10-7 10

-7

8 10-8 10

-6

9 10-9 10

-5

10 10-10 10

-4

Basic 11 10-11

10-3

12 10-12

10-2

13 10-13

10-1

14 10-14 10

0

From table 15.5 note the following.

1. As the [H+] increases the pH decreases.

2. The lower the pH, the greater the acidity of the solution.

3. A neutral solution is one with a pH of 7.

4. A change of one pH unit represents a tenfold change in the hydrogen ion

concentration.

Several methods can be used to estimate the pH of a sample. These include the following.

1. pH paper which turus a particular colour depending on the pH of the sample.

2. Universal indicator, a solution of several acid-base indicators, which also changes

colour depending on the pH of the solution.

3. A pH meter which can be used to obtain a more precise estimate of the pH.

The approximate pH values of some common substances are shown in Table

15.6.

Table 15.6 The pH values of some common substances

Substance pH Substance pH

Gastric juice 0.8 Urine 6.0

pH

7

Acidic Basic

0 14

Neutral

Lemon juice 2.3 Milk 6.4

Vinegar 2.8 Rain water 6.5

Aerated soft drinks 2.9 Pure water 7.0

Apples 3.2 Swimming pool water 7.2

Orange juice 3.5 Human blood 7.4

Grapes 4.0 Sea water 8.5

Tomatoes 4.2 Milk of magnesia 10.5

Bread 5.6 Household ammonia 11.9

Review exercise 15.4

1. Calculate the pH of following solutions

a. 5.0 x 10-1

mol L-1

HCl

b. 0.0065 mol L-1

KOH

c. 3.6 x 10-3

mol L-1

HNO3

d. 6.5 x 10-4

mol L-1

Ca(OH)2

2. Calculate the hydrogen ion and hydroxide ion concentrations in the following

solutions.

a. Orange juice with a pH of 3.50

b. Household ammonia with a pH of 11.90

c. Gastric juice with a pH of 0.80

3. If the pH of solution decreases from 7 to 5, by what factor does the [H+] change?

Strong acids

Strong acids are those which are essentially completely ionized to produce hydrogen ions

in aqueous solution. For example, in a solution of hydrochloric acid virtually all the

hydrogen chloride molecules are ionized to form hydrogen ions and chloride ions. This

ionization reaction can be represented by the following equations.

HCl(g) → H+(aq) + Cl

-(aq)

or

HCl(g) + H2O(l)→H3O+ (aq) + Cl

-(aq)

Using the Bronsted-Lowry approach, HCl is a stronger acid (protons donor) than H3O+

and H2O is a stronger base (proton acceptor) than Cl-, so the reaction is favoured in the

forward direction. Hydrochloric, sulfuric and nitric acids are all better proton donors than

the H3O+ ion. These acids are virtually completely ionized in aqueous solutions. Table

15.7 lists the formulas and ionization equations for the commonly encountered strong

acids.

Table 15.7 Common strong acids

Name Formula Ionization equation

Hydrochloric acid HCl HCl(g) → H+(aq) + Cl

-(aq)

Sulfuric acid H2SO4 H2SO4(l) → H+(aq) + HSO4

-(aq)

Nitric acid HNO3 HNO3(g) → H+(aq) + NO3

-(aq)

Weak Acids

Most acids are only partially ionized in water and are therefore classified as weak acids.

The ionization of the weak acid, ethanoic acid (CH3COOH), can be represented by the

following equations.

CH3COOH(aq) H+

(aq) + CH3COO- (aq)

Or

CH3COOH(aq) + H2O(l) H3O++ CH3COO

-

The reversible arrows indicate that an equilibrium is established between ethanoic acid

molecules and hydrogen ions and ethanoate ions. In terms of Bronsted-Lowry theory

CH3COOH is aweaker acid (proton donor) than H3O+ and H2O iis a weaker base

(proton acceptor) than CH3COO-. The reserve reaction is therefore favoured, so that at

equilibrium only a small proportion of ethanoic acid molecules are ionized. In fact, at

25ºC, a 0,10 mol L-1 ethanoic acid solution is only about 1.3% ionized. This means that

98.7% of the ethanoic acid exists in the molecular form and only 1.3% exist as ions.

Table 15.8 contains a selection of weak acids arranged in decreasing order of acid

strength.

Table 15.8 Weak acids arranged in decreasing order of acid strength

Name Formula Ionization equation

Oxalic acid H2C2O4 H2C2O4(aq) H+(

aq) + HC2O4-(aq)

Hydrogensulfate ion HSO4- HSO4-(aq) H

+ (aq) + SO4

2-(aq)

Phosphoric acid H3PO4 H3PO4(aq) H+ (aq) + H2PO4-(aq)

Iron (III) ion Fe3+

[Fe(H2O)6]3+

(aq) H+ (aq) + [Fe(OH)(H2O)5]

2+ (aq)

Hydrofluoric acid HF HF(aq) H+ (aq) + H2PO4-(aq)

Ethanoic acid CH3COOH CH3COOH(aq) H+ (aq) + CH3COO

- (aq)

Aluminum ion Al3+

[Al(H2O)6]3+

(aq) H+

(aq) + [Al(OH)(H2O)5]2+

(aq)

Carbonic ion H2CO3 H2CO3(aq) H+ (aq)+ HCO3

-(aq)

Hydrogen sulfide H2S H2S(aq) H+(aq) + HS

-(aq)

Dihydrogenphosphate ion H2PO4- H2PO4

-(aq) H

+(aq) +HPO4

2-(aq)

Hypochlorous acid HClO HClO(aq) H+ (aq)+ ClO

-(aq)

Ammonium ion NH4+ NH4

+(aq) H

+ (aq) + NH3(aq)

Hydrocyanic acid HCN HCN (aq) H+ (aq)+ CN-(aq)

From the table it is apparent that weak acids can be molecules, anions or cations. The

molecular acids are usually identified as acids from their names. The hydrogen-

containing anions such as hydrogensulfate (HSO4-) and dihydrogenphosphate (H2PO4

-)

can be considered to be derived from the molecular acids sulfuric (H2SO4) and

phosphoric (H3PO4) acid respectively. The ammonium ion and many metal cations,

except those in groups I and II, behave as weak acids in aqueous solution. The aquated

metal ions can donate a proton from one of their surrounding water molecules. This can

be an important factor contributing to the acidity of many soils.

Figure 15.8

In water a strong acid is virtually completely ionized but a weak acid is only slightly ionized

Strong Bases Strong bases are those which completely dissociate to produce hydroxide ions in aqueous

solution. For example, potassium hydroxide dissociates completely into potassium and

hydroxide ions according to the following equation.

KOH(s)→ K+(aq)+ OH

-(aq)

All group I and group II metal hydroxides are strong bases. In group II metal

hydroxides two moles of hydroxide ion are produced for every one mole of the

metal hydroxide which dissolves. For example, in a solution of barium hydroxide,

the dissociation equation is as follows.

BaOH)2 (s) →Ba2+

(aq)+ 2OH-(aq)

The group II metal hydroxides, however, have limited solubility. Magnesium

hydroxide is virtually insoluble, calcium hydroxide is slightly soluble and barium

hydroxide is still only moderately soluble. To extent that these compounds

dissolve they are completely dissociated and are therefore strong bases. However,

if a concentrated base solution is required, group I hydroxide would have to be

used.

The commonly used strong bases are listed in Table 15.9.

Table 15.9 Strong Bases

Name Formula Dissociation equation

Soluble

Sodium hydroxide NaOH NaOH(s) Na+(aq)+ OH

-(aq)

Potassium hydroxide KOH KOH(s) K+(aq)+ OH

-(aq)

Barium hydroxide Ba(OH)2 BaOH)2 (s) Ba2+

(aq)+ 2OH-(aq)

Low solubility

Magnesium hydroxide Mg(OH)2 Mg(OH)2 (s) Mg2+

(aq)+ 2OH-(aq)

Calcium hydroxide Ca(OH)2 Ca(OH)2 (s) Ca2+

(aq)+ 2OH-(aq)

Metallic oxides are also strong bases. In fact, the oxide ion is a stronger base than

hydroxide ion, as shown in Table 15.4. any soluble oxides such as sodium oxide (Na2O)

will therefore react completely with water to form a solution of the hydroxide.

Na2O + H2O(l)→ 2Na+

(aq) + 2OH-(aq)

Weak bases

Weak bases are substances in which only a small proportion of the molecules or ions

react with water to form hydroxide ions in aqueous solution. The weak base ammonia

reacts with water according to the following equation.

NH3(aq)+ H2O (l) NH4+

(aq) + OH-(aq)

Only a small fraction of ammonia molecules react and at equilibrium most of the

system exists as NH3 and H2O molecules. Using Bronsted-Lowry approach, NH4+

is a stronger acid (proton donor) than H2O and OH- is a stronger base (proton

acceptor) than NH3. Thus the reaction only occurs to a small extent.

Many weak bases are anions such as carbonate, ethanoate, fluoride and phosphate.

All these ions react with water to a small extent to produce hydroxide ions. For

example, carbonate ions react with water according to the equation:

CO32-

(aq) + H2O (l) HCO3-(aq) + OH

-(aq)

Although this reaction occurs only to a very limited extent, some hydroxide ions are

produced so the carbonate ions is acting as a weak base.

Review exercise 15.5

1. In the following pairs of solutions predict which solution would exhibit:

i. The higher [H+]

ii. The higher pH

iii. The greater electrical conductivity.

a. 0.1 mol L-1

HCl and 0.1 mol L-1 CH3COOH

b. 0.1 mol L-1

NaOH and 0.1 mol L-1 NH3

2. Give an example of each of the following

a. A concentrated solution of a strong acid

b. A dilute solution of a strong base

c. A concentrated solution of a weak base

d. A dilute solution of a weak acid

3. Write equations for the ionization or dissociation of the following.

a. The strong acid HBr

b. The weak acid H2SO3

c. The strong base RbOH

d. The weak base NaF

15.6 Acid ionization constants

The extent to which and an acid ionizes in aqueous solution can be determined from the

equilibrium constant for the ionization process. This equilibrium constant (Ka) is called

the acid ionization constant or the acid dissociation constant. In the ionization of

ethanoicacid:

CH3COOH(aq) H+ (aq) + CH3COO

- (aq)

The acid ionization constant (Ka) is given by:

Ka = ][

]][[

3

3

COOHCH

COOCHH

The Ka for ethanoic acid at 25ºC is 1.8 x 10-5

. This value indicates that reaction only

proceeds to a very limited extent. In a 0.1 mol L-1

solition only a little more than 1% of

ethanoic acid molecules are ionized. Table 15.10 contains the acid ionization constants

for some common acids. Note that the strong acids HCl, H2SO4 and HNO3 have no values

given for Ka. These acids are essentially completely ionized in aqueous solution and

would therefore have extremely large acid ionization constant.

Table 15.10 Acid ionization constants

Name Formula Ka

Hydrochloric acid HCl Large

Sulfuric acid H2SO4 Large

Nitric acid HNO3 Large

Oxalic acid H2C2O4 5.9 x 10-2

Sulfurous acid H2SO3 1.7 x 10-2

Hydrogensulfate ion HSO4- 1.2 x 10

-2

Phosphoric acid H3PO4 7.5 x 10-3

Iron (III) ion [Fe(H2O)6]3+

6.3 x 10-3

Hydrofluoric acid HF 7.2 x 10-4

Ethanoic acid CH3COOH 1.8 x 10-5

Aluminum ion Al3+

7.9 x 10-6

Carbonic ion H2CO3 4.2 x 10-7

Dihydrogenphosphate ion H2PO42-

6.2 x 10-8

Hydrogen sulfide H2S 1.0 x 10-7

Hypochlorous acid HClO 3.5 x 10-8

Ammonium ion NH4+ 5.6 x 10

-10

Hydrocyanic acid HCN 4.0 x 10-10

Hydrogencarbonate ion HCO3- 4.8 x 10

-11

Hydrogenphospate ion HPO42-

3.6 x 10-13

Hydrogensulfide ion HS- 1.3 x 10

-13

Calculation of [H+] and pH in a weak acid solution

From the value of Ka it is possible to calculate the hydrogen ion concentration and pH of

a solution of a weak acid.

Example 15.6

Calculate the [H+] and pH of a 0.10 mol L

-1 CH3COOH solution.

The equation for ionization is:

CH3COOH(aq) H+

(aq) + CH3COO- (aq)

If x moles of CH3COOH ionize per liter solution, the equilibrium concentration will be

[H+] = x, [CH3COO

-]=x, [CH3COOH]=0,10-x

Ka =1.8 x 10-5

= ][

]][[

3

3

COOHCH

COOCHH

= 1.8 x 10-5

= x

xx

10.0

..

1.8 x 10-5

= x

x

10.0

2

(1.8 x 10-5

) x (0.10-x) = x2

1.8 x 10-6

-1.8 x 10-5

x =x2

x2 + 1.8 x 10

-5 x-1.8 x 10

-6=0

Solving the quadratic equation for x

= 2

102.71024.3108.1 6105 xxx

=2

1068.2108.1 35 xx

=2

1066.2 3x

=1.33 x 10-3

[H+] = x = 1.3 x 10

-3 mol L

-1

pH = -log [H+]

= -log (1.33 x 10-3

)

= 2.89

A 0.10 mol L-1

CH3COOH solution therefore has a [H+] of only 0.0013 mol L

-1 and a pH

of approximately 3. In contrast, 0.10 mol L-1

HCl solution has a [H+] of 0.10 mol L

-1 and

a pH of 1. In general terms, acids with larger Ka values will have higher [H+], while for

acids with lower Ka values the [H+] will be lower.

1. Write out Ka expression for the following.

a. H3PO4 b. HClO c. HCO3 d. NH4+

2. Using the data table 15.10 arrange the following in order of decreasing strength

as acids:

CH3COOH, HS-, H2CO3, Fe(H2O)6

3+, H2SO4, NH4

+, H2PO4

3. For a 0.10 mol L-1

HF solution:

a. Calculate the [H+] and pH

b. Compare the [H+] and pH with that of 0.10 mol L

-1 CH3COOH and 0.10 mol

L-1 HCl solutions.

15.7 Polyprotic Acid and bases

Acids such as hydrochloric (HCl), nitric (HNO3) and ethanoic (CH3COOH) acid contain

only one acidic hydrogen atom per molecule which can be ionized in aqueous solution.

Despite the fact that hydrochloric and nitric acids are strong acids and ethanoic acid is a

weak acid, one mole of each of these acids will supply one mole of protons in their

reactions with sodium hydroxide. This is illustrated by the following equations.

OH- (aq) + HCl (aq) → H2O (l) + Cl

-(aq)

OH-(aq) + CH3COOH(aq)→H2O(l) + CH3COO

-(aq)

As hydrochloric, nitric and ethanoic acids only one acidic hydrogen atom per molecule of

acid, they are called monoprotic acids.

Other acids such as sulfuric (H2SO4) and carbonic (H2CO3) acid are diprotic, as they

contain two ionisable hydrogen atom per molecule of acid. In aqueous solutions of

sulfuric acid the first proton is completely ionized as shown by the following equation.

H2SO4 (l) + H2O (l) → H3O+ (aq) + HSO4

-(aq)

The hydrogensulfate ion is a weak acid, however, so that only a small proportion of these

ions ionize futher into hydrogen ions and sulfate ions. The equation for this reaction is as

follows.

HSO4-(aq) + H2O(l) H+ (aq) + SO4

2-(aq)

When it reacts with a strong base such as sodium hydroxide, one mole of sulfuric acid

will react with two moles of hydroxide ions.

2OH-(aq) + H2SO4(aq) → 2H2O (l) + SO4

2-(aq)

Phosphoric acid (H3PO4) is tripotic acid which contains three ionisable hydrogen atoms.

Phosphoric acid is weak acid so that ionization occurs to only a small extent. The

equations for the successive ionizations are as follows.

H3PO4 (aq) + H2O(l) H3O+(aq) + H2PO4

-(aq)

H2PO4- (aq) + H2O(l) H3O

+(aq) + HPO4

2-(aq)

HPO42-

(aq) + H2O(l) H3O+(aq) + PO4

3-(aq)

The acid ionization constant for these ionizations indicate that successive ionizations

occur to progressively smaller extents.

Because phosphoric acid is tripotic acid, one mole of acid will react with three moles of

sodium hydroxide as follows.

3OH- (aq) + H3PO4(aq) → 3H2O

+ (aq) + PO4

3-(aq)

It is often difficult to determine the number of acidic hydrogen atoms from a simple

molecular formula. For example, a molecule of ethanoic acid, CH3COOH, contains four

hydrogen atoms but only one of these is acidic. Oxalic acid, H2C2O4 or HOOCCOOH,

contains two hydrogen atoms and both of these are acidic. Generally

Figure 15.9

Acidic hydrogen atoms in ethanoic acid and oxalic acid

C C

H

H

H

O

C C

δO

O δ-O

-

δ+H

Hydrogen atoms are only acidic if they are attached to electronegative atoms, most

commonly oxygen. In CH3COOH only one hydrogen atom is bonded to an oxygen atom,

whereas in H2C2O4 both hydrogens are bonded to oxygen atoms.

Just as some acids are polyprotic it is possible for some bases to supply more than

one mole of hydroxide ion per mole of base. Sodium hydroxide (NaOH) and potassium

hydroxide (KOH) supply one mole of hydroxide per mole of base in aqueous solution.

One mole of these bases will react with one mole of hydrogen ions. Other bases such as

magnesium hydroxide (Mg(OH)2) and calcium hydroxide (Ca(OH)2) are able to supply

two moles of hydroxide ion per mole of base. One mol of magnesium hydroxide would

react with two moles of hydrogen ions according to the following equation.

Mg(OH)2 (s) +2H+(aq)→ 2H2O(l) + Mg

2+(aq)

Aluminium hydroxide and iron (III) hydroxide each contain three moles of hydroxide per

mole of base. One mole of these bases would react with three moles of hydrogen ions.

Review exercise

1. How many moles oh hydroxide ions would react with one mole of following

acids?

a. HNO3 b. H2SO4 c. H3PO4

d. HF e. CH3COOH f. HCOOH

2. How many moles of hydrogen ions would react with one mole of the following

bases?

a. NaOH b. Fe(OH)3 c. Ca(OH)2

3. Write equations for the successive ionizations of oxalic acid.

4. Using table 15.10 identify, in order of decreasing concentration, all the ions and

molecules present in the following solutions.

a. 1mol L-1

CH3COOH b.1mol L-1

H2SO4 c. 1mol L-1

H2S

15.8 Acid-base neutralization reactions

When solutions of acids and bases are mixed, reactions occur in which the acidic and

basic properties of the reactants are nullified. Such acid-base reactions are known as

neutralization reactions.

Most acid-base reactions result in the formation of a salt and water. The reaction

between a strong acid such as nitric acid and a strong base such as sodium hydroxide

produces water plus a salt, sodium nitrate, in solution. A non-ionic equation for this

reaction is:

HNO3(aq) + NaOH(aq)→NaNO3(aq) + H2O (l)

As nitric acid and sodium hydroxide are completely dissociated into ions in aqueous

solution, the reaction is better represented by the following equation.

δO δ+H δ+H O

Ethanoic acid Oxalic acid

Non-acidic hydrogen atoms acidic hydrogen atoms

15.7

H+(aq) + OH

-(aq)→H2O(l)

The sodium and nitrate ions are merely “spectator” ions which take no part in the

reactions but remain in the solution.

For a solution of a weak acid, such as ethanoic acid, the reaction with a

strong base such as sodium hydroxide is the best represented as :

OH-(aq) + CH3COOH(aq) H2O (l)+ CH3COO

- (aq)

The resulting solution is therefore one which contains sodium ethanoate ions ; that

is, a sodium ethanoate solution. Similar equations can be written for the reactions

of other weak acids with sodium hydroxide solution.

The reaction between hydrochloric acid and ammonia solution is an

example of a reaction between a strong acid and a weak base. As most of the

ammonia is present as molecules of NH3, and very little as NH4+ and OH

- ions,

the reaction is represented as follows.

NH3 (aq) + H+ NH4

+(aq)

The final solution is an ammonium chloride solution containing NH4 + and Cl

-

ions.

The reactions between sodium carbonate and hydrochloric acid is another

example of a reaction between a weak base and strong acid. The equation for this

reaction is as follows.

CO3 2-

(aq) + 2 H+(aq) H2CO3 (aq) CO2 (q) + H2O (l)

The reaction between a metallic oxide and acid is also a neutralization

reaction. For example insoluble copper (II) oxide will react with hydrochloric acid

to form a solution of copper (II) chloride.

Cu(s) + 2H+ Cu

2+ (aq) + H2O (l)

Review exercise 15.8

1. Write appropriate equations for the reactions between sodium hydroxide

solution and each of the following.

a. Hydrochloric acid

b. Sulfuric acid

c. Carbonic acid

d. Gaseous carbon dioxide

2. Write equation for the reaction between hydrochloric acid and each of the

following.

a. Calsium hydroxide solution

b. Solid aluminium hydroxide

c. Solid magnesium axide

d. Solid sodium carbonate

15.9 Salts

Sodium chloride (NaCl), common table salt, is a member of a class of

compounds called salt. A salt is an ionic compound containing a cation other than

H+ and an anion other than OH

- or O

2-. Pottasium nitrate (KNO3), magnesium

chloride (MgCl2), sodium sulfate (Na2SO4), copper (II) carbonate (CuCO3) and

ammonium ethanoate (NH4CH3COO) are examples of salt.

Another way of thinking about salts is to consider them as being formed

by the replacement of hydrogen in an acid by a metal ion or an ammonium ion.

Thus salts derived from hydrochloric acid are chlorides, from sulfuric acid are

sulfates, from nitric acid are nitrates, from carbonic acid are carbonates and so

In dilute aqueous solutions, salts are ompletely dissociated ions. Manesium

nitrate, for example, is cmpletely dissociated to the following equation.

Mg(NO3)2(s) Mg2+

(aq) + 2NO3-(aq)

Similarly, a solution of potassium chloride consists of K+ and Cl

- ions, an

ammonium ethanoate solution contains NH4+ and CH3COO

- ions and so on.

Acid-base properties of salt solutions

Aqueous solutions of salts can be acidic, neutral or basic depnding on the

particular ions in the salt. For examples, solutions of NaCl and Ca(NO3)2 are

neutral and have a pH of 7. In contras of NaCH3COO and KF are basic (pH > 7),

while solutions of NH4Cl and NaHSO4 are acidic (pH < 7), solutions of salts can

be acidic or basic when the ions in the salt react with water to produce H+(aq) or

OH-(aq), this type of reaction is often called hydrolysis.

In order to predict the acid-base nature of a salt it is necesary to consider

the acid-base prooperties of the individual ions making up the salt, the acid-base

properties of cation and anions are summarised in table 15.11.

Table 15.11 Acid-base properties of some common ions in aqueous solution.

Neutral Basic Acidic

Anion

derived

from

srong

acids

Cl-, NO3

-, Br

-, I

_

derived

from

weak

acids

F-, S

2-, SO4

2-, ClO

-,

CH3COO-, CO3

2-,

HCO3-, PO4

3-, HPO4

2-

derived

from

polyprotic

acids

HSO4-, H2PO4

2-

Cation

derived

from

strong

bases

Li+, Mg

2+, Na

+,

Ca2+

, K+, Ba

2+

none

NH4+

Al3+

Fe3+

The information in table 15.11 illustrates the following generalisations

about the acid-base properties of ions in aqueous solutions.

1. Neutral anions ore those derived from strong acids. Anions such as chloride

have no tendency to rect with water from HCl and hydroxide ion.

2. Neutral cations are group I and II cations. These cations are derived from

strong bases such as sodium hydroxide adn magnesium hyydroxide.

3. Basic anion are those which react with water to from some hydroxide ions in

aqueous solution. Basic anions are derived from week acids, for example, the

ethanoate ions react with water to produce ethanoic acid and hydroxide ion

according to the following equation.

CH3COO-(aq) + H2O(l) CH3COOH (aq)+ OH

– (aq)

This reaction occurs to only a limited extent. Therefore in a solution

containing ethanoate ions only a small proportion of these ion react with water

to form ethanoic acid and hydroxide ion. The reaction occurs to only a limited

extent because CH3COOH is stronger acid than H2O and OH – is a stronger

base than CH3COO-. Thus solution containing ethanoate ions are weakly

basic.

4. Acidic anions are those which contain hydrogen atoms which can transfer to

water molecules to form hydronium ions. Acidic anions are derived from

polyprotic acids. For instance, the hydrogensulfate ion HSO4- is derived from

sulfuric acid H2SO4 and the dihydrogensulfate ion H2PO4 from phosphoric

acid.

Solution containing the hydrogensulfate ion are acidic due to the following

reaction.

HSO4- (aq) + H2O(l) H3O

+(aq)

+ SO4

2-(aq)

Competing hydrolisis reactions

Anion which contain acidic hydrogen atoms and are derived from polyprotic

acids often undergo separate hydrolysis reaction which produce both H+ (aq)

and OH –

(aq). Examples are the H2PO4 -

and HCO3-

ions

with water as

follows.

H2PO4 -

(aq) + H2O(l) HPO4 2-

(aq) + H3O+(aq)

H2PO4 -

(aq) + H2O(l) HPO4 2-

(aq) + OH –

(aq)

HCO3 -

(aq) + H2O(l) CO3 2-

(aq) + H3O+(aq)

HCO3 - (aq) + H2O(l) H2CO3

(aq) + OH

– (aq)

Whether the anion exhibits acidic or basic properties in solution depends on

the relative tendencies of these competing hydrolysis reactions. H2PO4 –

is an

acidic anion because it is a better proton donor than proton acceptor.

5. Acidic cations can be classified as those derived from weak bases and aquated

metal ions. The ammonium ions (NH4+)is derived from the weak base

ammonia (NH3). The ammonium ions is acidic in aqueous solution because of

the following reactions.

NH4+(aq) + H2O(l) NH3(aq) + H3O

+(aq)

Again, this reaction only occurs to a limited extent as H3O+ is a stronger acid

than NH4+

and NH3 is a stronger base than H2O. Thus solutions containing

ammonium ions are only weakly acidic.

Small, highly charged metal ions are also acidic in aqueous solution. For

example, the aquated aluminium ion, [Al(H2O)6]3+

, produce hydronium ions

according to the reactions:

[Al(H2O)6]3+

(aq) + H2O(l) [Al(OH)(H2O)5]2+

(aq) + H3O+(aq)

Similarly, [Fe(H2O)6]3+

and other transition metal ions produce weakly acidic

solutions.

Predicting The Acid-Base Properties Of Salt Solutions

To determine whether a particular salt will undergo hydrolysis to form an acidic

or basic solution it is necessary to the cation and anion concerned. For example, a

NaCl solution will be neutral as both the Na+ + Cl

- have no tendency to react with

water to produce H3O+ or OH

-. in solution of KCH3COO and Na2S the CH3COO

-

and S2-

ions react with water to produce small amounts of hydroxide ion and

hence basic solution result. Solutions such as NH4Cl and Fe(NO3)3 contain the

cation NH4+

and [Fe(H2O)6]

3+ which produce small amounts of hydronium ion

and therefore weakly acidic solutions.

For solutions of salts such as ammonium ethanoate where one ion has acidic

properties and the other basic properties, the acidity or basidity of the solutions

will depend on the relative effect of the two ions. In ammonium ethanoate

solution the acidity due to the ion is approximately equal to the basicity of the

CH3COO- ion. As a result an NH4CH3COO solution is virtually neutral. A

solution of NH4CN however, is slightly basic as the cyanide ion has stronger

basic properties than the NH4+ions acidic properties.

Review exercise 15.9

1. Write equations for the hydrolysis of the following ions.

a. S2-

b. CO3 2-

c. NH4+

d. [Fe(H2O)6]3+

e. F-

f. HSO4

g. ClO-

h. CH3

2. Classify solution of the following as acidic, basic, or neutral.

a. KNO3

b. NH4NO3

c. Ca(ClO)2

d. Na3PO4

e. AlBr3

f. KH2PO4

g. MgI2

h. NaCO3

15.10 Acids and bases in the living world

Carbon dioxide is an important gas in the biological world. In photosynthesis it is

used by plantsas the primary source of food for living systems.

6CO2 (g) + 6H2O(l) + sunlight C6H12O6 (aq)+ 6O2 (g)

it is also a product ofrespiration through which living things gain the energy they

need to function.

C6H12O6 (aq)+ 6O2 (g) 6CO2 (g) + 6H2O(l) + energy

Carbon dioxide is slightly soluble in water and establisheds an equillibrium

between gaseous and dissolved CO2 as follows.

CO2 (g) CO2 (aq)

The dissolved CO2 react with water to form carbonic acid which is weak diprotic

acid

CO2 (g) + H2O(l) H2CO3 (aq)

The carbonic acid establishes equilibria involving the hydrogencarbonate and

carbonate ion.

H2CO3 -

(aq) + H2O(l) HCO3- (aq) + H3O

+(aq)

HCO3 - (aq) + H2O(l) H3O

+(aq) + CO3

2- (aq)

Because of these equilibria a solution of carbon dioxide is weakly acidic.

Rainwater, for example, has a pH about %.6. soda water is a supersaturated solution

of carbon dioxide and has a characteristic sour acid taste.

Within natural aquatic systems such as lakes, rivers, and thev ocean the CO2 /

H2CO3 / HCO3 -

/ CO3 2-

system provides

the carbon dioxide needed for

photosynthesis by aquatic plants and has a role in pH control.

Within the human body CO2 released by respiration processes is transported by the

blood to the lung where it is exhaled. The pH human blood needs to be in the range

7.35-7.45. The CO2 / H2CO3 / HCO3 -

/ CO3 2-

system plays an important role in

keeping the blood pH within this range.

The human stomach produces about 2-3 litres of gastric juice daily. This has a pH

af about 0.8 due to the presence of hydrochloric acid which is secreted by cells in the

stomach wall. The HCl helps to suppress the growth of bacteria in the stomach and

provides ideal pH conditions which asssist the enzyme pepsin, to digest protein

Normally the stomatch is not harmed by the HCl it contains because of the

presence of an inner protective lining. However, if too much HCl is produced this

may cause discomfort and, in severe cases, may lead to stomach ulcers. Antacids are

sometimes used to neutralise some of the HCl in the stomach and to decrease the

acidity. Several bases, incuding magnesium hydroxide, aluminium hydroxide,

calsium carbonate and sodium hydrogencarbonate are used in antacid preparation.

This react with hydrochloric acid in the stomach as follows.

Mg(OH)2 (s) + 2H+ Mg

2+ (aq) + 2H2O (l)

Al(OH)3 (s) + 3H+ Al

3+ (aq) + 3H2O (l)

CaCO3 (s) + 2H+ Ca

2+ (aq) + CO2 (g) + 2H2O (l)

NaHCO3 (s) + H+ Na

+ (aq) + CO2 (g) + 2H2O (l)

Figure 15.10

a. Cells in the stomach lining secreate gastric juice containing hydrochloric acid

b. Antacids are often used to reduce excess stomach acidity (June oxford)

Review exercise 15.10

1. For a solution of carbon dioxide identify the following

a. All the species present

b. The most abundant ionic species

2. Gastricis approximately 0.15 mol L-1

HCl. Calculate the volume of gastric juice

which would be neutralized by an antacid table containing 750 mg of CaCO3.

15.11 Acids, bases and the periodic table

Trends across periods

The acid base properties of the oxides and hydroxides of the period 3 elements are

shown in table 15.12

From the table it can be seen that the oxides and hydroxides have increasingly

acidic character going from left to right across the period. Sodium and magnesium

oxides and hydroxides are strong bases. Aluminium oxide and hydroxide are

amphoteric. They will dissolve in solutions of strong acids or strong bases. When

aluminium hydroxide dissolves in acid solution it is acting as a base.

Al(OH)3 (s) + 3H+ (aq) Al

3+ (aq) + 3 H2O(l)

When it dissolves in basic solution the aluminium hydroxide is acting as an acid

Al(OH)3 (s) + OH- [Al(OH)4]

- (aq)

The oxides and hydroxides of silicon, phosphorus, sulfur and chlorine are all

acidic but vary in strength from weakly acidic for silicon dioxide to strongly acidic

for sulfuric acid and perchloric acid.

Table 15.12 the acid base properties of the oxides and hydroxides of the period 3

elements.

Elements Na Mg Al Si P S Cl

Oxides

Formula

Product of

reaction

With water

Na2O

NaOH

MgO

insoluble

Al2O3

insoluble

SiO2

insoluble

P4O10

H3PO4

SO3

H2SO4

Cl2O7

HCO4

Hydroxide

Formula

NaOH

Mg(OH)2

Al(OH)3

Si(OH)4

PO(OH)3

SO2(OH)2

ClO3(OH)

Product of

reaction

With water

Soluble

Na+, OH

-

insoluble

insoluble

or H4SiO4

insoluble

or H3PO4

Soluble

mainly

molecules

Soluble,

H+ HSO4-

Or HClO4

Soluble

H+, ClO4

-

Reactions of

oxides/

hydroxides

Products of

reaction

With strong

acid

Products of

reaction

With strong

base

H2O

+ Na+

No

Reaction

H2O

+ Mg2+

No

Reaction

H2O

+ Al3+

[Al(OH)3]

No

Reaction

H2O

+SiO32-

No

Reaction

H2O

+PO4 3-

No

Reaction

H2O

+SO4 2-

No

Reaction

H2O

+ClO4-

Acid- base

properties of

oxides/

hydroxides

Stongly

Basic

Stongly

Basic

amphoteric Weakly

Acidic

Moderatel

y

acidic

Strongly

acidic

Strongly

acidic

Bonding in

solid oxide

ionic Ionic ionic Covalent

network

Covalent

molecular

Covalent

molecular

Covalent

molecular

Trend down groups

Trend in the acid-base properties of oxides down a group can be illustrated by the

group elements. These are shown in table 15.13. from the table it is apparent that

there is an increase in the basicity of the oxides going down the group. Carbon

dioxide is acidic in nature, producing carbonic acid when dissolved in water. Silicon

dioxide and germanium dioxide have successively less acidic properties. Tin and lead

dioxide are amphoteric in character, and hence will dissolve in solutions of strong

acid and strong bases.

Table 15.13 the acid-base properties of the oxides of the group IV elements

Elements C Si Ge Sn Pb

Formula of

oxide

CO2 SiO2 GeO2 SnO2 PbO2

Formula of

reaction

with water

Slightly

soluble

H2CO3

Insoluble Insoluble Insoluble Insoluble

Formula of

reaction

with strong

acid

no reaction No reaction no reaction H2O

+ Sn 4+

H2O

+ Pb 4+

Formula of

reaction

with strong

base

H2O

+ CO3 2-

H2O

+ SiO3 2-

H2O

+ GeO3 2-

[Sn(OH)6]2-

[Pb(OH)6]2-

Acid-base

properties

Moderately

acidic

Weakly

acidic

Weakly

acidic

amphoteric amphoteric

Bonding in

solid oxide

Covalent

molecular

Covalent

network

Covalent

network

ionic ionic

Periodic trends related to electronegativities and bonding

The oxides and hydroxides of elements therefore tend to increase in acidity across a

period but decrease in acidity down a group in the periodic table. These trends can be

understood in terms of the electronegativities of the elements increase gradually

across the periode from sodium to chlorine. This indicates that elements to the right

of the periode have a greater tendency to attract electrons than those at the left. In

sodium oxide the bonding is ionic as sodium has a low electronegativity and oxygen

has a high value. The electrons are completely transferred from the sodium atoms to

the oxygen has a high value. The electrons are completely transferredfrom the sodium

atoms to form Na+ and O

2- ions. When dissolved in waater the O

2- ion, being a strong

base, react to form hydroxide ions.

Figure 15. 11 Acid base properties of main group oxides

Basic oxide

Acidic oxide

Amphoteric acid

Li2O BeO B2O3 CO2 N2O5 OF2

Na2O MgO Al2O3 SiO2 P4O11 SO3 Cl2O7

K2O CaO Ga2O3 GeO2 As2O5 SeO3 Br2O7

Rb2O SrO In2O3 SnO2 Sb2O5 TeO3 I2O7

Cs2O BaO Tl2O3 PbO2 Bi2O5 PoO3 At2O7

In SO3 and Cl2O7 however, the difference in electronegativities of the

elements involved is not nearly so pronounced. In SO3 the bonding can be

represented as shown in the side column. The bonding is covalen and slightly polar.

When SO3 is dsissolved in water the sulfur atom, with a partial positive charge,

accepts a hydroxide ion as follows.

O -

S +

O - O -

O O -

S + H2O S + H+

O O O O OH

This results in the production of H+ and HSO4

– ions as found in a sulfuric acid

solution .

The decreasing acidity of oxides down a group can also be explained in terms

of trends in the elektronegativities of the elements within the group. There is gradual

decrease in elektronegativity down a group as is evident from the increasing metallic

nature of the elements in group IV. Consequently, the oxides will tend be more ionic

and display increasingly basic properties going down the group.

Review exercise 15.11

1. K2O, Ga2O3 and Br2O7 are three oxides from the fourth row of the periodic table.

Predict:

a. The nature of the bonding in these oxides

b. The acid –base properties of the oxides

2. Alumunium, boron and indium are all elements in group III. Al2O3 is an

amphoteric oxide. From the positions of Al, B and In in the periodic table, predict

the acid base characteristics of B2O3 and In2O3.

Major ideas

Properties of aqueous solutions of acids:

1. Turn litmus red

2. Conduct an electric current.

3. Taste sour

4. React with reactive metals to produce hydrogen gas .

5. React with carbonates and hydrogencarbonates to form carbon dioxide.

6. React with metal oxides to produce a salt and water.

7. React with metal hydroxides to form a salt and water.

Properties of aqueous solutions of bases:

1. Turn litmus blue

2. Conduct an electric current.

3. Taste bitter

4. React with amphoteric metals to produce hydrogen gas.

5. React with acid to form a salt and water.

6. Dissolve amphoteric metals hydroxides.

Arrhenius theory of acid and bases:

1. an acid is a substance which produces H+ (aq) in solution.

2. an base is a substance which produces OH- (aq) in solution.

Bronsted – lowry theory of acid and bases:

1. An acid is a proton donor

2. A base is a proton acceptor

3. Some substances can react as acids or basis in different reactions.

4. Every acid has a conjugate base which is related in the following way.

HX H+ + X

-

Acid proton conjugate base

5. The sronger an acid, the weaker its conjugate base.

6. Acid-base reactions tend to occur in the diretion in which a stronger acid

(proton donor) and stronger base (proton acceptor) react to form a weaker acid

and weaker base.

Water is a weak electrolyte which ionises to form H+ (aq) and OH

-(aq) to a small

extent.

The relationship between the concentration H+ (aq) and OH

-(aq) in any aqueous

solution at 25 0C is given by the following.

Kw= [H+] [OH

-] = 1.0 x 10

-14

The pH of any solution is calculated from :

pH = - log 10 [H+]

Table 15.14

pH Solution

>7

=7

<7

Basic

Neutral

Acidic

The relationship between pH and acidity is shown the table 15.14.

Strong acids (for example, HCl, H2SO4) and strong bases (for example, NaOH)

are essentially completely ionized or dissociated into ions in aqueous solution.

Weak acids (for example, CH3COOH, H3PO4) and weak bases (for example,

NH3) are those in which only a small proportion of the molecules are ionised in

aqueous solution.

The acid ionisation constant (Ka) is the equilibrium constant for the ionisation of

an acid (HX) into ions.

Ka = [H+][X

-]

[HX]

Polyprotic acids (for example, H2SO4, H3PO4 ) are those which contain more than

one acidic or ionisable hydrogen for formula unit of the acid

For polyprotic acids successive ionisation occur to successively smaller extents.

Some bases (for example, Ca(OH)2) contain more than one hydroxide ion per

formula unit of the base.

A neutralisation reaction is the reaction between an acid and a base.

Types of neutralisation reactions are shown in table 15.15

Table 15.15

Reaction Equation

Strong acid/strong base

Weak acid/strong acid

Strong acid/ weak base

Strong acid/insoluble metal oxide

H+ (aq) + OH

-(aq) H2O(l)

HX (aq) + OH-(aq) H2O(l) + X

-(aq)

H+ (aq) + B (aq) BH

+(aq)

2H

+ (aq) + MO(s) H2O(l) + M

2+ (aq)

Salts are ionic compounds containing a cation other than H+

and anion other that

O2-

or OH- .

Nearly all salts are strong electrolytes.

Salt solution can be acidic, basic or neutral depending on the tendencies of the

ions in the salt to undergo hydrolysis. The acid-base properties of ions are

summarized in table 15.16

Table 15.16

Neutral Basic Acidic

Anion Derived from strong

acids e.g. Cl-, NO3

-.

Derived from

weak acids e.g.

CO32-

, CH3COO-

Derived from

polyprotic acids e.g.

HSO4-, H2PO4

-

Cation Derived from strong

bases e.g. Na +, Ca

2+

None NH4+, Al

3+, transition

metal ions e.g. Fe 3+

Carbon dioxide is slightly soluble in water and dissolves to form carbonic acid,

H2CO3. this is a weak diprotic acid which undergoes successive ionisation

reaction to form hydrogencarbonate, HCO3- and carbonate, CO3

2- , ions.

Oxides and hydroxides increase in acidity across periods in the periodic table.

Oxides and hydroxides decrease in acidity down groups in the periodic table.

Question and problems

1. Write balanced chemical equations for the reaction between the following

a. Aluminium and hydrochloric acid.

b. Solid potassium hydrogencarbonate and nitric acid

c. Iron (III) oxide and sulfuric acid.

d. Barium hydroxide solution and hydrofluoric acid.

2. Write balanced chemical equations for the reactions between the following .

a. Sodium hydroxide solution and aluminium

b. Potassium hydroxide solution and Iron (III) hydroxide

c. Potassium hydroxide solution and zinc hydroxide

d. Sodium hydroxide solution and sulfur trioxide

3. Write balanced chemical equations for the following reactions.

a. Baking powder, which includes potassium hydrogen tartrate (KHC4H4O6) and

sodium hydrogencarbonate (NaHCO3), is moistened to allow reaction

between the components.

b. Rust (Fe2O3) on a wrought iron gate is trated with spirits of salt (HCl) to

remove the rust.

c. An antacid powder containing aluminium hydroxide is used to neutralise

excess stomach acidity (HCl).

d. Acid rain containing sulfurous acid (H2SO3) react with marblre (CaCO3)

statues in a city park.

4. Write the formulas for the following

a. The conjugate bases of HClO3, HS-, HH4

+, H2O.

b. The conjugate acids of HCO3-, HS

-, [Fe(OH)(H2O)5]

2+ , N2H4.

5. Write equations for the reactions which occur when the following are dissolved in

water

a. The strong acid HClO4

b. The strong base LiOH

c. The weak acid HCOOH

d. The weak base N2H4

6. Draw diagrams such as those shown in figure 15.8 to represent:

a. A concentrated solution of a strong acid

b. A dilute solution of a strong acid

c. A concentrated solution of a weak acid

d. A dilute solution of a weak acid

7. For each of the following reactions :

i. Identify the Bronsted-Lowry conjugate acid-base pairs

ii. Predict whether the reaction will occur to a small a large extent.

a. H2C2O4(aq) + H2O (l) H3O+ (aq) + HC2O4

-(aq)

b. H2O(l) + CN-(aq) HCN (aq) + OH

- (aq)

c. CH3COOH (aq) + S2-

(aq) CH3COO– (aq) + HS

- (aq)

d. HCl(aq) + F- (aq) HF(aq) + Cl

-(aq)

8. Calculate [H+], [OH

-] and pH of the following solutions

a. 0.50 mol L-1

HBr

b. 3.0 x 10-3

mol L-1

Ca(OH)2

c. a solution prepared by dissolving 1.60 gr of NaOH in water to make up 250.0

mL of solution

d. a solution prepared by diluting 2.0 mL of 12 mol L-1

HCl to form 250 mL of

solution.

9. Calculate the [H+] and [OH

-] concentration in each of the following

a. Coca cola with a pH of 3.00

b. Acid rain with pH of a 2.40

c. Baking soda solution with a pH 8.50

d. Dishwasing detergent with a pH of 12.10

10. For a 0.20 mol L-1

HNO3 solution calculate the following

a. The concentration of H+ and NO3

- ions and HNO3 molecules

b. The pH

11. For a 0.20 mol L-1

HNO2 solution (Ka = 4.5 10-4

) calculate the following

a. The concentration of H+ and NO2

- ions and HNO2 molecules

b. The pH

12. Without doing detailed calculations, use the acid ionisation constants to list the

following 0.1 mol L-1

solutions in order of increasing pH.

HCl, NH3, Na2CO3, H3PO4, NaCl, H2SO4, NaOH, NH4Cl, Ba(OH)2

13. For the following acid:

O H O H

C C C

H O O

a. Identify the number of hydrogen atoms

b. Identify the number of acidic hydrogen atoms

c. Determine the number of moles of hydroxide ions required to neutralise one

mole of the acid

d. Write an equation for the neutralisation reaction.

14. 1L of 2 mol L-1

NaOH is added to 1 L of 1 mol L-1

H2SO4. identify :

i. All of the spesies present

ii. The most abundant spesies other than water

a. In the separated solution before they are mixed

b. When half the NaOH solution has been added to the H2SO4 solution

c. When all the NaOH solution has been added to the H2SO4 solution

15. a. Write an equation for the hydrolysis of ethanoate ion (CH3COO-)

b. Explain why a solution of sodium ethanoate is basic when the hydrolysis of

ethanoate ion forms ethanoic acid molecules.

16. a. Write equations to represent the HCO3– ion acting as an acid and as a base in

aqueous solution.

b. Explain why HCO3– ion is classified as a basic anion.

17. Classify each of the following solution as acidic, neutral or basic. Where a

solution is not neutral write an equation far the hydrolysis reaction involved.

a. CaCl2

b. Cr(NO3)3

c. Na3PO4

d. Na2SO4

e. K2CO3

f. KCN

g. NH4Br

h. (NH4)2S

18. Write equation to represent the chemical equilibria between:

a. Gaseous and aqueous carbon dioxide

b. Aqueous carbon dioxide and carbonic acid

c. Carbonic acid and hydrogencarbonate ion

d. Hydrogencarbonate ion and carbonate ion.

19. An antacid tablet contains750 mg CaCO3 and 200 mg Al(OH)3.

a. Write equation to represent the reactions between CaCO3 and Al(OH)3 and gastric

juice acid, HCl.

b. Assuming gastric juice acid is 0.15 mol L-1

HCl, Calculate the volume of gastric juice

acid which would be neutralised by the antacid tablet.

20. Arrange the following oxides in order of increasing acidity.

a. Al2O3, Na2O, P4O11, Cl2O7

b. SiO2, SO3, K2O, Al2O3, Cl2O7