Assistant Lecture Aayad Amaar Acids and Bases The properties of ...

College of Dentistry 

Inorganic Chemistry 

Assistant Lecture 

Aayad Amaar 

Acids and Bases 

The properties of acids and bases are related to their chemical 

structure. All acids have common characteristics that enable them 

to increase the hydrogen ion concentration in water. All bases 

lower the hydrogen ion concentration in water. Two theories, one 

developed from the other, help us to understand the unique 

chemistry of acids and bases. 

Arrhenius Theory of Acids and Bases 

One of the earliest definitions of acids and bases is the 

Arrhenius theory. According to this theory, an acid, dissolved in 

water, dissociates to form hydrogen ions or protons (H + ), and a 

base, dissolved in water, dissociates to form hydroxide ions (OH - 

). For example, hydrochloric acid dissociates in solution 

according to the reaction 

Sodium hydroxide, a base, produces hydroxide ions in solution: 

The Arrhenius theory satisfactorily explains the behavior of 

many acids and bases. However, a substance such as ammonia, 

NH3, has basic properties but cannot be an Arrhenius base, 

because it contains no OH - . The Br? nsted-Lowry theory 

explains this mystery and gives us a broader view of acid-base 

theory by considering the central role of the solvent in the 

dissociation process. 

Br?nsted-Lowry Theory of Acids and Bases 

The Br?nsted-Lowry theory defines an acid as a proton (H + ) 

donor and a base as a proton acceptor. Hydrochloric acid in 

solution donates a proton to the solvent water thus behaving as a 

Br?nsted-Lowry acid: 

H3O + is referred to as the hydrated proton or hydronium ion.



The basic properties of ammonia are clearly accounted for by 

the Br?nsted-Lowry theory. Ammonia accepts a proton from the 

solvent water, producing OH - . An equilibrium mixture of NH3, 

H2O, NH4 + , and OH - results. 

For aqueous solutions, the Br?nsted-Lowry theory adequately 

describes the behavior of acids and bases. We shall limit our 

discussion of acid-base chemistry to aqueous solutions and use 

the following definitions: 

An acid is a proton donor. 

A base is a proton acceptor. 

Conjugate Acids and Bases 

The Br?nsted-Lowry theory contributed several fundamental 

ideas that broadened our understanding of solution chemistry. 

First of all, an acid-base reaction is a charge-transfer process. 

Second, the transfer process usually involves the solvent. Water 

may, in fact, accept or donate a proton. Last, and perhaps most 

important, the acid-base reaction is seen as a reversible process. 

Consequently, any acid-base reaction can be represented by the 

general equation 

In the forward reaction, the acid (HA) donates a proton (H + ) to the 

base (B) leading to the formation of BH + and A - . However, in the 

reverse reaction, it is the BH + that behaves as an acid; it donates 

its extra proton to A - . A - is therefore a base in its own right 

because it accepts the proton. These product acids and bases are 

termed conjugate acids and bases. 

A conjugate acid is the species formed when a base accepts 

a proton. 

A conjugate base is the species formed when an acid 

donates a proton.



The acid and base on the opposite sides of the equation are 

collectively termed a conjugate acid-base pair. In the above 

equation: 

BH + is the conjugate acid of the base B. 

A - is the conjugate base of the acid HA. 

B and BH + constitute a conjugate acid-base pair. 

HA and A - constitute a conjugate acid-base pair. 

Rewriting our model equation: 

Although we show the forward and reverse arrows to indicate 

the reversibility of the reaction, seldom are the two processes 

equal but opposite. One reaction, either forward or reverse, is 

usually favored. Consider the reaction of hydrochloric acid in 

water: 

HCl is a much better proton donor than H 3 O + . Consequently the 

forward reaction predominates, the reverse reaction is 

inconsequential, and hydrochloric acid is termed a strong acid. 

The dissociation of hydrochloric acid is so favorable that we 

describe it as 100% dissociated and use only a single forward 

arrow to represent its behavior in water: 

The degree of dissociation, or strength, of acids and bases has a 

profound influence on their aqueous chemistry. For example, 

vinegar (a 5% [w/v] solution of acetic acid in water) is a 

consumable product; aqueous hydrochloric acid in water is not. 

Why? Acetic acid is a weak acid and, as a result, a dilute solution 

does no damage to the mouth and esophagus. The following 

section looks at the strength of acids and bases in solution in more 

detail.



Acid-Base Properties of Water 

The role that the solvent, water, plays in acid-base reactions is 

noteworthy. In the example above, the water molecule accepts a 

proton from the HCl molecule. The water is behaving as a proton 

acceptor, a base. However, when water is a solvent for ammonia 

(NH 3 ), a base, the water molecule donates a proton to the 

ammonia molecule. The water, in this situation, is acting as a 

proton donor, an acid. Water, owing to the fact that it possesses 

both acid and base properties, is termed amphiprotic. The 

solvent properties of water are a consequence of this ability to 

either accept or donate protons. Water is the most commonly used 

solvent for acids and bases. These interactions promote solubility 

and dissociation of acids and bases 

Acid and Base Strength 

The terms acid or base strength and acid or base concentration 

are easily confused. Strength is a measure of the degree of 

dissociation of an acid or base in solution, independent of its 

concentration. Concentration, as we have learned, refers to the 

amount of solute (in this case, the amount of acid or base) per 

quantity of solution. The strength of acids and bases in water 

depends on the extent to which they react with the solvent, water. 

Acids and bases are classified as strong when the reaction with 

water is virtually 100% complete and as weak when the reaction 

with water is much less than 100% complete. 

Important strong acids include: 

Note that the equation for the dissociation of each of these 

acids is written with a single arrow. This indicates that the 

reaction has little or no tendency to proceed in the reverse 

direction to establish equilibrium. All of the acid molecules are 

dissociated to form ions. 

All common strong bases are metal hydroxides. Strong bases 

completely dissociate in aqueous solution to produce hydroxide



ions and metal cations. Of the common metal hydroxides, only 

NaOH and KOH are soluble in water and are readily usable strong 

bases: 

Weak acids and weak bases dissolve in water principally in the 

molecular form. Only a small percentage of the molecules 

dissociate to form the hydronium or hydroxide ion. 

Two important weak acids are: 

The double arrow implies an equilibrium between dissociated 

and undissociated species. We have already mentioned the most 

common weak base, ammonia. Many organic compounds 

function as weak bases. Several examples of weak bases follow: 

The fundamental chemical difference between strong and weak 

acids or bases is their equilibrium ion concentration. A strong 

acid, such as HCl, does not, in aqueous solution, exist to any 

measurable degree in equilibrium with its ions, H 3 O + and Cl - . On 

the other hand, a weak acid, such as acetic acid, establishes a 

dynamic equilibrium with its ions, H 3 O + and CH3COO - . The 

relative strength of an acid or base is determined by the ease with 

which it donates or accepts a proton. Acids with the greatest 

proton-donating capability (strongest acids) have the weakest 

conjugate bases. Good proton acceptors (strong bases) have weak 

conjugate acids. 

Solutions of acids and bases used in the laboratory must be 

handled with care. Acids burn because of their exothermic 

reaction with water present on and in the skin. Bases react with 

proteins, which are principal components of the skin and eyes. 

Such solutions are more hazardous if they are strong or 

concentrated. A strong acid or base produces more H3O + or OH - 

than does the corresponding weak acid or base. More-



concentrated acids or bases contain more H 3 O + or OH - than do 

less concentrated solutions of the same strength. 

The Dissociation of Water 

Aqueous solutions of acids and bases are electrolytes. The 

dissociation of the acid or base produces ions that can conduct an 

electrical current. As a result of the differences in the degree of 

dissociation, strong acids and bases are strong electrolytes; weak 

acids and bases are weak electrolytes. The conductivity of these 

solutions is principally dependent on the solute and not the 

solvent (water). 

Although pure water is virtually 100% molecular, a small number 

of water molecules do ionize. This process occurs by the transfer 

of a proton from one water molecule to another, producing a 

hydronium ion and a hydroxide ion: 

This process is the autoionization, or self-ionization, of water. 

Water is therefore a very weak electrolyte and a very poor 

conductor of electricity. Water has both acid and base properties; 

dissociation produces both the hydronium and hydroxide ion. 

Pure water at room temperature has a hydronium ion 

concentration of 1.0x10 -7 M. One hydroxide ion is produced for 

each hydronium ion. Therefore, the hydroxide ion concentration 

is also 1.0x10 -7 M. Molar equilibrium concentration is 

conveniently indicated by brackets around the species whose 

concentration is represented: 

The product of hydronium and hydroxide ion concentration in 

pure water is referred to as the ion product for water. 

The ion product is constant because its value does not depend on 

the nature or concentration of the solute, as long as the 

temperature does not change. The ion product is a temperature-



dependent quantity. The nature and concentration of the solutes 

added to water do alter the relative concentrations of H 3 O + and 

OH - present, but the product, [H 3 O + ][OH - ], always equals 1.0x10 - 

14 

at 25 °C. This relationship is the basis for a scale that is useful 

in the measurement of the level of acidity or basicity of solutions. 

This scale, the pH scale.

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