Basic definitions:

You might recall the Arrhenius definition of acids and bases (Module I):

Acid = proton (H⁺) donor

Base = hydroxide ion (OH^-) acceptor

The Arrhenius definition is sufficient to describe the reaction of strong acids and bases with water. However, it fails to describe the reaction of weak bases in water since weak bases do not have a hydroxide ion.
 

Strong acids = strong electrolytes:	Strong bases = strong electrolytes
Hydrochloric acid HCl	Sodium hydroxide (NaOH)
Hydrobromic acid HBr	Lithium hydroxide (LiOH)
Hydroiodic acid (HI)	Potassium hydroxide (KOH)
Perchloric acid (HClO4)	Barium hydroxide (Ba(OH)2
Nitric acid (HNO3)
Sulfuric acid (H2SO4)

Weak acids are weak electrolytes	Weak bases are weak electrolytes
Phosphoric acid (H3PO4)	Ammonia NH3
Carbonic acid (H2CO3)
All Organic acids (RCOOH) Formic acid HCOOH Acetic acid CH3COOH Benzoic acid C6H5COOH	All amines RNH2 Methyl amine CH3NH2

Bronsted (1879-1947, Denmark) and Lowry (1874-1936, Cambridge, England) expanded the acid/base definitions to make it more useful:

According to that definition we can write for the reaction of and acid (HA) with water (general expression):

Example:

Acids are not limited to neutral molecules as the following two examples show. Any compound that can donate a proton, is a Bronsted acid:

Examples:

For the reaction of any base (B) with water can be expressed in a similar way and we can write (general expression):

Hydroxide ions are generated in this process which make the aqueous solution basic.

Example:

Bases are not limited to neutral molecules. The example below shows how ions interact with water to form a basic solution:

Conjugate Acid/base pairs:

Every reaction between an acid an a base involves the transfer of protons from the acid to the base. We can write the chemical equation:

HA(aq) + B(aq) ó A-(aq) + BH+(aq)

Acid Base C. Base C. Acid C. Base = Conjugated base

C. Acid = Conjugated Acid

In this reaction a proton is transferred from the acid onto the base. The species that lost the proton is then called the "conjugated base" (A^-) and the species that gained the proton is called the "conjugated acid" (BH⁺).

Examples for acid/base reactions and the conjugated acid/base pairs:

Exercise:

Write the conjugated acid fo reach of the following:

H₂PO₄^-

HCO₃^-

CH₃NH₂^-

CH₃NH₃

Write the conjugated base for each of the following:

CH₃COOH

H₂PO₄^-

H₂S

H₃O⁺
 
 

Relative strength of Acids and Bases:

For each acid/conjugate base we can say that

the stronger the acid, the weaker its conjugate base

the weaker the acid, the stronger its conjugate base

and

the stronger the base, the weaker its conjugate acid

the weaker the base, the stronger its conjugate acid

The acid dissociation constants allow us to predict relative strengths of acids and bases (Module I). Table 17.4 on p. 799 of our textbook Chemistry and Chemical Reactivity by KOTZ and TREICHEL summarizes dissociation constants for acids and their conjugated bases.

Now we can apply this concept to the reactions of acids and bases and predict on which side the equilibrium for the reaction lies:

Strong acid in water:

Acid base conj. acid conj. base

From table 17.4 that H₃O⁺ is the strongest acid that can exist in water. Molecules such as HCl do not exist in water. We also find that H₂O is a stronger base than Cl^-, therefore the equilibrium shifts to the right, i.e. the acid is completely dissociated into its ions: HCl is a strong acid.

Weak Acid in Water:

Table 17.4 shows that the hydronium ion (H₃O⁺ ) is a stronger acid than acetic acid (CH₃COOH) and the acetate anion (CH₃COO^-) is a stronger base than water, therefore the equilibrium lies predominately to the left: the acid is mostly undissociated, it is a weak acid.

Exercise:

Predict the position of the equilibrium for each of the following acid/base reactions and determine whether it is a strong or weak acid or base:

The Bronsted acid/base definition is useful describe

• Neutral molecules as acids (SO₂, CO₂)
• Neutral molecules as bases (CH₃NH₂, NH₃)
• Cations as weak acids ([Cu(H₂O)₆]²⁺)
• Anions as weak bases (O^2-, H₂PO₄^-)

Acids that donate one proton = monoprotic acids (HNO₃, HCl)

Acids that can donate more than one proton = polyprotic acids (H₂SO₄, H₃PO₄)

Compounds or ions that can accept or donate a proton = amphoteres

Examples:
 

Acid Form	Amphoteric Form	Base Form
H3PO4	H2PO4-	HPO42-
H2PO4-	HPO2-	PO43-
H3O+	H2O	OH-

Exercises: Predicting the Direction of a Neutralization Reaction:

For each of these experiments, predict the position of the equilibrium.

Hint: First, identify which compound is the acid and which is the base and write the corresponding acid/base pairs to complete the equation. Then decide which side of the equation has the stronger acid and base and determine the position of the equilibrium.

1. You smell ammonia when NH₃ is dissolved in water


NH₃ + H₂O ó




2. The odor of the gas goes away when you add excess HCl:


NH₃ + HCl ó

3. The odor returns when you add an excess of NaOH:
4. NH₄⁺ + NaOH