Chemistry

a. To each chemical test whether it is an acid, a base or neutral by using blue lit-mus paper, red litmus paper and universal indicator (better).

Record your obser-vations and colour intensities in Table 16.1.

b. Comparing the colour intensities of the indicators, explain the source their difference.

Study Questions

1. Which of the chemicals are acids?

2. Which of the chemicals are bases?

3. Which of the chemicals are neutral?

16.1.1. Arrhenius’ Theory of Acid-Base

The first person to recognize the essential nature of acids and bases was the Swedish scientist Svante Arrhenius (1859–1927). On the basis of his experiments with electrolytes, Arrhenius postulated that acids produce hydrogen ions in aqueous solution, while bases produce hydroxide ions.

i. An acid is a substance which dissociates in aqueous solution to produce hydrogen ions (H⁺) as the only positive ions

ii. A base is any substance that dissociates in aqueous solution to produce hydroxyl ions (OH^-) as the only negative ions.Notice that the Arrhenius concept of acid and base requires the presence of water.In this theory, the products of an acid-base reaction are: salt and water:

This reaction acid-base shows that when an acid reacts with a base, the result is a neutral solution of a salt in water; this explains why the reaction is called an acid-base neutralization reaction: acidic properties and basic properties neutralize each other.

Notice

The observations in activity 16.1 show that all acids and bases do not show the same intensity of colours ; this means that although all are acids or bases they give different concentrations of H⁺ or OH^- ions.

This means that not all acids and bases dissociate into their ions at the same extent. Some dissociate completely in water, other dissociate partially.

Acids and bases that dissociate completely into their ions when dissolved in water are qualified as “strong”.

weak base.

Since strong acids and strong bases dissociate completely in water whereas weak acids and weak bases dissociate partially, the concentration [H⁺] in a solution of a strong acid will be higher than the concentration [H⁺] in the solution of a weak acid of the same concentration; in other words pH of the strong acid solution will be lower that pH of the weak acid solution of the same concentration.

For the same concentration: pH of HCl (aq) < pH of H₃PO₄(aq)In the same way, the concentration OH^- in a strong base solution will be higher than the concentration OH^- of a weak base solution of the same concentration; in this case the concentration H⁺ in a strong base solution will be lower than the concentration H⁺ in a weak base solution of the same concentration, or pH of strong base solution will be higher than pH of the weak base solution of the same concentration.

For the same concentration: pH of NaOH(aq) > pH of NH₄OH(aq).

a. Two strongest acidic substances

b. Two weakest acidic substances

c. Two most alkaline substances

d. Two least alkaline substances

e. Neutral substance (s).

3.i. Take two different solutions with equal concentrations of acids such as HCl and CH₃COOH, and measure their pH.

ii. Take two different solutions with equal concentrations of bases such as NaOH and NH₄OH, and measure their pH.

Questions:

a. What are the observations about the results?

b. How do you interpret those results?

16.1.2. Brønsted-Lowry’s Acid-Base Theory: Acid-base conjugate pairs

The weakness of the Arrhenius acid-base theory is that it requires the presence of water and the acid and base properties are due only to the presence of H⁺ and OH^-ions respectively in an aqueous solution. Hence, for example, hydrogen chloride, HCl(g), is not an acid and NH3 is not a base according to this theory.

In 1923, the Danish Chemist Johannes Nicolaus Brønsted (1879–1947) and English Chemist Thomas Martin Lowry(1874 – 1936) independently developed the definitions of acids and bases based on the compounds’ abilities to either donate or accept protons (H⁺ ions). In this theory, acids are defined as proton donors whereas bases are defined as proton acceptors. A compound that acts as both a Brønsted-Lowry acid and base is called amphoteric.

Briefly in this theory, an acid is a proton (H⁺) donor, and a base is a proton acceptor.

In this theory you notice that OH^- ion is not the only base; any chemical species that can accept a proton is a base, be it in aqeous solution or not. H⁺ continues to be the only source of acid properties except that the presence of water is no more a prerequisite.

Conjugate acids and bases

Let us consider an acid HA which ionizes partially in aqueous solution according to a reversible reaction below:

Forward reaction:

HA is an acid because it is donating a proton to H₂O.The water (H₂O) is a base because it is accepting a proton from the acid (HA).

•Reverse reaction

The H₃O⁺ is an acid because it is donating a proton to the A^- ion.

The A^- ion is a base because it is accepting a proton from the H₃O⁺.

The reversible reaction contains two acids and two bases. We take them as pairs, called conjugate pairs.

According to the forward reaction, HA is the acid and its conjugate base is A^-. According to the reverse reaction, H₃O⁺ is an acid whose conjugate base is H₂O.The above reaction can be represented as:

The acid-base conjugate pairs are written as Acid1/ Base 1 and Acid 2/ Base2.

The conjugate acid-base pairs for this reaction are: NH₄⁺/ NH3 and H₂O/ OH^-.

In general, when an acid is strong, its conjugate base is weak and vice-versa.

Remark

The observation of the above examples shows that water can act as both Brønsted-Lowry base or a Brønsted-Lowry acid. Thus, water is known as an amphoteric solvent.

Checking up 16.1 (b)

1. Label the acid (A), base (B), conjugate acid (CA), and conjugate base (CB) in each of the following reactions.

2. Give the conjugate base for each of the following Brønsted-Lowry acid.

a. HI

b. NH₄+c. H₂CO₃

d. HNO₃

3. Give the conjugate acid for each of the following Brønsted-Lowry bases.

a. CN^–

b. O₂^–

c. CH₃COO^–

d. NH₃

16.1.3. Lewis’s Acid-Base Theory

In 1938, an American chemist Gilbert Newton Lewis (1875–1946) introduced another theory which extends the concept of acids and bases further than those of Arrhenius and Brønsted-Lowry.

An acid is a chemical species that has a vacant orbital and can accept a pair of electrons from another chemical species. An acid is an electron pair acceptor or electron deficient.

In this theory, hydrogen ion, H⁺, is electron deficient and can accept a pair of electrons as any electron deficient species; it is no more the only responsible of acidic properties.Other various species can act as Lewis acids. They include the following chemical species.

a. All cations, but particularly cations of transition metals; e.g. Cu²⁺, Fe²⁺, Fe³⁺.

b. Chemical molecules whose central atom does not fulfill the octet rule; e.g BF₃, AlC_l3, BH₃,...

A base is a substance that possesses one or more lone pairs of electrons which may be used for the formation of a coordinate bond. A base is an electron pair donor.

A Lewis base can be an anion, such as OH^-, Cl^- , CN^- or a neutral molecule with lone pair of electrons such as H₂Ö:, :NH₃, etc...

Checking up 16.1(c)

1. Write equations of reactions between the following species and state which species are acids and which one are bases.

a. Fe³⁺ and H₂O to form hexaaquairon (III).

b. Cu²⁺ and NH₃ to form tetraamminecopper (II).

c. Cr³⁺ and H₂Oto form hexaaquachromium (III).

2. What is the name of the bond formed between a Lewis acid and a Lewis base?

(a) Covalent bond

(b) Double bond

(c) Coordinate covalent bond

(d) Ionic bond

3. Why is BH₃ (boron trihydride) a Lewis Acid? Identify the correct answer.

a. It is a cation.

b. It is electron deficient.

c. It can accept more than 8 valence electrons.

d. It contains double bonds

16.2. End unit Assessment

1. Write an equation for the dissociation of the following acids in water:

a) HClO₄

b) H₂SO₄

c) CH₃COOH

d) H₂S

e) HNO₃

2. Write the balanced reaction for what happens when nitric acid is put in water.

3. Write the balanced reaction for what happens when acetic acid is put in water.

4. Write an equation for the dissociation of Ba(OH)₂ in water.

5. What is the conjugate base of HSO^4-?

6. Write the formula and name for the conjugate acid for the following bases:

a) NH₃

b) PO^43-

c) CN^-

d) HCO^3-

7. Write a balanced equation for the Bronsted-Lowry acid HPO^42- in water.

8. Write and balance the reaction for ammonia in water.

9. The acidity strength of group VII oxoacids increases with the number of oxygen atoms. With reference to Bronsted-Lowry theory of acids and bases, rank their conjugate bases hereafter in increasing order of alkalinity strength.

a) ClO^-

b) ClO^2-

c) ClO^3-

d) ClO^4-

10. Rank the following chemical species in decreasing order of acidity strength.