Acids, Bases and Salts

Introduction

The word 'acid' is derived from the Latin word Acidus, meaning sour-tasting. The theories of acids and bases have been central to chemistry for centuries, shaping our comprehension of chemical reactions and their effects on the world around us. The historical journey of acid-base theories can be traced back to ancient civilizations. However, the formalization of acid-base theories began in the late 19^th Century with Svante Arrhenius.

 

THEORIES

Antoine Lavoisier

Antoine Lavoisier introduced one of the earliest modern acid-base theories in the late 18^th century. He proposed acids as substances that contained oxygen and believed that the presence of oxygen was responsible for their acidic properties.

Humphry Davy (1810)

Additional research in 1810 revealed that muriatic acid was a compound of hydrogen and chlorine (HCl), and contained no oxygen, and Davy’s research on chlorine led to a rejection of Lavoisier’s theory that oxygen was an essential constituent of acids.

Arrhenius Theory (1884)

The first comprehensive and widely accepted acid-base theory was proposed by Svante Arrhenius in 1884. Arrhenius defined acids as any substances that release or produce hydrogen ions (H⁺) in water, and bases as substances that release hydroxide ions (OH^-),  e.g.

This groundbreaking theory provided a clear definition of acids and bases in aqueous solutions, and the process demonstrated above is known as ionization. The characteristic properties of an acid in solution are due to the presence of these hydrogen ions.

 

Modification of Arrhenius' definition of an acid

The existence of hydronium (oxonium) ion, H₃O⁺, modifies the earlier definition of an acid. Thus, an acid is also defined as a substance which in aqueous solution produces hydronium (oxonium) ions, H₃O⁺ as the only positive ions, e.g.

HCl + H₂O → H₃O⁺ + Cl^-

HNO₃ + H₂O → H₃O⁺ + NO3-

H₂SO₄ + H₂O → H₃O⁺ + SO₄^2-

 

Bronsted- Lowry Theory (1923)

The Arrhenius Theory makes use of hydroxide ions, which may not exist in non-aqueous solvents, and does not cover weak bases, this prompted Bronsted and Lowry to put forward a more general theory of acids and bases which incorporates all protonic solvents, and not just water.

They defined an acids as substances that can donate protons (H⁺ ions), and bases are substances that can accept protons (H⁺). E.g.

HCl + H₂O ↔ H₃O⁺ + Cl^-

HCl + NH₃ ↔ NH₄⁺ + Cl^-

H₂O + NH₃ ↔ H₃O⁺ + NH₄⁺

In the forward reactions above a proton is being donated by (HCl in example 1, HCl in example 2 and H₂O in example 3) and accepted by (H₂O in example 1, NH₃ in example 2 and 3). Any substance which donates a proton to the other have the potential of being an acid while the other which accept the proton is the base.

If you study the forward reaction of equation 1, HCl donates a proton to H₂O which accepts it. HCl behaves as an acid while H₂O behaves as a base. In the backward reaction H₃O⁺ donates a proton behaving like an acid and Cl^- accepts a proton behaving like a base. From the above explanation, HCl and Cl- are a conjugate acid-base pair. Similarly, H₃O⁺ and H₂O are a conjugate acid-base pair.

Lewis Theory

In 1923, Gilbert N. Lewis proposed a generalized theory of acids and bases, expanding on the existing Arrhenius and Brønsted-Lowry definitions. Through the use of the Lewis definition of acids and bases, chemists are now able to predict a wider variety of acid-base reactions. Lewis' theory used electrons instead of proton transfer and specifically stated that:

An acid is a species that accepts an electron pair, while a base donates an electron pair.

The reaction of a Lewis acid and a Lewis base will produce a coordinate covalent bond as seen above. A coordinate covalent bond is just a type of covalent bond in which one reactant gives it electron pair to another reactant. In this case the lewis base donates its electrons to the Lewis acid. When they do react this way the resulting product is called an addition compound, or more commonly an adduct.

• Lewis Acid: a species that accepts an electron pair (i.e., an electrophile) and will have vacant orbitals
• Lewis Base: a species that donates an electron pair (i.e., a nucleophile) and will have lone-pair electrons

 

Several categories of substances can be considered Lewis acids:

1) All cations, i.e. positive ions (e.g., Cu²⁺, Fe²⁺, Fe³⁺)

2) An atom, ion, or molecule with an incomplete octet of electrons (e.g., BF₃, AlF₃).

3) Molecules that have multiple bonds between two atoms of different electronegativities (e.g., CO₂, SO₂)

4) Molecules where the central atom can have more than 8 valence shell electrons can be electron acceptors (e.g., SiBr₄, SiF₄).

 

Several categories of substances can be considered Lewis bases:

1) All anions, i.e. negative ions (e.g. OH⁻, CN⁻, CH3COO⁻⁾

2) Atom of one of more unshared pairs in the valence shell (e.g. NH₃, H₂O)

3) The presence of a double bond e.g. CO.

 

Amphoterism

As of now you should know that acids and bases are distinguished as two separate things however some substances can be both an acid and a base. You may have noticed this with water, which can act as both an acid or a base. This ability of water to do this makes it an amphoteric molecule. Water can act as an acid by donating its proton to the base and thus becoming its conjugate acid, OH-. However, water can also act as a base by accepting a proton from an acid to become its conjugate base, H₃O⁺.

• Water acting as an Acid:

H₂O + NH₃ → NH₄⁺ + OH⁻

• Water acting as a Base:

H₂O + HCl → Cl⁻ + H₃O⁺

You may have noticed that the degree to which a molecule acts depends on the medium in which the molecule has been placed in. Water does not act as an acid in an acid medium and does not act as a base in a basic medium. Thus, the medium which a molecule is placed in has an effect on the properties of that molecule. Other molecules can also act as either an acid or a base. For example,

Al(OH)₃ + 3H⁺ → Al³⁺ + 3H₂O

• where Al(OH)₃ is acting as a Lewis Base.

Al(OH)₃ + OH⁻ → Al(OH)₄^-

• where Al(OH)₃ is acting as an Lewis Acid.

Note how the amphoteric properties of the Al(OH)₃ depends on what type of environment that molecule has been placed in.

 

Classification of Acids

Acids can be categorized into different types based on various criteria, including their composition, strength, and occurrence. Here are some common types of acids:

A. Based on their Origin (Source)

1. Organic Acids: Organic acids are a type of weak acid that contains a carbon atom bonded to hydrogen and a carboxyl group (COOH). They are found in various natural sources, including fruits and other organic matter. Examples include acetic acid (CH₃COOH) found in vinegar and citric acid (C₆H₈O₇) found in citrus fruits.

2. Mineral Acids: Mineral acids are strong acids that are typically derived from minerals and inorganic compounds. They are highly corrosive and can cause severe chemical burns.

Examples include sulfuric acid (H₂SO₄), hydrochloric acid (HCl), and nitric acid (HNO₃)

 

B. Based on their Strength

1. Weak Acids: Weak acids are acids that only partially dissociate in water, releasing a limited number of hydrogen ions (H⁺). Their dissociation equilibrium lies more to the left.

Examples include acetic acid (CH₃COOH) and carbonic acid (H₂CO₃).

2. Strong Acids: Strong acids completely dissociate in water, releasing all their hydrogen ions (H⁺). Their dissociation equilibrium lies almost entirely to the right. Examples include hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃).

 

C. Based on elements present in the Acids 

1. Hydracids (Binary acids): are acids that consist of hydrogen and a non-metal element other than oxygen. 

Examples:

• Hydrochloric acid (HCl)
• Hydrogen sulphide (H₂S)
• Hydrobromic acid (HBr)
• Hydroiodic acid (HI) 

Oxyacids (Ternary acids):  These acids contain hydrogen, non-metal elements, and hydrogen. 

Examples:

• Sulphurous acid (H₂SO₃)
• Chloric acid (HClO₃)
• Perchloric acid (HClO₄)
• Sulphuric acid (H₂SO₄) 

 D. Based on Basicity of Acids 

Basicity is defined as the number of hydronium ions [H⁺ (aq)] released by one acid molecule to complete ionization. 

Based on basicity, acids are divided into three types. 

1. Monobasic acid: When one acid molecule on complete ionisation gives one hydronium ion [H⁺ (aq)], the acid is called monobasic acid. 

Examples: 

• HCl
• Hydrogen fluoride acid (HF)
• HBr
• Acetic acid (CH₃COOH)
• Formic acid (HCOOH)

Note:

        I.            A monobasic acid always ionizes in one step in an aqueous solution. 

      II.            A monobasic acid forms only a single salt. 

  

2. Dibasic acid: One molecule of dibasic acid on complete ionisation gives two hydronium ions [H⁺ (aq)]. 

Examples:

• These acids ionise in two steps. 

 

3. Tribasic acid: Here, one molecule of an acid gives three hydronium ions [H⁺ (aq)] on complete ionization. 

Complete ionisation is taking place in three steps for tribasic acids 

 

E. Based on Concentration (water content) of Acid

Concentration: It measures the amount of water present in a given sample of acid. 

Based on this, the acids are classified into two types as follows: 

1. Concentrated acid: A sample of an acid that contains only very little or no amount of water. 

2. Dilute acid:  For dilute acids, they contain a lot more amount of water than their weight 

 

  

General Properties of Acids

1. Taste

• Generally, acids taste sour. 

 

2. Effect on Skin

• Strong inorganic acids are corrosive action on the skin, and they can even damage other substances. 

 

 

3. Effect of Indicators

• Indicator: Indicators are the substances used to distinguish between acids and bases. They change their color when dissolved in the acids or bases; that is why they are known as indicators. 
• Effect Of Litmus Paper: Litmus is a natural indicator, and when the blue litmus paper is dipped in the acid solution, they turn blue litmus paper into the red. 

4. Effect of Electric Current

• Acids are good conductors of electricity when dissolved in water. This is because, in water, they get dissolved to release Hydrogen (H⁺ as Hydronium) ions. These free ions help in the conduction of electricity.  

 

Chemical Properties of Acid

 

Uses of Acids

1. Oxalic Acid

• Oxalic acid is used for removing food and ink stains. 
• It is also used in wastewater treatment to remove deposits of calcium. 

2. Sulphuric Acid

• It is very important industrial chemical and used in the preparation of many chemicals.  
• In the car batteries, sulphuric acid is used. 
• It is also used in the preparation of paints, dyes, and fertilizers. 

 

3. Acetic Acid

• The common name of acetic acid is vinegar and it used in the food industry to enhance the flavours of food. 
• Apart from that, it is also used to clean the utensil, floors. 
• It also used to remove stain of wooden furniture. 
• It is used as preservatives in many food items like ketchups, pickels, sauces. 

4. Boric Acid

• Boric acid is used as grain preservatives and used in the manufacture of glasses and adhesives. 

5. Hydrochloric Acid

• It is responsible for the digestion of the food we eat. 
• It is also used as a bathroom cleaner. 

6. Carbonic Acid

• In the preparation of soft drinks 

7. Tartaric Acid

• In the baking powder. 

8. Citric Acid

• In food prevention 

9. Ascorbic Acid

• In the treatment of scurvy disease and bone marrow. 

  

Methods of Preparation of Acids

Acids can be prepared in many ways; we will study the three most important methods for the preparation of acids:

1. Synthetic Method  

In the synthetic method, acids are prepared by a direct combination of elements, mostly non-metals. 

For example, when hydrogen gas and chlorine react with each other under the action of electricity, the hydrogen chloride gas is absorbed in water to produce hydrochloric acid. 

The chemical reaction can be written as below: 

Another is the formation of hydrogen sulphide; when the non-metals like sulphur and hydrogen gas are boiled, it forms hydrogen sulphide gas. 

  

2. By Dissolving Acidic Oxides in Water 

Acidic oxides: some of the oxides of non-metals are dissolved in the water. These oxides are known as acidic oxides. 

When acidic oxides of carbon, i.e., carbon dioxide, react with water, it gives carbonic acid. 

Similarly, sulphur trioxide dissolves in the water forms sulphuric acid. 

Detailed examples include:

 

3. The Reaction of an Acid with Salt of Another Acid

The principle of this preparation method is that the acid which has a higher boiling point reacts with the salt of an acid having the lower boiling point to produce an acid.  

For example, common salt, i.e., sodium chloride is salt and salt of hydrochloric acid (HCl).   

The boiling point of hydrochloric acid is lesser than sulphuric acid (H₂SO₄), and therefore, when sodium chloride (NaCl) reacts with sulphuric acid, it gives the product as hydrochloric acid (HCl). 

 

Other examples include: