CBSE Class 7 Science Advanced Chemistry Acids, Bases and Salts Notes Set 01

Acids, Bases and Salts

Introduction

Besides sorting matter into solids, liquids and gases (by physical state) and into elements, compounds and mixtures (by composition), chemists place many compounds into three more groups: acids, bases and salts.

Lemon, orange, grapes and tamarind taste sour. The Latin word acidus means sour, and the name acid was given to the substances that cause this taste. Substances with a bitter taste, such as bitter gourd, and a slippery feel, such as soap, are called bases.

Classification of Acids on the Basis of Origin

Organic acids come from living things (plants and animals), though they can also be made in a laboratory. Mineral acids are made from minerals found in the earth's crust and are also called inorganic acids.

Table 3.1 Some Organic Acids and Their Sources

Table 3.2 Common Mineral Acids

Hydracids and Oxyacids

Looking at what an acid is built from, there are two kinds. A hydracid has only hydrogen and one other nonmetal (for example HCl). An oxyacid has oxygen as well as hydrogen and a nonmetal (for example \(\ce{H2SO4}\)).

Definitions of an Acid and a Base

An acid is a substance that, when dissolved in water, gives \(\ce{H+}\) ions as the only positive ions. A base is a substance that, when dissolved in water, gives \(\ce{OH-}\) ions as the only negative ions. The breaking up of a substance into ions in water is called ionization. Acids and bases differ in how far they ionize, and this is the basis of their classification by strength.

Strong and Weak Acids and Bases

Strong acids and strong bases ionize completely in water and so give a high concentration of \(\ce{H+}\) or \(\ce{OH-}\) ions.

Strong acids:
\(\ce{HCl -> H+ + Cl-}\)
\(\ce{H2SO4 -> 2H+ + SO4^2-}\)
\(\ce{HNO3 -> H+ + NO3-}\)

Strong bases:
\(\ce{NaOH -> Na+ + OH-}\)
\(\ce{KOH -> K+ + OH-}\)

Weak acids and weak bases ionize only partly and give a low concentration of \(\ce{H+}\) or \(\ce{OH-}\) ions. All organic acids are weak acids.

Weak acids:
\(\ce{CH3COOH <=> CH3COO- + H+}\)
\(\ce{H2CO3 <=> 2H+ + CO3^2-}\)
\(\ce{H3PO4 <=> 3H+ + PO4^3-}\)

Weak bases:
\(\ce{Ca(OH)2 <=> Ca^2+ + 2OH-}\)
\(\ce{Al(OH)3 <=> Al^3+ + 3OH-}\)

Naming of Acids

A hydracid gets the name "hydro ... ic acid" with the name of the nonmetal in the middle.

Most nonmetals form two oxyacids. The acid with fewer oxygen atoms is named "-ous acid" and the one with more oxygen atoms is named "-ic acid".

Table 3.3 Naming of Oxyacids

Naming of Bases

Oxides and hydroxides of metals are basic. A base is named by writing the name of the metal (positive radical) followed by "oxide" or "hydroxide". Examples: NaOH is sodium hydroxide and \(\ce{NH4OH}\) is ammonium hydroxide. A base that dissolves in water is called an alkali.

Aqueous Solutions, Concentrated and Dilute

Acids and alkalies dissolve in water, and these solutions are called aqueous solutions. A concentrated solution has relatively little water for the amount of acid or base, whereas a dilute solution has a lot of water.

Basicity of Acids and Acidity of Bases

How many \(\ce{H+}\) or \(\ce{OH-}\) ions a solution contains also depends on how many such ions one molecule can give. The number of \(\ce{H+}\) ions given by one molecule of an acid is its basicity. The number of \(\ce{OH-}\) ions given by one molecule of a base is its acidity. This leads to another way of classifying.

Table 3.4 Acids Classified by Basicity

Table 3.5 Bases Classified by Acidity

Methods of Preparation of Acids and Bases

Acids

1. Direct combination of elements: hydrogen gas and chlorine gas combine to give hydrogen chloride gas, which dissolves in water to form hydrochloric acid.
\(\ce{H2 + Cl2 -> 2HCl}\)

2. From nonmetals: a nonmetal burns in oxygen to form an oxide. If the oxide dissolves in water and gives an acid, it is an acidic oxide.
\(\ce{C + O2 -> CO2}\)
\(\ce{CO2 + H2O -> H2CO3}\)
\(\ce{S + O2 -> SO2}\)
\(\ce{SO2 + H2O -> H2SO3}\)
\(\ce{2SO2 + O2 -> 2SO3}\)
\(\ce{SO3 + H2O -> H2SO4}\)

Oxides that do not dissolve in water and give no acid are called neutral oxides. Examples: NO and CO.

Bases

1. From metals: a metal heated in oxygen gives a metal oxide, which is basic. The oxide that dissolves in water gives the corresponding base.
\(\ce{4Na + O2 -> 2Na2O}\)
\(\ce{Na2O + H2O -> 2NaOH}\)
\(\ce{2Mg + O2 -> 2MgO}\)
\(\ce{MgO + H2O -> Mg(OH)2}\)

2. Slaked lime from quicklime: quicklime (calcium oxide) combines with water to give slaked lime (calcium hydroxide) and a lot of heat.
\(\ce{CaO + H2O -> Ca(OH)2 + heat}\)
Slaked lime dissolves only a little and mostly settles at the bottom. The clear water above it is called lime water.

Physical Properties of Acids and Bases

Table 3.6

Chemical Properties of Acids

Table 3.7

Chemical Properties of Bases

Table 3.8

Comparative Study of Acids and Bases

Table 3.9

Example

Why is KOH called caustic potash?
KOH burns (causticises) the skin, and so it is called caustic potash.

Example

Give an example of dibasic acid. How is it obtained from the corresponding nonmetal? Give balanced equations.
Sulphuric acid is a dibasic acid. Sulphur burns in oxygen to give sulphur dioxide; this is oxidised further to sulphur trioxide, and sulphur trioxide dissolves in water to give sulphuric acid.
\(\ce{S + O2 -> SO2}\)
\(\ce{2SO2 + O2 -> 2SO3}\)
\(\ce{SO3 + H2O -> H2SO4}\)

Neutralization

An acid and a base always react to give a salt and water. This is called neutralization. In it, one \(\ce{H+}\) ion of the acid is cancelled by one \(\ce{OH-}\) ion of the base. When all the \(\ce{H+}\) ions of the acid solution have been cancelled by an equal number of \(\ce{OH-}\) ions, the neutralization is complete. The quantities of acid and base needed for this depend on the total number of \(\ce{H+}\) and \(\ce{OH-}\) ions that each can produce.

Indicators

The point at which the \(\ce{H+}\) ions of an acid have been exactly cancelled by the \(\ce{OH-}\) ions of a base is the neutralization point. Adding more base (or more acid) beyond this point puts an excess of \(\ce{OH-}\) (or \(\ce{H+}\)) and turns the solution basic (or acidic). To spot the neutralization point, substances that have different colours in acidic and basic solutions are added. Because they show the change at that point, they are called indicators. Some natural materials, such as red cabbage juice, work as indicators. Others in common use are litmus, methyl orange and phenolphthalein.

Table 3.10 Action of Indicators

Salts

A salt is a compound made of positive and negative ions (radicals) that ionizes completely in its aqueous solution. Salts are formed by neutralization:
Acid + Base \(\ce{->}\) Salt + Water

If the reaction is stopped midway so that some replaceable hydrogen ions of the acid, or some replaceable hydroxyl ions of the base, are left, it is called partial neutralization. It produces acidic or basic salts.

Types of Salts

1. Normal salts: every \(\ce{H+}\) of the acid has been replaced by metal ions. They form by complete neutralization.
\(\ce{2NaOH + H2SO4 -> Na2SO4 + 2H2O}\)
\(\ce{2H3PO4 + 3Ca(OH)2 -> Ca3(PO4)2 + 6H2O}\)

2. Acidic salts: a dibasic or tribasic acid is only partly neutralized, so the salt still has at least one replaceable hydrogen. Such a salt gives \(\ce{H+}\) ions in water and can be neutralized further until a normal salt forms.
\(\ce{NaOH + H3PO4 -> NaH2PO4 + H2O}\) (sodium dihydrogen phosphate, acidic salt)
\(\ce{NaOH + NaH2PO4 -> Na2HPO4 + H2O}\) (disodium hydrogen phosphate, acidic salt)
\(\ce{NaOH + Na2HPO4 -> Na3PO4 + H2O}\) (sodium phosphate, normal salt)

3. Basic salts: a diacidic or triacidic base is only partly neutralized, so the salt still has at least one \(\ce{OH-}\). It gives \(\ce{OH-}\) ions in water and can be neutralized further to a normal salt.
\(\ce{Ca(OH)2 + HCl -> Ca(OH)Cl + H2O}\) (calcium hydroxychloride, basic salt)
\(\ce{Ca(OH)Cl + HCl -> CaCl2 + H2O}\) (calcium chloride, normal salt)

Table 3.11 Acidic Salts and Basic Salts Compared

4. Double salts: a salt that is a combination of two simple salts. It dissociates completely in water and gives the reactions of all the ions of both salts. Examples: Mohr's salt, \(\ce{FeSO4.(NH4)2SO4.6H2O}\), which gives \(\ce{Fe^2+}\), \(\ce{NH4+}\) and \(\ce{SO4^2-}\) ions and their tests; and potash alum, \(\ce{K2SO4.Al2(SO4)3.24H2O}\).

5. Mixed salts: salts that contain more than one basic radical or more than one acidic radical. They dissociate completely in water and give the reactions of all their ions. Examples: sodium potassium carbonate, \(\ce{NaKCO3}\) (radicals \(\ce{Na+}\), \(\ce{K+}\), \(\ce{CO3^2-}\)); bleaching powder, \(\ce{CaOCl2}\) (radicals \(\ce{Ca^2+}\), \(\ce{OCl-}\), \(\ce{Cl-}\)).

6. Complex salts: salts that contain one or more simple ions along with a complex ion. A complex ion is a group of atoms that stays together as one unit and does not split into separate atoms or ions. The salt dissociates into the complex ion and the simple ion. Examples: potassium ferrocyanide, \(\ce{K4[Fe(CN)6]}\), which dissociates as \(\ce{K4[Fe(CN)6] -> 4K+ + [Fe(CN)6]^4-}\) (simple ion and complex ion); and sodium argentocyanide, \(\ce{Na[Ag(CN)2]}\).

Derivation of Names of Salts

The name of a salt comes from the acid that produced it. Salts of hydracids end in "-ide". If hydrogen is still present in the salt, the word "hydrogen" is added.

Table 3.12 Naming of Salts of Hydracids

Salts of oxyacids depend on the kind of oxyacid. Salts of "-ous" acids end in "-ite" and salts of "-ic" acids end in "-ate". If hydrogen is present, "hydrogen" or "bi" is added to the name.

Table 3.13 Naming of Salts of Oxyacids

In the same way, salts of nitrous acid are nitrites and salts of nitric acid are nitrates. Salts of carbonic acid are carbonates and bicarbonates (hydrogen carbonates). In a salt, the positive part is the basic radical and the negative part is the acidic radical.

Preparation of Simple Salts

1. Sodium chloride from sodium hydroxide and hydrochloric acid: \(\ce{NaOH + HCl -> NaCl + H2O}\)
2. Potassium chloride from potassium hydroxide: \(\ce{KOH + HCl -> KCl + H2O}\)
3. Calcium chloride from calcium hydroxide: \(\ce{Ca(OH)2 + 2HCl -> CaCl2 + 2H2O}\)
4. Zinc sulphate from zinc oxide: \(\ce{ZnO + H2SO4 -> ZnSO4 + H2O}\)
5. Copper sulphate from copper oxide: \(\ce{CuO + H2SO4 -> CuSO4 + H2O}\)
6. Ammonium chloride from ammonium hydroxide: \(\ce{NH4OH + HCl -> NH4Cl + H2O}\)

Properties of Salts

1. Salts are solid at room temperature.
2. Salts of sodium and potassium are colourless, while those of iron, copper and chromium are coloured. Examples: ferrous sulphate is green, aqueous copper sulphate is blue and potassium dichromate is orange-red.
3. Salts have high melting points.
4. Most salts dissolve in water. The exceptions are the carbonates, sulphides and sulphates of some metals.
5. Aqueous solutions of salts conduct electricity well, so salts are called electrolytes.
6. When a salt is dissolved in a fixed amount of water, three kinds of solution are possible:
(a) Unsaturated solution: it can still dissolve more solute.
(b) Saturated solution: it cannot dissolve any more solute.
(c) Supersaturated solution: under special conditions, a saturated solution is made to hold even more salt.

Hydrated Salts

The crystals of some salts hold water molecules. These are hydrated salts. Each molecule of the salt carries a fixed number of water molecules, called the water of crystallization. This is written after the formula of the salt.
Examples: \(\ce{CuSO4.5H2O}\), \(\ce{FeSO4.7H2O}\), \(\ce{Na2SO4.10H2O}\).

Salts without water of crystallization are anhydrous salts, for example \(\ce{CuSO4}\) and \(\ce{FeSO4}\).

Every hydrated salt loses its water of crystallization on heating and becomes anhydrous. In some salts the colour is due to this water, so they become colourless on heating. Blue \(\ce{CuSO4.5H2O}\) loses its blue colour and turns into white anhydrous \(\ce{CuSO4}\).

Efflorescent substances: some salts lose their water of crystallization, fully or partly, just by standing in the air. For example, \(\ce{Na2CO3.10H2O}\) gives up nine water molecules in air and the last one on heating.

Deliquescent substances: these absorb moisture from the air and dissolve in it, so they turn into a solution. Examples: calcium chloride, magnesium chloride, mercuric nitrate, zinc chloride.

Hygroscopic substances: these absorb moisture from the air but show no change in physical state. Examples: calcium oxide, sulphuric acid, cobalt chloride.

Uses of Acids, Bases and Salts

Uses of Acids

Table 3.14

Uses of Bases

Table 3.15

Uses of Salts

Table 3.16

Example

Why is nitric acid used in the purification of gold?
Nitric acid is a good oxidising agent. When impure gold is treated with it, the impurities are oxidised and removed, while gold remains.

Example

What is meant by neutralisation? Give balanced equations showing neutralisation reactions of the following. (a) Slaked lime and sulphuric acid (b) Sodium bicarbonate and hydrochloric acid. (c) Magnesium carbonate and sulphuric acid.
An acid reacts with a base (alkali) to give a salt and water. This reaction is called neutralization.
(a) \(\ce{Ca(OH)2 + H2SO4 -> CaSO4 + 2H2O}\)
(b) \(\ce{NaHCO3 + HCl -> NaCl + CO2 + H2O}\)
(c) \(\ce{MgCO3 + H2SO4 -> MgSO4 + CO2 + H2O}\)

Points to Remember

1. How strong an acid or base is depends on how much of it ionizes.
2. Metal oxides dissolved in water give bases, and nonmetal oxides dissolved in water give acids.
3. A metal that is more reactive than hydrogen releases hydrogen from dilute hydrochloric or dilute sulphuric acid.
4. Neutralization is the reaction of an acid with a base to form a salt and water.
5. Indicators help to find the neutralization point.
6. Calcium hydroxide is used for whitewashing and for making mortar.
7. Acidic salts have replaceable \(\ce{H+}\), basic salts have replaceable \(\ce{OH-}\), and normal salts have neither.
8. Aqueous solutions of salts conduct electricity.
9. Hydrated salts hold water of crystallization; anhydrous salts do not.
10. Efflorescent substances lose water of crystallization in air, deliquescent substances absorb moisture and dissolve, and hygroscopic substances absorb moisture without any change in state.
11. Salt suffixes from the acid: sulphuric acid gives sulphate, sulphurous acid gives sulphite, hydrochloric acid gives chloride, nitric acid gives nitrate, nitrous acid gives nitrite and phosphoric acid gives phosphate.
12. Many salts are known by common names: baking soda is sodium bicarbonate and nitre is potassium nitrate.