CHEMISTRY ACIDS, BASES & SALTS

Acids, Bases & Salts

Introduction

Elements combine to form numerous compounds. On the basis of their chemical properties, compounds
can be classified into three categories:
 Acids
 Bases
 Salts

Acids and Bases in the Laboratory

Indicators
An indicator tells us whether a substance is acidic or basic in nature, by the change in colour.

Common Indicators
 An acid turns blue litmus red and a base turns red litmus blue.
 Methyl orange indicator gives a red colour in an acidic solution and gives a yellow colour in a basic
solution.
 Phenolphthalein is colourless in an acidic solution and gives a pink colour in a basic solution.

Olfactory Indicators
 Those substances whose odour changes in acidic or basic media are called olfactory indicators. For
example: onion, vanilla and clove oil.
 On adding sodium hydroxide solution to a cloth strip treated with onion, the smell of the onion is not
detected. An acidic solution does not eliminate the smell of the onion.

Reaction of Acids & Bases with Metals

Acids react with metals to produce salt by displacing hydrogen.

For Example:
i. When dilute sulphuric acid reacts with the metal zinc, zinc sulphate is formed with the evolution of
hydrogen gas.
Zn + H2SO4 → ZnSO4 + H2

ii. Zinc is the only metal which reacts with sodium hydroxide to form sodium zincate with the release of
hydrogen gas.
Zn + 2NaOH → Na2ZnO2 + H2

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Reaction of Metal Carbonates & Bicarbonates with Acids

Acids react with metal carbonates or bicarbonates to form salt and water with the evolution of carbon
dioxide gas.
For Example:
i. Hydrochloric acid reacts with sodium carbonate to form sodium chloride and water with the release of
carbon dioxide gas.
Na2CO3(s) + 2 HCl(aq) → 2NaCl(aq) + CO2(g) + H2O(l)

ii. Similarly, sodium bicarbonate also reacts with hydrochloric acid to form sodium chloride and water
with the release of carbon dioxide gas.
NaHCO3(s) + HCl (aq) → NaCl (aq) + CO2(g) + H2O(l)

Neutralisation

The reaction between an acid and a base to form salt and water is called a neutralisation reaction.
For example:
Hydrochloric acid reacts with sodium hydroxide to form sodium chloride and water.
HCl + NaOH → NaCl + H 2O

Reaction of Metallic Oxides with Acids

Acids react with metallic oxides to form salt and water.

For Example:
Copper oxide (II), a black metal oxide reacts with dilute hydrochloric acid to form a blue-green coloured
copper chloride (II) solution.
CuO + 2HCl → CuCl2(aq) + H2O

Reaction of Non-Metallic Oxides with Base

Bases react with non-metallic oxides to form salt and water.

For Example:
Calcium hydroxide reacts with non-metallic oxides like carbon dioxide to form calcium carbonate salt and
water.
Ca(OH)2 + CO2 → CaCO3 + H2O

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Acids and Bases in Water

Acids
An acid is a substance which dissociates (or ionises) when dissolved in water to release hydrogen ions.
For Example:
An aqueous solution of hydrochloric acid dissociates to form hydrogen ions. Since hydrogen ions do not
exist as H+ in solution, they combine with polar water molecules to form hydronium ions [H 3O+].
HCl (aq) → H+(aq) + Cl-(aq)
H+ + H2O → H3O+
The presence of hydrogen ions [H+] in hydrochloric acid solution makes it behave like an acid.

Bases
A base is a substance which dissolves in water to produce hydroxide ions [OH- ions].
Bases which are soluble in water are called alkalis.
For Example:
Sodium hydroxide dissolves in water to produce hydroxide and sodium ions.
NaOH (aq) → Na+(aq) + OH-(aq)
The presence of hydroxide ions [OH-] in sodium hydroxide solution makes it behave like a base.

pH Scale

 pH of a solution: pH of a solution is the negative logarithm to the base 10 of the hydrogen ion
concentration expressed in mole per litre.
pH = –log (H+)
10
p H = 7 - Neutral [H+] = [OH-]
pH less than 7 - [H+] more than [OH-]
Acidic
pH more than 7 - [OH-] more than [H+]
Basic

 pH scale : It is a scale showing the relative strength of acids and alkalis.

 The normal pH scale ranges from 0 to 14 as given below.

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Universal Indicator
In case of a colourless liquid, the accurate pH can be obtained by adding a universal indicator.
It is a mixture of several indicators and shows different colours at different concentration of hydrogen ions
in a solution.
For Example:
i. A universal indicator produces green colour in a neutral solution, pH = 7.
ii. The colour changes from blue to violet as pH increases from 7 to 14.
iii. The colour changes from yellow to pink and then to red as pH decreases from 7 to 1.

Importance of pH in everyday life

pH change and survival of animals
 Our body works well within a narrow pH range of 7.0 to 7.8.
 When the pH of rain water is less than 5.6, it is known as acid rain.
 When this acid rain flows into rivers, it lowers the pH of the river water making the survival of
aquatic life difficult.

pH in our digestive system

 Our stomach produces hydrochloric acid which helps in the digestion of food without harming the
stomach.
 Sometimes excess acid is produced in the stomach which causes indigestion.
 To get rid of this pain, bases called antacids are used.
 Antacids are a group of mild bases which react with the excess acid and neutralise it.
 Commonly used antacids are magnesium hydroxide [Mg(OH)2] & sodium bicarbonate[NaHCO3]

pH change - Cause of tooth decay

 Tooth decay starts when the pH in the mouth falls below 5.5.
 Tooth enamel is made up of calcium phosphate which is the hardest substance in the body.
 It is insoluble in water but gets corroded when the pH in the mouth falls below 5.5.
 The bacteria present in the mouth produce acids due to the degradation of sugar and food particles
after eating.
 Hence, to prevent tooth decay, the mouth should be rinsed after eating food and toothpastes which
are basic should be used cleaning teeth to neutralise the excess acid.

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More about Salts

Salts having same positive ions (or same negative ions) are said to belong to a family of salts.

pH of Salts
 Salts of strong acid and a strong base are neutral, with a pH value of 7.
For Example: NaCl, Na2SO4
 Salts of strong acid and weak base are acidic, with a pH value less than 7.
For Example: Ammonium chloride solution has pH value of 6.
 Salts of weak acid and strong base are basic, with a pH value more than 7.
For Example: Sodium carbonate solution has a pH value of 9.

Common Salt
 Chemical name: Sodium chloride
 Common salt is a neutral salt and can be prepared in the laboratory by the reaction of sodium
hydroxide and hydrochloric acid.
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(aq)
 It is an important raw material for products of daily use such as NaOH, baking soda, washing soda and
bleaching powder.

Sodium Hydroxide
 Sodium hydroxide is produced by the electrolysis of an aqueous solution of sodium chloride (called
brine).
 The process is called the chlor-alkali process because of the products formed, i.e. ‘chlor’ for chlorine
and ‘alkali’ for sodium hydroxide.

2NaCl(aq) + 2H2O(aq) → 2NaOH(aq) + H2(g) + Cl2(g)

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Bleaching Powder
 Bleaching powder is manufactured from chlorine gas.
 It is produced by the action of chlorine on dry slaked lime [Ca(OH)2].
Ca(OH)2 + Cl2 → CaOCl2 + H2O
 It is represented as CaOCl2

Uses
 For bleaching cotton and linen in the textile industry and for bleaching wood pulp in the paper industry.
 Used for disinfecting drinking water to make it free of germs.

Baking Soda
 Chemical formula: NaHCO3
 It is produced on a large scale by treating cold and concentrated solution of sodium chloride (brine)
with ammonia and carbon dioxide.
NaCl + H2O + CO2 + NH3 → NH4Cl + NaHCO3
 On heating, it decomposes to give sodium carbonate with the evolution of carbon dioxide.
2NaHCO3 Na2CO3 + H2O + CO2

Uses
 Used as an antacid to treat acidity of the stomach.
 Used to make baking powder, which is used in preparation of cakes, breads, etc.
 Used in soda-acid fire extinguishers.

Washing Soda
 Chemical formula: Na2CO3.10H2O
 Sodium hydrogen carbonate, on heating decomposes to give sodium carbonate with the release of
hydrogen gas. Re-crystallisation of sodium carbonate produces washing soda.
2NaHCO3 Na2CO3 + H2O + CO2
Na2CO3 + 10H2O Na2CO3. 10H2O

Uses
 Used in glass, soap and paper industries.
 Employed in the manufacture of sodium compounds such as borax.

Water Of Crystallisation
 Water molecules which form a part of the structure of a crystal are called water of crystallisation.
 The salts which contain water of crystallisation are called hydrated salts.
 Every hydrated salt has a fixed number of molecules of crystallisation in its one formula unit.
For Example: CuSO4.5H2O, Na2CO3.10H2O, CaSO4.5H2O, and FeSO4.7H2O
 Copper sulphate crystals (CuSO4.5H2O) are blue in colour, and on heating strongly they lose all the
water of crystallisation and form anhydrous copper sulphate, which is white. On adding water to
anhydrous copper sulphate, it gets hydrated and turns blue.
CuSO4.5H2O CuSO4 + 5H2O
CuSO4 + 5H2O CuSO4.5H2O

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Plaster of Paris
Plaster of Paris is prepared by heating gypsum at 373 K. On heating, it loses water molecules and
becomes calcium sulphate hemihydrate (CaSO4.1/2 H2O) which is called Plaster of Paris.
CaSO4.2H2O CaSO4. ½ H2O + 1 ½ H2O
Gypsum Plaster of Paris

Uses
 Used in hospitals as plaster for supporting fractured bones in the right position.
 Used as a fire-proofing material.

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