Revision Notes: Acids, Bases, and Salts

Understanding the fundamentals of chemical properties

1. Introduction to Acids and Bases

  • Acids are substances that taste sour.
  • Bases are substances that taste bitter.

Acids and bases have the ability to nullify each other's effect.

Remedy for acidity: Baking soda solution is suggested to nullify the effect of excess acid.

2. Indicators

Indicators are substances that change color (or odor) in the presence of an acid or a base, allowing us to test them without tasting.

Natural Indicators:

  • Litmus: A purple dye extracted from lichen.
    • Acids change blue litmus to red.
    • Bases change red litmus to blue.
    • Its color is purple when neither acidic nor basic.
  • Turmeric: Becomes reddish-brown with basic soap, turns yellow with water.
  • Other natural materials: Red cabbage leaves, colored petals of flowers like Hydrangea, Petunia, and Geranium.

Synthetic Indicators:

  • Methyl orange.
  • Phenolphthalein.

Olfactory Indicators:

Substances whose odour changes in acidic or basic media.

Examples: onion, vanilla essence, and clove oil.

3. Chemical Properties of Acids and Bases

3.1 How do Acids and Bases React with Metals?

Acids react with metals to produce salt and hydrogen gas.

Acid + Metal → Salt + Hydrogen gas

Example: Reaction of dilute sulphuric acid with zinc granules produces hydrogen gas.

Testing Hydrogen Gas: Hydrogen gas burns with a 'pop' sound.

Bases react with some metals to produce salt and hydrogen gas.

Example: Sodium hydroxide (NaOH) reacts with zinc (Zn) to form sodium zincate (Na₂ZnO₂) and hydrogen gas (H₂).

Note: Such reactions are not possible with all metals.

3.2 How do Metal Carbonates and Metal Hydrogencarbonates React with Acids?

Metal carbonates and metal hydrogencarbonates react with acids to produce a corresponding salt, carbon dioxide gas, and water.

Metal carbonate/Metal hydrogencarbonate + Acid → Salt + Carbon dioxide + Water

Examples:

  • Sodium carbonate (Na₂CO₃) + HCl → Sodium chloride (NaCl) + Water (H₂O) + Carbon dioxide (CO₂).
  • Sodium hydrogencarbonate (NaHCO₃) + HCl → Sodium chloride (NaCl) + Water (H₂O) + Carbon dioxide (CO₂).

Testing Carbon Dioxide Gas: Carbon dioxide gas turns lime water (calcium hydroxide solution) milky due to the formation of white precipitate of calcium carbonate.

Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

Excess CO₂: If excess carbon dioxide is passed, the white precipitate of calcium carbonate dissolves to form soluble calcium hydrogencarbonate.

CaCO₃(s) + H₂O(l) + CO₂(g) → Ca(HCO₃)₂(aq)

Limestone, chalk, and marble are different forms of calcium carbonate.

3.3 How do Acids and Bases React with Each Other? (Neutralisation)

The reaction between an acid and a base to give a salt and water is called a neutralisation reaction.

Base + Acid → Salt + Water

Example: NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)

The effect of a base is nullified by an acid and vice-versa.

This reaction essentially involves the combination of hydrogen ions (H⁺) from the acid and hydroxide ions (OH⁻) from the base to form water:

H⁺(aq) + OH⁻(aq) → H₂O(l)

3.4 Reaction of Metallic Oxides with Acids

Metallic oxides react with acids to give salts and water.

Metal oxide + Acid → Salt + Water

Example: Copper oxide (CuO) + Hydrochloric acid (HCl) → Copper(II) chloride (CuCl₂) + Water (H₂O). The solution turns blue-green due to copper(II) chloride.

Since metallic oxides react like bases with acids, they are considered basic oxides.

3.5 Reaction of Non-metallic Oxides with Bases

Non-metallic oxides react with bases to produce a salt and water.

Example: Calcium hydroxide (base) reacts with carbon dioxide (non-metallic oxide) to produce calcium carbonate (salt) and water.

Since this reaction is similar to a base reacting with an acid, non-metallic oxides are considered acidic in nature.

4. What do all Acids and all Bases have in Common?

4.1 Ions in Aqueous Solutions

All acids have similar chemical properties because they produce hydrogen ions (H⁺) in solution.

These H⁺ ions combine with water molecules to form hydronium ions (H₃O⁺).

HCl + H₂O → H₃O⁺ + Cl⁻

The H⁺(aq) or H₃O⁺ ions are responsible for the acidic properties.

All bases generate hydroxide ions (OH⁻) in water.

Example: NaOH(s) → Na⁺(aq) + OH⁻(aq).

Effect of Water: Hydrogen ions (H⁺) from acids are produced only in the presence of water; they cannot separate from HCl molecules in its absence. Dry HCl gas does not change the color of dry litmus paper.

4.2 Conductivity

Acidic and basic solutions in water conduct electricity because they produce hydrogen and hydroxide ions respectively, which carry the electric current.

Solutions of compounds like glucose and alcohol do not conduct electricity and are not categorized as acids, even though they contain hydrogen, because they do not produce H⁺ ions in solution.

4.3 Dilution

The process of dissolving an acid or a base in water is a highly exothermic process.

Safety Precaution: Acid must always be added slowly to water with constant stirring, not the other way around. Adding water to concentrated acid can cause splashing, burns, and breakage of the glass container due to excessive heat.

Dilution is the process where mixing an acid or base with water results in a decrease in the concentration of ions (H₃O⁺/OH⁻) per unit volume. The acid or base is then said to be diluted.

5. How Strong are Acid or Base Solutions? (pH Scale)

5.1 The pH Scale

A universal indicator (a mixture of several indicators) shows different colors at different concentrations of hydrogen ions in a solution.

The pH scale measures the hydrogen ion concentration in a solution.

The 'p' in pH stands for 'potenz' (German for power).

pH generally ranges from 0 (very acidic) to 14 (very alkaline).

pH Values and Nature of Solution:

  • pH < 7: Acidic solution. Higher hydronium ion (H₃O⁺) concentration means a lower pH value.
  • pH = 7: Neutral solution.
  • pH > 7: Basic/Alkaline solution. As pH increases from 7 to 14, it represents an increase in OH⁻ ion concentration and an increase in the strength of the alkali.

Strength of Acids and Bases:

  • Strong Acids: Produce more H⁺ ions (e.g., HCl, H₂SO₄, HNO₃).
  • Weak Acids: Produce less H⁺ ions (e.g., CH₃COOH).
  • Strong Bases: Produce more OH⁻ ions.
  • Weak Bases: Produce less OH⁻ ions.

Alkalis: Bases that are soluble in water. They are soapy, bitter, and corrosive.

6. Importance of pH in Everyday Life

Plants and Animals are pH Sensitive:

  • Our body works within a narrow pH range of 7.0 to 7.8.
  • Acid Rain: Rainwater with pH less than 5.6 is called acid rain. It lowers river pH, making aquatic life survival difficult.
  • Soil pH: Plants require a specific pH range for healthy growth. Farmers might treat acidic soil with quicklime, slaked lime, or chalk (bases) to adjust pH.

pH in our Digestive System:

  • Stomach produces hydrochloric acid (HCl) for food digestion.
  • During indigestion, excess acid causes pain. Antacids (mild bases like Magnesium hydroxide/Milk of Magnesia) are used to neutralise the excess acid and provide relief.

pH Change as the Cause of Tooth Decay:

  • Tooth decay starts when mouth pH is lower than 5.5.
  • Bacteria degrade sugar and food particles to produce acids, corroding tooth enamel (calcium hydroxyapatite).
  • Basic toothpastes are used to neutralise excess acid and prevent decay.

Self-defence by Animals and Plants through Chemical Warfare:

  • Bee sting leaves an acid (causes pain/irritation), relieved by a mild base like baking soda.
  • Nettle leaves inject methanoic acid (causes burning pain). A traditional remedy is rubbing the area with a dock plant leaf, which is basic in nature.

7. More About Salts

7.1 Family of Salts

Salts are formed by the combination of an acid and a base.

Salts having the same positive or negative radicals belong to a family (e.g., NaCl and Na₂SO₄ are sodium salts; NaCl and KCl are chloride salts).

7.2 pH of Salts

The pH of a salt solution depends on the strength of the acid and base from which it is formed.

  • Strong acid + Strong base = Neutral salt (pH = 7). Example: Sodium chloride (NaCl).
  • Strong acid + Weak base = Acidic salt (pH < 7).
  • Strong base + Weak acid = Basic salt (pH > 7).

7.3 Chemicals from Common Salt (Sodium Chloride, NaCl)

Common salt is a raw material for various important chemicals.

Sodium Hydroxide (NaOH):

  • Produced by the chlor-alkali process (electrolysis of an aqueous solution of sodium chloride, called brine).
  • Products: NaOH (near cathode), chlorine gas (Cl₂) at anode, hydrogen gas (H₂) at cathode. All three products are useful.

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

Bleaching Powder [Ca(ClO)₂]:

  • Produced by the action of chlorine gas (from chlor-alkali process) on dry slaked lime [Ca(OH)₂].

2Ca(OH)₂ + 2Cl₂ → Ca(ClO)₂ + CaCl₂ + 2H₂O

Uses:

  • Bleaching cotton, linen, wood pulp, washed clothes.
  • Oxidising agent in chemical industries.
  • Making drinking water free from germs.

Baking Soda (Sodium Hydrogencarbonate, NaHCO₃):

  • Produced using sodium chloride as a raw material.

NaCl + H₂O + CO₂ + NH₃ → NH₄Cl + NaHCO₃

It is a mild non-corrosive basic salt. This basic nature allows it to neutralise acids.

When heated during cooking, it decomposes:

2NaHCO₃(s) + Heat → Na₂CO₃(s) + H₂O(l) + CO₂(g)

Uses:

  • Making baking powder (mixture of baking soda and a mild edible acid like tartaric acid). When heated or mixed with water, baking powder produces CO₂, making bread/cake soft and spongy.
  • Ingredient in antacids to neutralise excess stomach acid.
  • Used in soda-acid fire extinguishers (reacts to produce carbon dioxide).

Washing Soda (Sodium Carbonate Decahydrate, Na₂CO₃.10H₂O):

  • Obtained by recrystallization of sodium carbonate (which is produced by heating baking soda).

Na₂CO₃ + 10H₂O → Na₂CO₃.10H₂O

It is also a basic salt.

Uses:

  • Glass, soap, and paper industries.
  • Manufacture of sodium compounds like borax.
  • Cleaning agent for domestic purposes.
  • Removing permanent hardness of water.

7.4 Water of Crystallisation

Water of crystallisation is the fixed number of water molecules present in one formula unit of a salt.

Crystals that appear dry often contain water of crystallisation.

When hydrated crystals are heated, this water is removed, and the salt changes color (e.g., blue copper sulphate turns white). Adding water restores the original color.

Examples:

  • Copper Sulphate: CuSO₄.5H₂O (five water molecules).
  • Washing Soda: Na₂CO₃.10H₂O (ten water molecules).
  • Gypsum: CaSO₄.2H₂O (two water molecules).

7.5 Plaster of Paris (Calcium Sulphate Hemihydrate, CaSO₄.½H₂O)

  • Produced by heating gypsum (CaSO₄.2H₂O) at 373 K.

CaSO₄.2H₂O(s) + Heat (at 373 K) → CaSO₄.½H₂O(s) + 1½H₂O(g)

It is a white powder.

When mixed with water, it changes back to gypsum, forming a hard solid mass:

CaSO₄.½H₂O + 1½H₂O → CaSO₄.2H₂O

Uses:

  • Supporting fractured bones (as plaster).
  • Making toys, materials for decoration, and smoothing surfaces.