Acids, Bases, and Salts - Comprehensive Notes

Acids, Bases, and Salts

Acids

• An acid is a compound that, when dissolved in water, yields hydronium ions (H_3O^+) as the only positively charged ions.

• Examples:

  • HCl dissociates into H^+ and Cl^-

  • H^+ combines with H2O to form H3O^+

Classification of Acids

1. Depending on Sources

   • Organic Acids:

     • Usually obtained from plants.

     • Contain carbon and hydrogen atoms.

     • Examples:

       • Acetic acid (Vinegar)

       • Citric acid (Citrus fruits like oranges and lemons)

   • Inorganic (Mineral) Acids:

     • Obtained from minerals.

     • Examples:

       • Hydrochloric acid (HCl)

       • Sulfuric acid (H2SO4)

       • Nitric acid (HNO_3)

2. Depending on Strength

   • Strength of an acid: Depends on the concentration of hydronium ions (H_3O^+) present in the aqueous solution.

   • Strong Acids:

     • Vigorously ionize in aqueous solution, producing a high concentration of hydronium ions (H_3O^+).

     • Examples: HNO3, HCl, H2SO_4

   • Weak Acids:

     • Ionize only partially in aqueous solution to produce ions and molecules.

     • Examples: H2CO3, CH_3COOH, HCOOH

3. Depending on Basicity

   • Basicity of an acid: The number of hydronium ions (H_3O^+) that can be produced by the ionization of one molecule of that acid in aqueous solution.

   • Monobasic Acids:

     • Produce one hydronium ion (H_3O^+) per molecule of the acid.

     • Example: HCl + H2O \rightarrow H3O^+ + Cl^- [Basicity = 1]

   • Dibasic Acids:

     • Produce two hydronium ions (H_3O^+) per molecule of the acid.

     • Examples:

       • H2SO4 + H2O \rightarrow H3O^+ + HSO_4^-

       • HSO4^- + H2O \rightarrow H3O^+ + SO4^{2-} [Basicity = 2]

   • Tribasic Acids:

     • Produce three hydronium ions (H_3O^+) per molecule of the acid.

     • Examples:

       • H3PO4 + H2O \rightarrow H3O^+ + H2PO4^-

       • H2PO4^- + H2O \rightarrow H3O^+ + HPO_4^{2-}

       • HPO4^{2-} + H2O \rightarrow H3O^+ + PO4^{3-} [Basicity = 3]

4. Depending on Concentration

   • Concentrated Acid: Contains a very small amount of water or no water.

   • Dilute Acid: Contains far more water than its own mass.

5. Depending on Molecular Composition

   • Hydracids:

     • Contain hydrogen, a non-metallic element, and no oxygen.

     • Examples: HCl, H_2S, HBr, HI

   • Oxyacids:

     • Contain oxygen, hydrogen, and a non-metallic element.

     • Examples: H2SO4, HNO3, H2CO_3

Preparation of Acids

1. By Synthesis:

   • H2 + Cl2 \rightarrow 2HCl

2. By the Action of Water on Non-metallic or Acidic Oxides:

   • SO3 + H2O \rightarrow H2SO4

   • N2O5 + H2O \rightarrow 2HNO3

3. By Oxidation of Non-metals:

   • S + 6HNO3 \rightarrow H2SO4 + 2H2O + 6NO_2

   • P + H3PO4 \rightarrow H3PO4 + H2O + 5O2

4. By Displacement:

   • NaCl + H2SO4 \rightarrow NaHSO_4 + HCl

   • NaNO3 + H2SO4 \rightarrow NaHSO4 + HNO_3

Properties of Acids

• Physical Properties:

  1. Sour taste in aqueous solution.

  2. Turns blue litmus red.

  3. Some are solids, and some are liquids at room temperature.

  4. Strong mineral acids have corrosive action on the skin and cause painful burns.

  5. Electrolytes: conduct electricity in the aqueous state.

• Chemical Properties:

  1. Reaction with active metals:

     • Mg + 2HCl \rightarrow MgCl2 + H2

  2. Reaction with bases - Neutralization:

     • NaOH + H2SO4 \rightarrow NaNO3 + H2O

  3. Reaction with carbonates and bicarbonates:

     • CaCO3 + 2HCl \rightarrow CaCl2 + H2O + CO2

  4. Reaction with sulphites and bisulphites:

     • CaSO3 + 2HCl \rightarrow CaCl2 + H2O + SO2

     • NaHSO3 + HCl \rightarrow NaCl + H2O + SO_2

  5. Reaction with sulphides:

     • ZnS + 2HCl \rightarrow ZnCl2 + H2S

  6. Reaction with chlorides

  7. Reaction with nitrates:

     • Pb(NO3)2 + 2HCl \rightarrow PbCl2 + 2HNO3

Uses of Some Acids

• Boric acid: Eye wash/antiseptic.

• Citric acid: Food preservation.

• Oxalic acid: Ink stain remover.

• Carbonic acid: Flavored drinks.

Bases

• A base is either a metallic oxide, a metallic hydroxide, or ammonium hydroxide that reacts with hydronium ions of an acid to form salt and water only.

• Basic Oxide: A metallic oxide containing the O^{2-} ion that reacts with an acid to form salt and water.

• Alkalis: A basic hydroxide that, when dissolved in water, produces hydroxyl (OH^-) ions as the only negatively charged ions.

  • NaOH (aq) \rightarrow Na^+ + OH^-

• Note: All alkalis are bases, but not all bases are alkalis.

Classification of Bases

1. On the Basis of Strength

   • Strong Base:

     • Undergoes almost complete ionization in aqueous solution to produce a high concentration of OH^- ions.

     • Example: NaOH (aq) \rightarrow Na^+ (aq) + OH^- (aq)

   • Weak Base:

     • Undergoes only partial ionization in aqueous solution to produce a low concentration of OH^- in solution.

     • Example: NH4OH (aq) \rightarrow NH4^+ (aq) + OH^- (aq)

2. On the Basis of Acidity

   • Acidity of a base: The number of hydroxyl ions (OH^-) that can be produced per molecule of the base in aqueous solution.

     • Monoacidic Base:

       • Produces one hydroxyl ion (OH^-) per molecule.

       • Example: NaOH \rightarrow Na^+ + OH^- [Acidity = 1]

     • Diacidic Base:

       • Produces two hydroxyl ions (OH^-) per molecule.

       • Example: Ca(OH)_2 \rightarrow Ca^{2+} + 2OH^- [Acidity = 2]

     • Triacidic Base:

       • Produces three hydroxyl ions (OH^-) per molecule.

       • Example: Al(OH)_3 \rightarrow Al^{3+} + 3OH^- [Acidity = 3]

3. On the Basis of Composition

   • Concentrated Alkali: An alkali with a relatively high percentage of alkali in its aqueous solution.

   • Dilute Alkali: An alkali with a relatively low percentage of alkali in its aqueous solution.

Preparation of Bases

1. From Metals:

   • 2Mg + O_2 \rightarrow 2MgO

2. By Action of Water or Steam on Reactive Metals:

   • 2Na + 2H2O \rightarrow 2NaOH + H2

3. By the Action of Water on Soluble Metallic Oxides:

   • Na2O + H2O \rightarrow 2NaOH

4. By Double Decomposition:

   • FeCl3 + 3NaOH \rightarrow Fe(OH)3 + 3NaCl

5. By the Action of Oxygen on Metal Sulphides:

   • 2ZnS + 3O2 \rightarrow 2ZnO + 2SO2

6. By Decomposition of Salts:

   • CaCO3 \rightarrow CaO + CO2

Properties of Bases

• Physical Properties:

  1. Sharp and bitter taste.

  2. Change red litmus blue.

  3. Soapy and slippery to touch.

  4. Strong electrolytes.

  5. Mild corrosive action on the skin.

• Chemical properties

  1. Reaction with carbon dioxide:

     • 2NaOH + CO2 \rightarrow Na2CO3 + H2O

  2. Reaction with acids - Neutralisation:

     • Ca(OH)2 + 2HCl \rightarrow CaCl2 + 2H_2O

  3. Reaction with metallic salts:

     • CuSO4 + 2NH4OH \rightarrow (NH4)2SO4 + Cu(OH)2

Uses of Some Bases

• Sodium hydroxide: Manufacture of soaps.

• Potassium hydroxide: Manufacture of salts and soaps, in batteries.

• Magnesium hydroxide: An antacid.

• Magnesia: In making refractory bricks.

pH Value

• Represents the strength of acids and alkalis in terms of hydrogen ion concentration.

• pH of Solution: The negative logarithm (base 10) of the hydrogen ion concentration in moles per liter.

  • pH = -log_{10}(H^+)

• pH Scale: A scale showing the relative strength of acids and alkalis, ranging from 0 to 14.

Indicators

• Complex substances that acquire separate colors in acidic and basic media.

Types of Indicators

• Acid-base indicators:

  • Common indicators like litmus, methyl orange, and phenolphthalein distinguish between acid and basic solutions, but do not determine the strength of the solution.

• Universal indicator:

  • A mixture of organic dyes that gives a definite color change over a wide range of pH.

Salts

• A salt is a compound formed by the partial or total replacement of the ionisable hydrogen atoms of an acid by a metallic ion or an ammonium ion.

Classification of Salts

1. Normal Salts:

   • Formed by the complete replacement of the replaceable hydrogen ion of an acid molecule by a basic radical.

   • Example: HCl + NaOH \rightarrow NaCl + H_2O

2. Acid Salts:

   • Formed by the partial replacement of the replaceable hydrogen ion of an acid molecule by a basic radical.

   • Example: NaOH + H2SO4 \rightarrow NaHSO4 + H2O

3. Basic Salts:

   • Formed by the partial replacement of the hydroxyl group of a di- or tri-acidic base by an acidic radical.

   • Example: Mg(OH)2 + HCl \rightarrow Mg(OH)Cl + H2O

4. Double Salts:

   • Formed by the union of two simple salts which dissolve in water and crystallize.

   • Example: Potash alum: K2SO4 . Al2(SO4)3 . 24H2O

5. Mixed Salts:

   • Salts containing more than one basic or acidic radical.

   • Example: Sodium potassium carbonate NaKCO_3

6. Complex Salts:

   • Salts that, upon dissociation, yield one simple ion and one complex ion.

   • Example: Na[Ag(CN)2] \rightarrow Na^+ + [Ag(CN)2]^-

Preparation of Soluble Salts

• Method 1: Direct Combination

  • Metal + Non-metal → Salt

  • 2Na + Cl_2 \rightarrow 2NaCl

• Method 2: Simple Displacement

  • Active metal + Acid → Salt + Hydrogen

  • Zn + H2SO4 \rightarrow ZnSO4 + H2

• Method 3: Decomposition

  • a. Decomposition of bicarbonates

  • b. Decomposition of carbonates

  • c. Decomposition of chlorides

  • d. Decomposition of nitrates

  • NaHCO3 + HCl \rightarrow NaCl + H2O + CO_2

  • CuCO3 + 2HCl \rightarrow 2CuCl2 + H2O + CO2

• Method 4: Neutralization

  • HNO3 + NaOH \rightarrow NaNO3 + H_2O

Preparation of Insoluble Salts

1. By Direct Combination

   • Reaction: Pb + S \rightarrow PbS

2. By Combination of an Acidic Oxide with a Basic Oxide

   • Reaction: SO2 + CaO \rightarrow CaSO3

3. Double Decomposition

   • Reactions: BaCl2 + H2SO4 \rightarrow BaSO4 + 2HCl

Laboratory Preparation of some Normal and Acid Salts

1. Iron (III) chloride or anhydrous ferric chloride

   • Prepared by passing dry chlorine gas over heated iron.

   • Fe + Cl2 \rightarrow FeCl3

2. Copper (II) sulphate

   • Prepared by the reaction of copper oxide, copper hydroxides or copper carbonates with dilute sulphuric acid.

   • CuO + H2SO4 \rightarrow CuSO4 + H2O

   • Cu(OH)2 + H2SO4 \rightarrow CuSO4 + 2H_2O

   • CuCO3 + H2SO4 \rightarrow CuSO4 + H_2O

   • CuSO4 + 5H2O \rightarrow CuSO4 . 5H2O

3. Zinc sulphate and iron (II) sulphate

   • Prepared by the reaction of metals with dilute sulphuric acid.

   • Zn + H2SO4 \rightarrow ZnSO4 + H2O

   • ZnSO4 + 7H2O \rightarrow FeSO4 . 7H2O

4. Lead chloride

   • Prepared by adding either dilute hydrochloric acid or sodium chloride solution to a solution of lead nitrate.

   • Pb(NO3)2 + 2HCl \rightarrow PbCl2 + 2HNO3

5. Calcium carbonate

   • Prepared by adding sodium carbonate solution to a hot solution of calcium chloride.

   • CaCl2 + Na2CO3 \rightarrow CaCO3 + 2NaCl

6. Sodium bicarbonate

   • Prepared by passing excess carbon dioxide gas through a saturated solution of sodium carbonate.

   • Na2CO3 + CO2 + H2O \rightarrow 2NaHCO_3

7. Neutralisation

   • It is the process by which H+ ions of an acid react completely with the [OH−] ions of a base to give salt and water only.

   • Example: HCl (Acid) + NaOH (Base) \rightarrow NaCl (Salt) + H_2O (water)

Water of Crystallisation

• The amount of water molecules which enter into loose chemical combination with one molecule of the substance on crystallisation from its aqueous solution.

Hydrated Salt

• Salts containing a definite number of water molecules as water of crystallisation.

• Examples: Na2CO3 . 10H2O (washing soda), CuSO4 . 5H_2O (blue vitriol)

Anhydrous Salt

• A salt which does not contain any water of crystallisation.

• Examples: NaCl, NaNO3, Pb(NO3)_2

Deliquescence

• Water-soluble salts that absorb moisture from the atmosphere, dissolve in the same and change into a solution.

• The phenomenon is called deliquescence, and the salts are deliquescent.

• Examples: CaCl2, MgCl2, ZnCl_2

Efflorescence

• Crystalline hydrated salts that lose their water of crystallisation partly or completely on exposure to the atmosphere and change into a powder.

• This phenomenon is called efflorescent, and the salts are efflorescent.

• Examples: CuSO4 . 5H2O, MgSO4 . 7H2O, Na2CO3 . 10H_2O