Acids, Bases and Salts

(i) Pink solution becomes colourless.

(ii) Orange solution changes to pink colour.

(iii) Purple solution changes to red colour.

Copper hydroxide [Cu(OH)₂]
Reason — Copper hydroxide [Cu(OH)₂] is an example of insoluble base and it is not an alkali.

A solution X turns blue litmus red, so it must contain (i) hydronium ions; another solution Y turns red litmus blue and therefore, must contain (ii) hydroxide ions. When solutions X and Y are mixed together the products will be a (iii) salt and (iv) water. If a piece of magnesium were put into solution X, (v) hydrogen gas would be evolved.

Pb(NO₃)₂ + 2NaCl ⟶ PbCl₂ + 2NaNO₃

Alkali is a base which is soluble in water.

Acetic acid contains four hydrogen atoms.

Copper [II] oxide
Reason — Copper [II] oxide is black in colour and the following reaction takes place when it is treated with dilute sulphuric acid —

CuO + H₂SO₄ ⟶ CuSO₄ + H₂O

CuSO₄ is a blue coloured soln.

(i) Solution B — weak acid
Reason — Weak Acid is an acid which dissociates only partially in aqueous solution thereby producing a low concentration of hydrogen [H⁺] ions [or H₃O⁺ ions]. For example — CH₃COOH ⇌ CH₃COO^- + H⁺ [contains molecules and ions]

(ii) Solution C — strong alkali
Reason — Alkalis react with certain salt solutions to precipitate insoluble hydroxide. Hence,
ZnSO₄ + 2NaOH ⟶ Na₂SO₄ + Zn(OH)₂ [gelatinous white precipitate]

(iii) Solution B — weak acid
Reason — Anhydrous acetic acid on cooling forms crystals of glacial acetic acid and acetic acid is a weak acid.

(iv) Ammonium hydroxide (NH₄OH)

PbCO₃ + 2HNO₃ ⟶ Pb(NO₃)₂ + H₂O + CO₂

Pb(NO₃)₂ + Na₂SO₄ ⟶ PbSO₄ + 2NaNO₃

Neutralization — It is the process due to which [H⁺] ions of an acid react completely or combine with [OH^-] ions of a base to give salt and water only.

Acid + Base ⟶ Salt + Water

HCl + NaOH ⟶ NaCl + H₂O

H⁺Cl^- + Na⁺OH^- ⟶ Na⁺Cl^- + H₂O

[H⁺ (aq) + OH^- (aq) ⇌ H₂O (l)]

(i) D: Lead [II] chloride

(ii) E: Sodium chloride

(i) Formic acid

(ii) Tetrammine copper (II) sulphate

(i) ZnSO₄ + (NH₄)₂CO₃ ⟶ (NH₄)₂SO₄ + ZnCO₃

(ii) ZnCO₃ + 2HNO₃ ⟶ Zn(NO₃)₂ + H₂O + CO₂

Weak acid
Reason — An acid which dissociates only partially in aqueous solution thereby producing a low concentration of hydrogen [H⁺] ions [or H₃O⁺ ions] is a weak acid.

Example — CH₃COOH ⇌ CH₃COO^- + H⁺ [contains molecules and ions]

1. CuO + H₂SO₄ ⟶ CuSO₄ + H₂O

2. 2Fe + 3Cl₂ ⟶ 2FeCl₃

3. 2KOH + H₂SO₄ [dil.] ⟶ K₂SO₄ + 2H₂O

4. PbCO₃ + 2HNO₃ ⟶ Pb(NO₃)₂ + H₂O + CO₂
   Pb(NO₃)₂ + 2NaCl ⟶ 2NaNO₃ + PbCl₂

Pb(NO₃)₂ + 2NaCl ⟶ PbCl₂ + 2NaNO₃

(i) Sodium nitrate — B : Neutralization

(ii) Iron (III) chloride — E : Direct synthesis

(iii) Lead chloride — D : Double decomposition

(iv) Zinc sulphate — A : Simple displacement

(v) Sodium hydrogen sulphate — C : Decomposition by acid

1. Acid salt — B : Sodium hydrogen sulphate
2. Double salt — A : Ferrous ammonium sulphate

(i) A solultion M turns blue litmus red, so it must contain (i) hydronium ions ; another solution O turns red litmus blue and hence, must contain, (ii) hydroxide ions.

(ii) When solution M and O are mixed together, the prod­ucts will be (iii) salt and (iv) water.

(iii) If a piece of magnesium was put into a solution M, (v) hydrogen gas would be evolved.

(i) Acid salt

(ii) Water of crystallization

1. Sodium sulphate
   Na₂CO₃ + H₂SO₄ (dil.) ⟶ Na₂SO₄ + H₂O + CO₂

2. Zinc carbonate
   Zn(NO₃)₂ + Na₂CO₃⟶ 2NaNO₃ + ZnCO₃

3. Copper (II) sulphate
   CuCO₃ + H₂SO₄ (dil.) ⟶ CuSO₄ + H₂O + CO₂

4. Iron (II) sulphate
   Fe + H₂SO₄ (dil.) ⟶ FeSO₄ + H₂

Acetic acid

The basicity of acetic acid is 1.

Hydronium ion is the stable positive ion formed when an acid dissolves in water. Its structure is shown below:

Salt S is Copper sulphate CuSO₄

CuO + H₂SO₄ (dil.) ⟶ CuSO₄ + H₂O

1. Lead sulphate from lead carbonate.
   PbCO₃ + 2HNO₃ ⟶ Pb(NO₃)₂ + H₂O + CO₂
   Pb(NO₃)₂ + Na₂SO₄ ⟶ 2NaNO₃ + PbSO₄

2. Sodium sulphate using dilute sulphuric acid.
   Na₂CO₃ + H₂SO₄(dil.) ⟶ Na₂SO₄ + H₂O+ CO₂

3. Copper chloride using copper carbonate.
   CuCO₃ + 2HCl (dil) ⟶ CuCl₂ + H₂O + CO₂

2Fe + 3Cl₂ ⟶ 2FeCl₃

AgCl

SO₂

The following reaction takes place.

SO₂ + H₂O ⟶ H₂SO₃

Higher the pH value of a solution, the more alkaline it is.

(i) Pb(NO₃)₂ from PbO — (C) Neutralization

(ii) MgCl₂ from Mg — (A) Simple displacement

(iii) FeCl₃ from Fe — (E) Combination

(iv) NaNO₃ from NaOH — (B) Titration

(v) ZnCO₃ from ZnSO₄ — (D) Precipitation

When a metallic oxide is dissolved in water, the solution formed has a high concentration of OH^- ions.

(i) An alkali
Reason — The pH of alkali solutions is more than 7 therefore in order to increase the pH value of a neutral solution, an alkali should be added.

(ii) Direct combination
Reason — Anhydrous iron [III] chloride is prepared by direct combination of iron and chloride as follows —
2Fe + 3Cl₂ ⟶ 2FeCl₃

(i) CuSO₄ + Na₂CO₃ ⟶ Na₂SO₄ + CuCO₃

(ii) 2NH₄OH + H₂SO₄ ⟶ (NH₄)₂SO₄ + 2H₂O

Efflorescence

It absorbs moisture from the atmosphere, dissolve in the same and change into a solution. This property is known as deliquescence

The salt prepared by the method of direct combination is iron [III] chloride

(i) Q
Reason — On a pH scale, acids have pH less than 7 whereas weak acids have pH towards 7. Hence, Q will be a weak acid with pH 5.2 .

(ii) R
Reason — On a pH scale, alkali have pH more than 7 and alkalinity increases as the pH value moves away from 7. Hence, R will be a strong alkali with pH 12.2 .

(i) CuCO₃ + H₂SO₄ ⟶ CuSO₄ + H₂O + CO₂

(ii) ZnSO₄ + (NH₄)₂CO₃ ⟶ (NH₄)₂SO₄ + ZnCO₃

Acid
Reason — An acid is a compound which when dissolved in water yields hydronium ions [H₃O⁺] as the only positively charged ion.

HCl (aq) ⇌ H⁺ + Cl^-

H⁺ + H₂O ⇌ H₃O⁺ [hydronium ion]

HCl + H₂O ⇌ H₃O⁺ + Cl^-

(i) Solution C : pH value 7

(ii) Solution B : pH value 2

(iii) Solution A : pH value 12

(i) Lead chloride — B: Precipitation

(ii) Iron [II] Sulphate — D: Substitution

(iii) Sodium nitrate — A: Neutralization

(iv) Iron [III] chloride — C: Direct combination

Ferrous chloride

pH of acetic acid is greater than dilute sulphuric acid. So, acetic acid contains low concentration of H⁺ ions.

1. Preparation of copper sulphate by neutralization
   CuO + H₂SO₄ ⟶ CuSO₄ + H₂O

2. Preparation of zinc carbonate by precipitation
   Zn(NO₃)₂ + Na₂CO₃ ⟶ 2NaNO₃ + ZnCO₃

3. Preparation of ammonium sulphate by titration
   2NH₄OH + H₂SO₄ ⟶ (NH₄)₂SO₄ + 2H₂O

(i) Acid — An acid is a compound which when dissolved in water yields hydronium ions [H₃O⁺] as the only positively charged ion.

HCl (aq) ⇌ H⁺ + Cl^-

H⁺ + H₂O ⇌ H₃O⁺ [hydronium ion]

HCl + H₂O ⇌ H₃O⁺ + Cl^-

(ii) Base — A base is a compound which reacts with hydronium ions of an acid to give salt and water only.

CuO + 2HCl ⟶ CuCl₂ + H₂O

Cu(OH)₂ + H₂SO₄ ⟶ CuSO₄ + 2H₂O

• Bases are oxides or hydroxides of a metal [including ammonium hydroxide]

• Examples of insoluble bases [i.e., not alkalis] — ZnO, PbO, CuO, Fe(OH)₂, Pb(OH)₂, Cu(OH)₂

(iii) Alkali — An alkali is a compound which when dissolved in water yields hydroxyl ions [OH^-] as the only negatively charged ions.

NaOH [aq.] ⇌ Na⁺ + OH^- [Hydroxyl or hydroxide ion]

Alkali is a base, soluble in water. [All alkalis are bases, but all bases are not alkalis.]

Examples of soluble bases [i.e., alkalis] — KOH, NaOH [strong alkalis] , Ca(OH)₂, NH₄OH (weak alkalis).

(iv) Neutralization — It is the process due to which [H⁺] ions of an acid react completely or combine with [OH^-] ions of a base to give salt and water only.

Acid + Base ⟶ Salt + Water

HCl + NaOH ⟶ NaCl + H₂O

H⁺Cl^- + Na⁺OH^- ⟶ Na⁺Cl^- + H₂O

[H⁺ (aq) + OH^- (aq) ⇌ H₂O (l)]

(i) Difference between organic and inorganic acids are as follows :

(ii) Difference between hydracids and oxyacids are as follows :

Strength of acids depends on the concentration of hydronium ion [H₃O⁺] present in an aqueous solution of an acid.

Strength of alkali depends on the concentration of the hydroxyl ions [OH^-] present in an aqueous solution of the alkali.

(i) Differences between strong and weak acid are as follows :

(ii) Differences between strong alkali and and weak alkali are as follows :

(i) When HCl is dissolved in water, it is ionised into hydrogen ion [or H₃O⁺ ion] and chloride ion.

HCl ⟶ H⁺ + Cl^-

The H⁺ cannot exist independently, therefore, it combines with water molecule to form hydronium ion (H₃O⁺)

H⁺ + H₂O ⟶ H₃O⁺

(ii) When HNO₃ is dissolved in water, it is ionised into hydronium ion and nitrate ion.

HNO₃ ⟶ H⁺ + NO₃^-

H⁺ + H₂O ⟶ H₃O⁺

(iii) When H₂SO₄ is dissolved in water, it is ionised into hydronium ion and sulphate ion.

H₂SO₄ ⟶ 2H⁺ + SO₄^2-

H⁺ + H₂O ⟶ H₃O⁺

(iii) When CH₃COOH is dissolved in water, it is ionised into hydronium ion and acetate ion.

CH₃COOH ⟶ CH₃COO^- + H⁺

H⁺ + H₂O ⟶ H₃O⁺

(iii) When NaOH is dissolved in water, it is ionised into sodium ion and hydroxyl ion.

NaOH ⟶ Na⁺ + OH^-

(iv) When NH₄OH is dissolved in water, it is ionised into ammonium ion and hydroxyl ion.

NH₄OH ⟶ NH₄⁺ + OH^-

Dilute HCl is a stronger acid than concentrated H₂CO₃
Reason — HCl dissociates almost completely in aqueous solution and produces a high concentration of H⁺ ions and Cl^- ions, hence is a strong acid. Whereas, H₂CO₃ is a weak acid because it dissociates partially yielding H⁺ ions and bicarbonate HCO₃^- ion and hence, contains ions as well as molecules. Therefore, dil. HCl is a stronger acid than conc. H₂CO₃.

Basicity of acid — is the number of hydrogen ions [H⁺] which can be produced per molecule of the acid in aq. soln. Acetic acid [CH₃COOH] ionises in aq. soln. and gives one hydrogen ion per molecule of the acid, hence acetic acid is monobasic i.e., it's basicity is one.

Acidity of base — is the number of hydroxyl ions [OH^-] which can be produced per molecule of the base in aq. soln. Calcium hydroxide Ca(OH)₂ ionises in aq. soln. and gives two hydroxyl ions per molecule of the base, hence calcium hydroxide is a diacidic base i.e., it's acidity is two.

Sulphuric acid (H₂SO₄) is a dibasic acid as :

1. It ionises in aq. soln. to produce two hydrogen ions per molecule of the acid.
2. It dissociate in two steps in aq. soln. as shown below:
   H₂SO₄ + H₂O ⇌ H₃O⁺ + HSO₄^-
   HSO₄^- + H₂O ⇌ H₃O⁺ + SO₄^2-
   H₂SO₄ + 2H₂O ⇌ 2H₃O⁺ + SO₄^2-
3. It contains two replaceable hydrogen ions per molecule of the acid so forms two types of salt [acid and normal salt] :
   NaOH + H₂SO₄ ⟶ NaHSO₄ (Acid Salt) + H₂O
   2NaOH + H₂SO₄ ⟶ Na₂SO₄ (Normal Salt) + 2H₂O

Arrhenius Theory — Acids are substances which dissociate in aqueous solution to give H⁺ ions.
Strong acids dissociate almost completely, while weak acids dissociate partially.

Lowry – Bronsted's theory — Acids are proton donors and bases are proton acceptors [proton = H⁺].
HCl [aq.] ⟶ H⁺ + Cl^- [acid - proton donors]
NH₃ + H⁺ ⟶ NH₄⁺ [bases - proton acceptors]

Oxygen atom in water has two 'lone pair of electrons' implies that two pairs of electrons on oxygen are not shared with any other atom as shown below:

Formation of hydronium ion

Formation of ammonium ion

1. Acidic oxides dissolve in water to give an acid.
   SO₃ + H₂O ⟶ H₂SO₄
2. Basic oxides [soluble] dissolve in water to give a base i.e., alkali.
   K₂O + H₂O ⟶ 2KOH

(i) Chemical properties of a solution containing H⁺ (acids) are as follows —

1. Neutralization — acids neutralizes base to give salt and water only.
   Example : CuO + H₂SO₄ ⟶ CuSO₄ + H₂O
2. Reaction with active metals — Acids react with active metals [e.g., Mg, Al, Zn, Fe] to liberate hydrogen.
   Example : Zn + 2HCl ⟶ ZnCl₂ + H₂

(ii) Chemical properties of a solution containing OH^- (bases/alkalis) are as follows —

1. Neutralization — Alkalis neutralize acids to form salt and water.
   Example :
   Ca(OH)₂ + 2HCl ⟶ CaCl₂ + 2H₂O
2. Reaction with metallic salt — Alkalis react with certain matallic salt solutions to precipitate insoluble hydroxide.
   Example :
   FeCl₃ + 3NaOH ⟶ 3NaCl + Fe(OH)₃ ↓ red brown ppt.

(i) Decomposition of a metallic chloride —

$\text{NaCl} + \text{H}_2\text{SO}_4 [\text{conc.}] \xrightarrow{\lt 200 ^{\circ}\text{C}}\text{NaHSO}_4 + \text{HCl}$

(ii) Decomposition of a metallic nitrate with conc. H₂SO₄ —

$\text{NaNO}_3 + \text{H}_2\text{SO}_4 [\text{conc.}] \xrightarrow{\lt 200 ^{\circ}\text{C}}\text{NaHSO}_4 + \text{HNO}_3$

As H₂SO₄ is a less volatile acid and displaces the more volatile acid on heating with the salt, hence the respective acids are formed.

AB [salt I] + HX [acid I] ⟶ AX [salt II] + HB [acid II]

2NaHCO₃ + H₂SO₄ [dil.] ⟶ Na₂SO₄ + 2H₂O + 2CO₂ ⟶ H₂CO₃ (unstable)

Na₂CO₃ + H₂SO₄ [dil.] ⟶ Na₂SO₄ + H₂O + CO₂ ⟶ H₂CO₃ (unstable)

Ammonia is evolved when an ammonium salt and alkali are heated because a less volatile base (e.g.,NaOH) displace the more volatile base, NH₄OH and we get the products NH₃ and H₂O.

NH₄Cl + NaOH ⟶ NaCl + H₂O + NH₃

pH is defined as the negative logarithm [to the base 10] of the hydrogen ion concentration expressed in moles/litre. Thus, pH = -log₁₀H⁺. It represents the strength of acids and alkalis, expressed in terms of hydrogen ion concentration [H⁺ aq.]

(i) When, H⁺ ions = OH^- ions, the solution is neutral with pH = 7.

(ii) When the solution evolves CO₂ when heated with Na₂CO₃, it is acidic in nature with pH less than 7.

(iii) when, OH^-ions > H⁺ aq. ions, the solution is basic in nature with pH more than 7.

Litmus is a common acid-base indicator and not a universal indicator. It only indicates whether a solution is acidic or alkaline. It cannot be utilized for determining the strength of the acidic or alkaline solution.

(i) A neutral litmus is purple in colour. In presence of an acidic medium, the colour of neutral litmus changes from purple to red.

(ii) In an alkaline medium, colour of neutral litmus changes from purple to blue.

(i) The pH paper changes to yellow colour on a slightly acidic soil.

(ii) The pH paper changes to blue colour on a slightly alkaline soil.

(iii) The pH paper changes to green colour in a dairy milk.

(iv) The pH paper changes to green colour in human blood tested for medical diagnosis.

(i) Salt — A salt is a compound formed by partial or complete replacement of the replaceable hydrogen ions of an acid by a metallic ion or ammonium ion [basic radical].

NaOH + H₂SO₄ ⟶ NaHSO₄ + H₂O [Partial replacement]

2NaOH + H₂SO₄ ⟶ Na₂SO₄ + 2H₂O [Complete replacement]

(ii) Normal salt — The salt formed by complete replacement of the replaceable hydrogen ion of an acid molecule by a basic radical [metallic or ammonium ion].
For example,

2NaOH + H₂SO₄ ⟶ Na₂SO₄ + 2H₂O

2NaOH + H₂SO₃ ⟶ Na₂SO₃ + 2H₂O

[Both H ions in sulphuric and sulphurous acid are replaced by metallic radical — sodium. ]

(3) Acid salt — The salt formed by partial replacement of the replaceable hydrogen ion of an acid molecule by a basic radical [metallic or ammonium ion]. For example,

NaOH + H₂SO₄ ⟶ NaHSO₄ + H₂O

NaOH + H₂SO₃ ⟶ NaHSO₃ + H₂O

[Only one H ion in sulphuric and sulphurous acid is replaced by metallic radical — sodium. ]

(i) Formation of a basic salt — A basic salt is formed by partial replacement of hydroxyl radicals of a diacidic or triacidic base with an acid radical.

(ii) Components of a basic salt — A basic salt contains a cation [metallic], a hydroxyl ion [of a base] and an anion [of an acid].

For example — Basic copper nitrate Cu[OH]NO₃ , basic copper chloride Cu[OH]Cl

Acid, normal or basic salt —

(i) bleaching powder — Normal salt (Mixed salt)

(ii) potassium mercuric iodide — Normal salt (Complex salt)

(iii) sodium sulphite — Normal salt

(iv) sodium hydrogen sulphite — Acid salt

(v) sodium silver cyanide — Normal salt (Complex salt)

(vi) basic lead nitrate — Basic salt

(vii) potassium zincate — Normal salt

(viii) alum — Normal salt (Double salt)

(ix) calcium bicarbonate — Acid salt

(x) basic copper chloride — Basic salt

(xi) trisodium phosphate — Normal salt

Insoluble in Water

(i) Three sulphates insoluble in water are:

1. Lead sulphate (PbSO₄)
2. Calcium sulphate (CaSO₄)
3. Barium sulphate (BaSO₄).

(ii) Three chloride insoluble in water are:

1. Silver chloride (AgCl)
2. Lead chloride (PbCl₂)
3. Mercury chloride (Hg₂Cl₂).

Soluble in Water

(i) Two oxides soluble in water are:

1. Sodium oxide (Na₂O)
2. Potassium oxide (K₂O)

(ii) Two carbonates soluble in water are:

1. Sodium carbonate (Na₂CO₃)
2. Ammonium carbonate [(Na₄)₂CO₃]

Method of preparation are as follows :

(i) Zn(NO₃)₂ — By action of dilute acids on carbonate.

(ii) NH₄Cl — By neutralisation of an alkali (titration)

(iii) ZnSO₄ — By displacement of active metal and acid

(iv) ZnS — Direct combination (Synthesis)

(v) CaCO₃ — By precipitation (double decomposition)

(vi) FeCl₃ — Direct combination (Synthesis)

(vii) PbCl₂ — Precipitation (double decomposition)

(viii) Pb(NO₃)₂ — Action of dil. acid on carbonates and bicarbonates

(a) Using dil. H₂SO₄ :

1. Preparation of CuSO₄:
   CuO + H₂SO₄ ⟶ CuSO₄ + H₂O
2. Preparation of NaHSO₄:
   NaOH + H₂SO₄ ⟶ NaHSO₄ + H₂O
3. Preparation of Na₂SO₄:
   Na₂CO₃ + H₂SO₄ ⟶ Na₂SO₄ + H₂O + CO₂
4. Preparation of FeSO₄:
   Fe + H₂SO₄ ⟶ FeSO₄ + H₂
5. Preparation of BaSO₄:
   BaCl₂ + H₂SO₄ ⟶ BaSO₄ + 2HCl
6. Preparation of PbSO₄:
   Pb(NO₃)₂ + H₂SO₄ ⟶ PbSO₄ + 2HNO₃

(b) Using conc. H₂SO₄ :

1. Preparation of NaHSO₄:

$\text{NaCl} + \text{H}_2\text{SO}_4 [\text{conc.}] \xrightarrow{\lt 200 ^{\circ}\text{C}}\text{NaHSO}_4 + \text{HCl}$

2. Preparation of CuSO₄:
   CuO + H₂SO₄ ⟶ CuSO₄ + H₂O

ZnO + 2HCl ⟶ ZnCl₂ + H₂O

ZnCl₂ + Na₂CO₃ ⟶ ZnCO₃ + 2NaCl

Dissolve zinc oxide in dil. HCl. Add to it a saturated solution of Na₂CO₃. The precipitate formed by the interchange of radicals is filtered. It is dried to obtain the zinc carbonate.

(i) ZnCO₃ + 2HNO₃ (dil) ⟶ Zn(NO₃)₂ + H₂O + CO₂

(ii) 2KHCO₃ + H₂SO₄ (dil) ⟶ K₂SO₄ + 2H₂O + 2CO₂

1. Decomposition of calcium bicarbonate by dil. HCl
   Ca(HCO₃)₂ + 2HCl ⟶ CaCl₂ + 2H₂O+ 2CO₂
2. Decomposition of calcium carbonate by dil. HNO₃
   CaCO₃ + 2HNO₃ ⟶ Ca(NO₃)₂ + H₂O+ CO₂
3. Decomposition of sodium sulphite by dil. H₂SO₄
   Na₂SO₃ + H₂SO₄ (dil.) ⟶ Na₂SO₄ + H₂O + SO₂
4. Decomposition of zinc sulphide by dil. H₂SO₄
   ZnS + H₂SO₄ ⟶ ZnSO₄ + H₂S

(i) As K₂CO₃ is a salt of a strong base (KOH) and weak acid (H₂CO₃), it hydrolyses in water to give alkaline solutions. They have pH more than 7, hence it will have no effect on blue litmus.

(ii) KCl is a salt of a strong acid (HCl) and a strong base (KOH). Hence it's aqueous solution will be neutral in nature and will have no effect on blue litmus.

(iii) NH₄NO₃ is a salt of a strong acid (HNO₃) and weak base (NH₄OH), it hydrolyses in water to give an acidic solution. They have pH less than 7, hence will turn blue litmus red.

An example of an acid derived from a mineral is nitric acid.

An example of a base which is not an alkali is zinc hydroxide.

An example of a strong acid is dilute sulphuric acid.

An example of a weak alkali is calcium hydroxide solution.

An acid having basicity 1 is acetic acid.

An acid obtained by dissolving sulphur trioxide in water is sulphuric acid.

A volatile acid obtained when nitre reacts with non­-volatile concentrated sulphuric acid on heating is nitric acid.

A base obtained when lead nitrate undergoes thermal decomposition is lead (II) oxide.

An acid obtained when concentrated nitric acid is heated with sulphur is sulphuric acid.

The more volatile acid obtained when the less volatile acid reacts with sodium bicarbonate is carbonic acid.

The insoluble base obtained when sodium hydroxide reacts with iron (III) chloride is Iron (III) hydroxide.

A solution whose pH is above 7 is liquor ammonia.

The salt formed when sulphuric acid reacts with excess caustic soda solution is sodium sulphate.

An example of an acid salt is Na₂HPO₄.

An example of a soluble salt is CaCl₂.

An example of an insoluble salt is MgCO₃.

A salt prepared by neutralization in which titration is involved is NH₄Cl.

An insoluble salt prepared by direct combination or synthesis is FeS.

A salt prepared by precipitation i.e. by double decomposition of two salt solutions is PbSO₄.

A salt prepared by simple displacement i.e. action of dilute acid on a metal is AlCl₃.

Decomposition of calcium hydrogen carbonate with dil. HCl results in formation of calcium chloride.

Action of dilute acid on a metallic sulphide results in evolution of H₂S gas.

A salt which on hydrolysis produces a neutral solution is sodium chloride .

1. Ammonium Hydroxide (NH₄OH) is a basic solution which does not contain a metallic element.
2. Sodium acetate (CH₃COONa) is a normal salt of sodium formed from acetic acid.
3. Ammonium hydroxide (NH₄OH) is a base which reacts with an acid to give a salt which on hydrolysis gives a slightly acidic solution.
4. Hydrogen ion (H⁺) is an ion which combines with a polar covalent molecule to form an ammonium ion.
5. Aluminium chloride (AlCl₃) is a soluble salt formed by direct combination between a light metal & a greenish yellow gas where, light metal is Al and greenish yellow gas is Cl.

1. B: Monobasic acid
2. D: Weak acid
3. A: Hydracid
4. C: Less volatile acid
5. F: Dibasic acid

1. E: Precipitation (Double decomposition)
2. C: Simple displacement — active metal and dil. acid
3. B: Neutralisation — alkali and dil. acid
4. A: Neutralisation — insoluble base and dil. acid
5. D: Direct combination

1. Calcium oxide ⟶ Calcium chloride ⟶ Calcium carbonate
   1. CaO + 2HCl ⟶ CaCl₂ + H₂O
   2. CaCl₂ + Na₂CO₃ ⟶ 2NaCl + CaCO₃
2. Zinc sulphide ⟵ Zn ⟶ Zinc sulphate
   1. Zn + S ⟶ ZnS
   2. Zn + H₂SO₄ ⟶ ZnSO₄ + H₂
3. Iron [II] chloride ⟵ Fe ⟶ Iron [III] chloride
   1. Fe + 2HCl ⟶ FeCl₂ + H₂
   2. 2Fe + 3Cl₂ ⟶ 2FeCl₃
4. Lead [II] oxide ⟶ Lead nitrate ⟶ Lead sulphate
   1. PbO + 2HNO₃ ⟶ Pb(NO₃)₂ + H₂O
   2. Pb(NO₃)₂ + Na₂SO₄ ⟶ 2NaNO₃ + PbSO₄
5. Copper [II] oxide ⟶ Copper [II] sulphate ⟵ Copper [II] hydroxide
   1. CuO + H₂SO₄ ⟶ CuSO₄ + H₂O
   2. Cu(OH)₂ + H₂SO₄ ⟶ CuSO₄ + 2H₂O

1. Burette
2. Dil. sulphuric acid (H₂SO₄) is substance 'X' placed in 'A' and sodium hydroxide (NaOH) is the substance 'Y' placed in B.
3. Titration is conducted to determine the completion of the neutralization reaction, i.e. to determine the amount of sulphuric acid required to neutralize a known amount of sodium hydroxide.
4. Aqueous sodium sulphate (Na₂SO₄.10H₂O) is transferred to the evaporating dish and evaporated to point of crystallisation for obtaining the salt.
5. As copper (II) oxide is not soluble in water hence titration is not conducted for the preparation of copper (II) sulphate crystals by neutralization.

Strength of an acid is determined by the amount of hydronium ions it produces in it's aqueous solution. As dilute sulphuric acid dissociates almost completely thereby producing a high concentration of hydronium ions hence is a stronger acidic than concentrated sulphuric acid.

As ammonium chloride is a salt of a weak alkali and strong acid, hence it's aqueous solution is acidic in nature. On the other hand, sodium chloride is a product of strong alkali and strong acid hence, it's aqueous solution is neutral in nature.

In the preparation of an insoluble salt from another insoluble salt by precipitation, if dil. sulphuric acid is directly used then it forms an insoluble precipitate and slows the reaction.

For example — Direct addition of dil. sulphuric acid to PbCO₃ is an impractical method of preparing lead sulphate since PbSO₄ is insoluble and forms a coating on PbCO₃, thereby the reaction slowly comes to a stop.

Acetic acid (CH₃COOH) is a monobasic acid and contains only one replaceable hydrogen atom, hence it can only form normal salt and not acid salts.

Sulphurous acid (H₂SO₃) is a dibasic acid, i.e., it contains two replaceable hydrogen ions per molecule of the acid. Hence, it can form normal salt as well as acid salt on reaction with an alkali.