Properties of Ionic Compounds

Physical Properties of Ionic Compounds — Introduction

Ionic compounds — familiar in everyday life: salt, lime, baking soda. All their properties trace to one root cause: the crystal lattice of ions.

The Core Idea

In ionic compounds:

• There are no separate 'molecules'.
• Ions are bonded in a three-dimensional lattice.
• One Na⁺ — surrounded by 6 Cl⁻.
• One Cl⁻ — surrounded by 6 Na⁺.
• (In NaCl — face-centered cubic structure.)

This structure — the root of all properties.

Main Physical Properties

1. Physical State:

• Ionic compounds are typically solid.
• Hard but brittle.
• Crystalline.

2. Melting Point and Boiling Point:

• High melting and boiling points.
• Because of strong attraction between ions.

Examples:

• NaCl: melting point 801°C
• MgO: melting point 2852°C
• Al₂O₃: melting point ~2050°C
• CaO: melting point 2572°C

3. Brittleness:

• Ionic crystals are brittle.
• Break with slight force.
• Why? Pressure shifts ions, bringing like charges face-to-face — repulsion breaks the crystal.

4. Hardness:

• Generally hard.
• Strong forces between ions.

Properties Linked to Strong Ionic Attraction

Coulomb's Law: $F = \frac{kq_1q_2}{r^2}$

That is:

• More charge → more attraction → higher melting point.

Comparison:

Compound	Charges	Melting Point
NaCl	+1, -1	801°C
MgO	+2, -2	2852°C
Al₂O₃	+3, -2	2050°C

Melting points of MgO and CaO are much higher than NaCl — because charges are larger.

Solubility

Solubility in Water

Most ionic compounds — are soluble in water.

Examples:

• NaCl, KCl, KNO₃ — fully soluble.
• $CuSO_4$, $ZnSO_4$, $FeSO_4$ — soluble.
• $MgCl_2$, $CaCl_2$ — soluble.

Insoluble ones:

• AgCl, BaSO₄, PbI₂, CaCO₃ — sparingly soluble or insoluble.

Why Soluble in Water?

Nature of water: Water is a polar molecule.

• $H$ end — partial positive.
• $O$ end — partial negative.

Mechanism of dissolution:

1. Water's $\delta^-$ ($O$) — surrounds $Na^+$.
2. Water's $\delta^+$ ($H$) — surrounds $Cl^-$.
3. Ions are pulled out of crystal lattice.
4. Aqueous solution is formed.

This process = Hydration.

Solubility in Organic Solvents

Kerosene, petrol, benzene — non-polar.

Ionic compounds are insoluble in these.

Rule: 'Like dissolves like'.

• Polar substance — in polar solvent.
• Non-polar substance — in non-polar solvent.

An Interesting Fact

Why don't oil and water mix?

• Oil: non-polar (organic).
• Water: polar.
• Rule: polar and non-polar don't mix.

Solubility Table — Some Famous Ionic Compounds

Compound	In Water	In Kerosene
NaCl	✓ Highly soluble	✗
KNO₃	✓	✗
CuSO₄	✓ (blue)	✗
AgCl	✗ Insoluble	✗
CaCO₃	✗	✗
MgO	✗ Sparingly	✗

Electrical Conductivity

Main Question: When are ionic compounds electrical conductors?

Ionic compounds — in three states:

1. In solid state: Not electrical conductors. Why? Ions are fixed in crystal lattice — cannot move.

2. In molten state: Electrical conductors. Why? On melting, lattice breaks — ions become mobile.

3. In aqueous solution: Electrical conductors. Why? Ions move freely in water.

Mechanism of Conduction

Positive ions (cations) → move to negative end (cathode). Negative ions (anions) → move to positive end (anode).

This movement of ions = electric current.

Practical Demonstration

Simple Experiment:

• A battery, bulb, and two electrodes.
• Three media: dry NaCl solid, molten NaCl, NaCl solution.
• In which will the bulb glow?

Answer:

• Dry NaCl: No.
• Molten NaCl: Yes! (brightly)
• NaCl solution: Yes.

Electrolysis

This principle — the basis of many industrial processes:

1. Extraction of metals (Na, K, Al — by electrolysis) (Section 7)
2. Electroplating (Section 8)
3. Decomposition of water (refer Chapter 1)
4. Manufacturing of NaOH and Cl₂ (Chapter 2 — chlor-alkali)

Contrast — Covalent Compounds

Covalent compounds — generally not electrical conductors. (No ions formed.)

Exception: Acids like HCl, H₂SO₄ — get ionised in water.

Chemical Nature and Other Properties

Nature of Metallic Oxides

Oxides of metals — mostly basic. (Refer to Section 2.)

Examples:

• $Na_2O + H_2O \rightarrow 2NaOH$ (alkali)
• $CaO + H_2O \rightarrow Ca(OH)_2$ (alkali)

Exceptions — Amphoteric:

• $Al_2O_3$, $ZnO$ — react with both acids and bases.

Metallic Hydroxides

Most are basic (alkali).

Examples:

• $NaOH$ (caustic soda)
• $KOH$ (caustic potash)
• $Ca(OH)_2$ (slaked lime)
• $Mg(OH)_2$ (milk of magnesia)

Non-metallic Hydrides

Non-metal + H — covalent. Examples: H₂O, NH₃, HCl, HF

HCl and HF — become ionic in water → acidic.

A Comparison — Ionic vs Covalent

Property	Ionic	Covalent
Bond	Electron transfer	Electron sharing
Melting point	High	Low
Boiling point	High	Low
State	Solid	Solid/Liquid/Gas
Solubility in water	Soluble	Mostly insoluble
In non-polar solvent	Insoluble	Soluble
Conductivity in solid	No	No
In molten/solution	Yes	Mostly No
Examples	NaCl, MgO	H₂O, CO₂

Some Interesting Facts

1. Diamond:

• Covalent crystal — but very hard!
• Exception — not all covalents are soft.

2. Graphite:

• Covalent, but — electrical conductor!
• Because of delocalised electrons.

3. Taste of salt:

• NaCl — salty.
• KCl — bitter-salty.
• Human tongue specifically recognises Na⁺.

Practical Uses of Ionic Compounds

1. NaCl (Common Salt)

Uses:

• In food — taste, digestion.
• Pickling, food preservation.
• Raw material for NaOH, Cl₂, Na.
• To prevent dehydration in hot countries.

2. CaO (Quicklime)

Uses:

• Cement manufacturing.
• Sugar refining.
• Paper manufacturing.
• Increasing pH of acidic soil.

3. Ca(OH)₂ (Slaked Lime)

Uses:

• Whitewashing walls.
• Making cheese.
• pH adjustment of water.
• Manufacturing bleaching powder.

4. CaCO₃ (Calcium Carbonate)

Sparingly-soluble ionic compound.

Uses:

• Major component of cement.
• Marble, limestone.
• Food additive (E170).
• In classroom dusters as 'chalk'.
• Component of toothpaste.

5. Na₂CO₃ (Washing Soda)

Uses:

• Washing clothes.
• Glass manufacturing.
• Paper industry.

6. NaHCO₃ (Baking Soda)

Uses:

• Cooking.
• Fire extinguishers.
• Antacid (stomach acidity).

7. CuSO₄ (Blue Vitriol)

Uses:

• Pesticide.
• Fungicide.
• Electroplating.
• Pigment.

Summary Table — Properties and Examples

Property	Ionic	Reason
Physical state	Solid	Strong attraction
Melting point	High	Strong ionic bond
Brittleness	Brittle	Charge repulsion
In water	Soluble	Water is polar
In non-polar	Insoluble	Opposite nature
Solid conductivity	No	Ions fixed
Aqueous/Molten	Yes	Ions mobile

[Board Important] Cause of all properties = crystal lattice + strong ionic forces.

🧠 Memory Capsule

A quick glance just before the board exam.

1. 5 Main Physical Properties of Ionic Compounds

1. Solid state — hard, brittle, crystalline.
2. High melting/boiling points — strong attraction.
3. Soluble in water — water is polar.
4. Insoluble in non-polar solvents — kerosene, benzene.
5. Electrical conductivity: No in solid, Yes in molten/solution.

2. Examples of Melting Points

Compound	Melting Point
NaCl	801°C
MgO	2852°C
CaO	2572°C
Al₂O₃	~2050°C

3. Rule of Electrical Conductivity

Solid — no (ions fixed) Molten — yes (ions free) Aqueous — yes (ions free)

4. Mechanism of Dissolution

• Polar water → surrounds ions → ions leave the lattice.
• 'Like solvent dissolves like solute.'

5. Reason for Brittleness

• Slight force → ions shift → like charges face-to-face → repulsion → crystal breaks.

6. Ionic vs Covalent (Quick recall)

Property	Ionic	Covalent
Melting point	High	Low
Water	Soluble	Insoluble
Solid conductivity	No	No
Water conductivity	Yes	No

7. Board's 'Golden' Questions

1. Five properties of ionic compounds.
2. Why is solid NaCl not an electrical conductor?
3. Why is melting point of MgO higher than NaCl?
4. Why are ionic compounds brittle?
5. Why are ionic compounds soluble in water?

Final Formula: All properties = crystal lattice + strong ion-ion attraction.

Solved Examples

Example 1: NCERT — Properties of Ionic Compounds

Write five main physical properties of ionic compounds.

Solution:

5 Physical Properties of Ionic Compounds:

1. Physical State and Structure

• Ionic compounds are typically solid.
• Hard but brittle.
• Crystalline. Examples: NaCl, KCl — table salt.

2. Melting and Boiling Point

• High melting and boiling points.
• Due to strong electrostatic attraction between ions. Examples: NaCl - 801°C; MgO - 2852°C.

3. Solubility

• Soluble in water (mostly).
• Insoluble in non-polar solvents (kerosene, petrol). Reason: water is polar.

4. Electrical Conductivity

• In solid state: No (ions fixed).
• In molten/aqueous: Yes (ions mobile).

5. Nature

• Mainly metal + non-metal.
• Ions in a crystal lattice.

[NCERT textbook — asked every year]

Example 2: NCERT — Comparison of Melting Points

Which has a higher melting point — NaCl or MgO? Why?

Solution:

Comparison:

Compound	Ions	Charges
NaCl	Na⁺, Cl⁻	+1, -1
MgO	Mg²⁺, O²⁻	+2, -2

Melting Points:

• NaCl: 801°C
• MgO: 2852°C (3.5× higher than NaCl!)

Reason — Coulomb's Law

$$F = \frac{kq_1q_2}{r^2}$$

Force of attraction:

In NaCl: $F \propto (+1)(-1) = 1$ (unit)

In MgO: $F \propto (+2)(-2) = 4$ (unit)

Attraction in MgO is 4 times stronger.

More attraction → more energy needed to melt → higher melting point.

Additional Reason — Size of Ions

• Mg²⁺ and O²⁻ — both smaller ions (smaller radii).
• $r$ (distance) less → $F$ more (Coulomb's Law).
• Na⁺ and Cl⁻ — relatively larger → $F$ less.

Final Answer

MgO has a higher melting point because:

1. Charges of ions are higher (±2 vs ±1).
2. Attraction between ions is 4× stronger.
3. Ions are smaller — distance is less.

The same principle explains the very high melting point of Al₂O₃ (Al³⁺O²⁻).

[Board: 3-mark]

Example 3: NCERT — Electrical Conductivity

In an experiment, current was passed through dry NaCl, molten NaCl, and aqueous NaCl. In which did the bulb glow?

Solution:

Setup: Battery + bulb + two electrodes + test medium.

(a) Dry NaCl Solid

Result: Bulb did not glow.

Reason:

• In solid NaCl, Na⁺ and Cl⁻ ions are fixed in the crystal lattice.
• They cannot move.
• Electric current = movement of ions. No movement here → no current.

(b) Molten NaCl

Result: Bulb glowed! (very brightly)

Reason:

• Above the melting point (801°C), NaCl melts.
• Crystal lattice breaks down.
• Ions become mobile in the melt.
• Na⁺ → goes to negative electrode (cathode).
• Cl⁻ → goes to positive electrode (anode).
• This ionic movement = current.

(c) Aqueous Solution of NaCl

Result: Bulb glowed. (moderate)

Reason:

• In water, NaCl → Na⁺ + Cl⁻ (ionised).
• Aqueous ions move freely.
• Current flows.
• (Slightly less than molten — because ions are surrounded by water molecules, slight resistance.)

Summary

Medium	Conductor?	Reason
Dry NaCl	No	Ions fixed
Molten NaCl	Yes (bright)	Ions mobile
NaCl solution	Yes	Aqueous ions

Key Insight

Conductivity of ionic compounds depends on movement of ions. This is the basis of industrial 'electrolysis'.

[NCERT textbook — important experiment]

Example 4: NCERT — Solubility

Why is NaCl soluble in water but insoluble in kerosene?

Solution:

Principle — 'Like dissolves like'

Polar solute — in polar solvent. Non-polar solute — in non-polar solvent.

NaCl: Ionic → 'Highly polar'

Na⁺ and Cl⁻ — fully charged ions.

Water: Polar Molecule

• $H$ end: partial positive ($\delta^+$).
• $O$ end: partial negative ($\delta^-$).

That is, water — a 'dipole'.

Mechanism of Dissolution

Step 1: Add NaCl to water.

Step 2: Water molecules surround the ions.

• $H_2O$'s $O$ ($\delta^-$) → towards $Na^+$.
• $H_2O$'s $H$ ($\delta^+$) → towards $Cl^-$.

Step 3: Water's attraction — stronger than ion-ion attraction.

• Ions are pulled out of crystal lattice.

Step 4: Aqueous ions = $Na^+(aq)$ and $Cl^-(aq)$.

• This is 'Hydration'.

Kerosene: Non-polar

Kerosene — hydrocarbon (C-H bonds). Non-polar — no partial charges.

Reaction with kerosene:

• No attraction towards Na⁺ or Cl⁻.
• Ions remain in crystal lattice.
• NaCl insoluble.

An Interesting Demonstration

Oil and water with NaCl:

• NaCl dissolves in water.
• Not in oil.
• Finally: bottom — water + NaCl; top — oil.

Summary

NaCl + water → solution (soluble) — water is polar. NaCl + kerosene → insoluble — kerosene is non-polar.

'Like dissolves like' — fundamental rule of chemistry.

[NCERT — 3-mark]

Example 5: NCERT — Brittleness

Why are ionic compounds brittle?

Solution:

Crystal Lattice

In NaCl: each Na⁺ surrounded by 6 Cl⁻. Each Cl⁻ by 6 Na⁺. This structure — 3D cubic lattice.

Normal State

Ions are arranged so that:

• Opposite charges are close.
• Like charges are far. = electrostatic stability.

When Force is Applied

1. We apply slight force on the crystal (hammer, pressure).

2. Ions of one layer — shift relative to another layer.

3. Now like charges come face-to-face!

• Na⁺ in front of Na⁺ — repulsion.
• Cl⁻ in front of Cl⁻ — repulsion.

4. Strong repulsion → crystal breaks.

Comparison — with Metals

In metals: sea of electrons.

• Layers can shift — electrons fill the gap.
• Metals — malleable and ductile.

In ionic compounds: fixed ions.

• Shift = immediately breaks.
• Ionic — brittle.

Practical Demonstration

Hammer on a large salt crystal:

• Immediately — breaks into many small pieces.
• Pieces are also — cubic! (because they break along the crystal lattice).

Other Famous Brittle Substances

• Glass (covalent, yet brittle — crystal-like structure).
• Ceramics (Al₂O₃ based — ionic, brittle).
• Salt crystals.
• Sugar (although covalent, crystal-like).

Summary

Ionic compounds are brittle because:

1. Ions are fixed in crystal lattice.
2. Force → displacement of layers.
3. Like charges face each other → repulsion.
4. Crystal breaks.

[Board: 3-mark]

Example 6: NCERT — Two Demonstrations

(a) Connect a dry NaCl rod to a bulb and battery. What happens? (b) Place two graphite rods in NaCl solution and light the bulb — what happens?

Solution:

(a) Dry NaCl Rod

Result: Bulb did not glow.

Reason:

• In solid NaCl, ions are fixed in crystal lattice.
• No mobile charge carriers.
• Electrical conductivity = 0.

Conclusion: solid ionic compound — non-conductor.

(b) NaCl Solution

Setup:

• Beaker with NaCl + water solution.
• Two graphite rods — as electrodes.
• Connected to battery.
• Bulb in circuit.

Result: Bulb glowed! (with moderate brightness)

Mechanism:

1. NaCl ionised in water: $$NaCl \rightarrow Na^+ + Cl^-$$

2. Movement of ions:

• Na⁺ → at cathode (-)
• Cl⁻ → at anode (+)

3. Reactions at electrodes:

• Cathode: $Na^+ + e^- \rightarrow Na$ (or actually, in aqueous — $H^+$ instead of less reactive Na)
• Actually: $2H_2O + 2e^- \rightarrow H_2 + 2OH^-$
• Anode: $2Cl^- \rightarrow Cl_2 + 2e^-$

4. Observations:

• Bubbles at cathode (H₂).
• Bubbles and pale yellow gas at anode (Cl₂).
• Bulb glows.

Important Point

This is the basic method of the 'chlor-alkali process'! (Refer to Chapter 2)

In industry: $$2NaCl + 2H_2O \xrightarrow{\text{electricity}} 2NaOH + H_2 + Cl_2$$

Summary

Experiment	Conductor?
Dry NaCl rod	No
NaCl solution	Yes

[NCERT textbook experiment]

Example 7: NCERT — Identifying Ionic vs Covalent

Which of the following have ionic properties? NaCl, sugar (C₁₂H₂₂O₁₁), MgO, glucose, KBr, $CO_2$

Solution:

Analysis

1. NaCl:

• Elements: Na (metal) + Cl (non-metal)
• ✓ Ionic.

2. Sugar (C₁₂H₂₂O₁₁):

• Elements: C, H, O — all non-metals.
• ✗ Covalent.

3. MgO:

• Elements: Mg (metal) + O (non-metal).
• ✓ Ionic.

4. Glucose (C₆H₁₂O₆):

• Elements: C, H, O.
• ✗ Covalent.

5. KBr:

• Elements: K (metal) + Br (non-metal).
• ✓ Ionic.

6. CO₂:

• Elements: C + O (both non-metals).
• ✗ Covalent.

Confirmation by Properties

Confirming three ionic ones:

Property	NaCl	MgO	KBr
High melting point?	801°C ✓	2852°C ✓	734°C ✓
Soluble in water?	✓	✗ Sparingly	✓
Conducts in water?	✓	Slightly	✓

Confirming three covalent ones:

Property	Sugar	Glucose	CO₂
Melting point	186°C (low)	146°C	-78°C
Conducts in water?	✗	✗	✗

Final List

Ionic: NaCl, MgO, KBr Covalent: sugar, glucose, CO₂

A Caveat

Not all ionic compounds have all properties identical. Example: AgCl — ionic, but insoluble in water. MgO — ionic, but sparingly soluble in water. These exceptions due to specific reasons.

[Board: 5-mark]

Example 8: An Interesting — Ranking Melting Points

Arrange the following ionic compounds in increasing order of melting point — with reasons: NaCl, MgO, NaF, MgCl₂

Solution:

Analysis of Each

1. NaCl:

• Charges: +1, -1
• Attraction $\propto 1$

2. MgO:

• Charges: +2, -2
• Attraction $\propto 4$ (very high!)

3. NaF:

• Charges: +1, -1
• Attraction $\propto 1$
• But F⁻ is small, smaller than Cl⁻ → distance less → attraction slightly higher.

4. MgCl₂:

• Charges: +2, -1 (less 'pure' ionic)
• Attraction $\propto 2$

Melting Point Values

Compound	Melting Point
MgO	2852°C
NaF	993°C
NaCl	801°C
MgCl₂	714°C

Increasing Order

$$\boxed{MgCl_2 < NaCl < NaF < MgO}$$

Reasoning

MgO is highest:

• ±2 charges, small ions — strongest attraction.

NaF higher than NaCl:

• Both ±1, but F⁻ is smaller.
• Slightly more attraction.

MgCl₂ is lowest:

• Special: Mg²⁺ is very strong — slightly 'pulls' on the electron.
• Result — partial covalent character.
• (Fajans' Rules.)
• Slightly lower melting point.

Key Insight

Melting points depend on:

1. Charges of ions (more → higher).
2. Size of ions (smaller → higher).
3. Partial covalent character (more → lower).

[Board: 5-mark reasoning]

Example 9: NCERT — Choice of Solvent

Choose the best solvent — with reason:

(a) For NaCl — water or kerosene? (b) For naphthalene (non-polar) — water or benzene? (c) For sodium chloride — ethanol (mildly polar) or water (highly polar)?

Solution:

Principle

'Like solvent dissolves like solute.'

(a) For NaCl

Options:

• Water: polar.
• Kerosene: non-polar.

NaCl is ionic — highly polar. Polar solvent needed.

Answer: Water.

Evidence: NaCl is soluble in water (35-40 g per 100 g water). Insoluble in kerosene.

(b) For Naphthalene (C₁₀H₈)

Options:

• Water: polar.
• Benzene: non-polar.

Naphthalene — non-polar (hydrocarbon). Non-polar solvent needed.

Answer: Benzene.

Evidence: Naphthalene (camphor / 'moth-balls') is insoluble in water. Highly soluble in benzene.

(c) For NaCl — ethanol or water?

Options:

• Ethanol (C₂H₅OH): mildly polar (OH group).
• Water: highly polar.

NaCl needs the more polar solvent.

Answer: Water.

Evidence: NaCl highly soluble in water (~360 g/L). Very low in ethanol (~0.65 g/L).

Summary Table

Solute	Nature	Good Solvent
NaCl	Ionic	Water
MgSO₄	Ionic	Water
Naphthalene	Non-polar	Benzene
Iodine	Non-polar	CCl₄
Sugar	Polar (OH)	Water

Final Insight

This rule — not just chemistry — daily life.

'Oil and water don't mix' — proof of this rule!

[NCERT textbook]

Example 10: An Interesting — Role of Electrolysis

What will be obtained on electrolysis of an aqueous NaCl solution?

Solution:

Setup

• Aqueous NaCl solution (brine).
• Two graphite electrodes.
• DC power source.

Mechanism

NaCl ionised in water: $$NaCl \rightarrow Na^+ + Cl^-$$

Water also slightly ionised: $$H_2O \rightleftharpoons H^+ + OH^-$$

Reactions at Electrodes

At Cathode (-): Less reactive ion is liberated.

Here — Na⁺ vs H⁺ (from water):

• Na — highly reactive (no H₂).
• H₂ — easily formed.

$$2H_2O + 2e^- \rightarrow H_2\uparrow + 2OH^-$$

At Anode (+): Less reactive ion is oxidised.

Here — Cl⁻ vs OH⁻:

• Cl⁻ → Cl₂ (more likely in concentrated solution).

$$2Cl^- \rightarrow Cl_2\uparrow + 2e^-$$

Overall Reaction

$$2NaCl + 2H_2O \xrightarrow{\text{electricity}} 2NaOH + H_2\uparrow + Cl_2\uparrow$$

Observations

• At Cathode: Colourless gas (H₂) — pop test.
• At Anode: Pale yellow gas (Cl₂) — suffocating smell.
• In solution: NaOH (base).

Industrial Use

This is the 'chlor-alkali process':

• NaOH (caustic soda): soap, paper, textile.
• Cl₂: water purification, bleaching powder.
• H₂: fuel, ammonia.

Some Interesting Facts

'Chlor-alkali':

• Chlor = chlorine (Cl₂)
• Alkali = base (NaOH)
• One process, three products!

Volume calculation: If 1 mol NaCl is taken (NTP):

• 0.5 mol H₂ = 11.2 L
• 0.5 mol Cl₂ = 11.2 L
• 1 mol NaOH (in solution).

[Board + Industrial]

Example 11: Numerical — Electrical Conductivity

0.1 mol NaCl was dissolved in 1 L of water. How many free ions are present?

Solution:

Ionisation:

$$NaCl \rightarrow Na^+ + Cl^-$$

From 0.1 mol NaCl:

• 0.1 mol Na⁺
• 0.1 mol Cl⁻
• Total = 0.2 mol ions

Number of ions:

1 mol = $6.022 \times 10^{23}$.

• Na⁺ = $0.1 \times 6.022 \times 10^{23} = 6.022 \times 10^{22}$
• Cl⁻ = $6.022 \times 10^{22}$
• Total = $1.2 \times 10^{23}$ ions

Concentration:

$1.2 \times 10^{23}$ ions in 1 L solution.

$\frac{1.2 \times 10^{23}}{1000}$ = $1.2 \times 10^{20}$ ions per mL.

Contribution to Electrical Conduction:

Na⁺ → at cathode. Cl⁻ → at anode.

Both together — generate electric current.

Comparison — Effect of Concentration

1 mol/L NaCl: more conductivity. 0.1 mol/L NaCl: less conductivity. 0.001 mol/L NaCl: very low conductivity.

That is — concentration ↑ → ions ↑ → conductivity ↑.

Final Numbers

0.1 mol NaCl in 1 L solution:

• Total ions: $1.2 \times 10^{23}$
• Na⁺ ions: $6.022 \times 10^{22}$
• Cl⁻ ions: $6.022 \times 10^{22}$

[Board: 3-mark numerical]

Example 12: NCERT — Daily Life Compounds

Mention daily uses of the following ionic compounds:

(a) NaCl (b) CaO (c) NaHCO₃ (d) Na₂CO₃

Solution:

(a) NaCl — Sodium Chloride (Common Salt)

Daily Uses:

1. In food: taste, digestion, electrolyte.
2. Food preservation: fish, pickles, meat.
3. Industrial: for manufacturing NaOH, Cl₂, Na₂CO₃, Na.
4. Medical: saline — IV fluid.
5. In winter: for melting ice.

In India: most obtained by evaporation of seawater.

(b) CaO — Calcium Oxide (Quicklime)

Daily Uses:

1. Cement manufacturing: main ingredient.
2. Sugar refining: removing impurities.
3. Water pH: making acidic water alkaline.
4. Lime: for whitewashing (as Ca(OH)₂).
5. Paper manufacturing.
6. Soil fertility.

(c) NaHCO₃ — Sodium Bicarbonate (Baking Soda)

Daily Uses:

1. In cooking: for cakes, biscuits — to make them rise.
2. Antacid: stomach acidity.
3. Fire extinguisher: $$2NaHCO_3 \xrightarrow{\Delta} Na_2CO_3 + H_2O + CO_2$$ CO₂ extinguishes fire.
4. Toothpaste.
5. Cleaning fabrics.

(d) Na₂CO₃ — Sodium Carbonate (Washing Soda)

Daily Uses:

1. Washing: with detergents for clothes.
2. Glass manufacturing: main raw material.
3. Paper industry.
4. Removing water hardness.
5. Manufacturing borax.

'Washing soda' — Na₂CO₃·10H₂O.

Summary

These four compounds — everywhere from kitchen to industry!

[NCERT — chemistry in daily life]

Example 13: NCERT — A Comparative Question

NaCl and sugar (sucrose) — both white solids, both soluble in water. But what is the difference in physical properties?

Solution:

Fundamental Difference — Bond Type

NaCl: Ionic (Na — metal, Cl — non-metal). Sugar (C₁₂H₂₂O₁₁): Covalent (C, H, O — all non-metals).

Detailed Comparison

Property	NaCl	Sugar
Bond	Ionic	Covalent
Melting point	801°C	186°C
Boiling point	1465°C	Decomposes
Solubility in water	36 g/100mL	200 g/100mL
Conducts in water?	Yes	No!
Conducts in molten state?	Yes	No
Crystal	Cubic	Tricyclic
Taste	Salty	Sweet
Brittleness	Brittle	Brittle (less so)

Special — Electrical Conductivity

NaCl + water → Na⁺ + Cl⁻ → conducts electricity.

Sugar + water → C₁₂H₂₂O₁₁ molecules (intact) → no ions → no conduction.

A Demonstration

Two beakers:

• One: NaCl + water
• Other: sugar + water

Connect to bulb-battery:

• NaCl: bulb glows.
• Sugar: bulb does not glow.

A Confusion — Sugar 'dissolves' in water, yet doesn't conduct?

Sugar dissolves in water as molecules — not as ions. Ions = required for conduction. Molecules — not for conduction.

Key Insight

Ionic compound: gives ions → conducts. Covalent compound: molecules → does not conduct. Exception: Acids like HCl — covalent, but ionised in water → conducts.

[Board: 5-mark comparative]

Example 14: An Interesting — KMnO₄ vs MgO

Both are ionic. Yet KMnO₄ gives violet colour in water, MgO is sparingly soluble. Why?

Solution:

KMnO₄

Ions: $K^+$ (+1) and $MnO_4^-$ (-1)

$MnO_4^-$ — polyatomic ion.

Solubility:

• K⁺ small but charge +1.
• MnO₄⁻ large.
• Moderate attraction.
• Water easily separates them.

That is — highly soluble in water.

MnO₄⁻ gives violet colour.

MgO

Ions: $Mg^{2+}$ and $O^{2-}$

Solubility:

• Both ±2 charge — strong attraction.
• Both small — distance less.
• Attraction $\propto 4$.
• Water's attraction — less than Mg²⁺-O²⁻ attraction.
• Ions remain in crystal lattice.

That is — sparingly soluble in water.

Principle

'Solubility' depends on:

1. Lattice Energy ($U_{lattice}$): ion-ion attraction.
2. Hydration Energy ($U_{hydration}$): ion-water attraction.

If $U_{hydration} > U_{lattice}$ → soluble. If $U_{lattice} > U_{hydration}$ → insoluble.

Comparison

Compound	$U_{lattice}$	$U_{hydration}$	Solubility
KMnO₄	Low	High	Soluble (violet)
NaCl	Moderate	High	Soluble
MgO	Very high	Moderate	Insoluble
AgCl	High	Low	Insoluble

Key Insight

'All ionic — soluble in water' — not entirely true. Depends on this balance.

More charge + smaller ions → higher lattice energy → less soluble.

[Board + Practical]

Example 15: A Challenging Question

Is it true that — "All ionic compounds are soluble in water, all are electrical conductors"? Explain with reasoning.

Solution:

Original Statement

"All ionic compounds are soluble in water, all are electrical conductors."

Answer: This is not entirely true.

Analysis

'Soluble in water' — exceptions exist:

Compound	Solubility	Reason
NaCl, KCl	Highly soluble	Balanced charges
AgCl	Insoluble	High lattice energy
BaSO₄	Insoluble	Very stable lattice
PbI₂	Insoluble	Large ions, strong attraction
CaCO₃	Insoluble	Strong forces
CaSO₄	Sparingly soluble	Moderate
MgO	Sparingly soluble	±2 charges

Therefore: About 70-80% of ionic compounds are soluble, 20-30% are not.

Electrical Conductivity — Correct?

Two situations:

1. In solid state:

• Ions fixed — no conductivity.
• All solid ionic compounds — non-conductors.

2. In molten or aqueous:

• Ions free — conductors.
• All ionic — conductors in molten/aqueous.

Exception:

• If insoluble (AgCl), then aqueous not conductor (ions don't come out).
• In molten state — still a conductor.

Updated Statements

Correct statements:

'Most ionic compounds are soluble in water.' 'All ionic compounds — conductors in molten state.' 'Soluble ionic compounds — conductors in aqueous solution.' 'Solid ionic compounds — non-conductors of electricity.'

Final Reasoning

The original statement is a 'generalisation'. Several exceptions exist. The actual test depends on properties.

[Board: 5-mark]

Example 16: A Concluding Question

(a) 5 physical properties of ionic compounds. (b) Why is solid NaCl not a conductor, but molten is? (c) Why is the melting point of MgO higher than NaCl — explain. (d) Ionic vs Covalent — most important difference.

Solution:

(a) 5 Physical Properties

1. Solid state: hard, crystalline, brittle.
2. High melting/boiling points: strong ion-ion attraction.
3. Soluble in water: polar water surrounds ions.
4. Insoluble in non-polar: kerosene, benzene.
5. Electrical conductivity: No in solid, Yes in molten/aqueous.

(b) Conductivity of NaCl

Solid NaCl:

• Ions fixed in crystal lattice.
• No motion → no current.
• ✗ Not a conductor.

Molten NaCl:

• Above 801°C — melts.
• Crystal lattice breaks.
• Ions free — mobile.
• Na⁺ → cathode, Cl⁻ → anode.
• ✓ Highly conducting.

This is the basic principle of 'electrolysis'.

(c) Melting Point of MgO — Higher than NaCl

Comparison:

• NaCl: ±1 charges, melting point 801°C
• MgO: ±2 charges, melting point 2852°C

Reason (Coulomb): $F \propto q_1q_2$

Attraction in MgO = 4 × NaCl.

More attraction → more energy needed to melt → higher melting point.

+ Smaller ion size → less distance → more attraction.

(d) Ionic vs Covalent — Most Important Difference

Most fundamental difference:

Ionic: complete transfer of electrons → ions. Covalent: sharing of electrons → molecules.

Practical evidence: Ionic — conducts in aqueous solution. Covalent — does not conduct (exception: acids).

Quick Table

Property	Ionic	Covalent
Melting point	High	Low
Conducts in water	Yes	No (generally)
Bond	Transfer	Sharing
Examples	NaCl, MgO	H₂O, CO₂

[Board: 5-mark mixed question]