Bonding in Carbon — The Covalent Bond

Carbon — A Special Element

Carbon is everywhere — in our food, fuel, clothes, body, and even in the breath we exhale.

Look around — most things you see contain carbon!

How Special is Carbon?

Earth's crust has only 0.02% carbon (as carbonates, hydrocarbons, and CO₂). Atmosphere has only 0.03% carbon (as CO₂).

Yet — all known life is carbon-based!

Why? Because carbon forms an enormous variety of compounds — millions of them.

Most other elements form only a few hundred compounds. Carbon alone — more than 10 million!

Position of Carbon in the Periodic Table

Carbon (C):

• Atomic Number: 6
• Electronic Configuration: 2, 4
• Group: 14 (IVA)
• Period: 2

Outer shell has 4 electrons.

Why Doesn't Carbon Form Ionic Bonds?

To attain a noble gas configuration (octet), carbon would need to either:

Option 1: Lose 4 electrons → become C⁴⁺ (like helium 'He', 2 electrons). Problem: needs enormous energy to remove 4 electrons.

Option 2: Gain 4 electrons → become C⁴⁻ (like neon 'Ne', 10 electrons). Problem: nucleus with only 6 protons cannot hold 10 electrons stably.

Both options — energetically unfavourable.

The Solution — Covalent Bond

Carbon shares its 4 electrons with other atoms. This way, both atoms attain a stable octet without losing or gaining electrons.

This shared electron-pair bond = Covalent Bond.

Why is Carbon Unique?

Carbon's 4 valence electrons mean it can:

• Share with 4 other atoms (single bonds).
• Share 2 pairs with one atom (double bond).
• Share 3 pairs with one atom (triple bond).

This versatility is unique among elements.

What is a Covalent Bond?

Definition: A chemical bond formed by the sharing of electrons between two atoms — so that both can attain a stable noble gas configuration.

Simple Examples

1. Hydrogen Molecule ($H_2$)

Two H atoms — each has 1 electron, needs 1 more for stability (like helium).

Each shares its 1 electron — both now have 2 electrons (helium configuration).

Lewis dot structure: H : H or H — H

One shared pair = single bond (one line).

2. Oxygen Molecule ($O_2$)

Two O atoms — each has 6 valence electrons, needs 2 more for octet.

Each shares 2 electrons — total 4 shared (2 pairs).

Lewis dot structure: O :: O or O = O

Two shared pairs = double bond (two lines).

3. Nitrogen Molecule ($N_2$)

Two N atoms — each has 5 valence electrons, needs 3 more for octet.

Each shares 3 electrons — total 6 shared (3 pairs).

Lewis dot structure: N ::: N or N ≡ N

Three shared pairs = triple bond (three lines).

Carbon Compounds — Examples

Methane ($CH_4$)

Carbon (4 valence electrons) + 4 Hydrogens (1 each).

Each H shares 1 electron with C. C ends up with 8 electrons (octet). Each H ends up with 2 electrons (helium config).

Structure:

       H
       |
   H — C — H
       |
       H

All 4 bonds = single bonds.

Ethene ($C_2H_4$)

Two C atoms + 4 H atoms. C-C bond is a double bond (2 shared pairs).

Structure:

    H       H
     \     /
      C = C
     /     \
    H       H

Ethyne / Acetylene ($C_2H_2$)

Two C atoms + 2 H atoms. C-C bond is a triple bond (3 shared pairs).

Structure: H — C ≡ C — H

Bond Strength Order

Triple bond > Double bond > Single bond. (More shared pairs = stronger bond.)

But also: triple > double > single — shorter bond length.

Common Covalent Compounds

Water ($H_2O$)

O has 6 valence electrons; needs 2 more for octet. Each H shares 1 electron with O.

Structure: H — O — H

Two single bonds.

Ammonia ($NH_3$)

N has 5 valence electrons; needs 3 more for octet. Three H atoms each share 1 electron.

Structure:

    H — N — H
        |
        H

Three single bonds.

Carbon Dioxide ($CO_2$)

C shares 2 pairs with each O atom (double bonds with both).

Structure: O = C = O

Two double bonds.

Hydrogen Chloride ($HCl$)

H shares 1 electron with Cl. Cl already has 7 electrons; gains 1 more for octet.

Structure: H — Cl

Single bond.

General Rule for Covalent Compounds

Covalent compounds form between non-metals.

Examples:

• C—H, C—C, C—O, C—N, C—Cl bonds.
• H—H, O—O, N—N, F—F bonds.
• H₂O, CO₂, NH₃, HCl, etc.

Octet Rule — In Practice

Octet Rule: Atoms tend to have 8 electrons in their outermost shell (or 2 for elements close to helium).

Achieved by:

• Ionic bond (transfer of electrons) — for metal + non-metal.
• Covalent bond (sharing of electrons) — for non-metal + non-metal.

Properties of Covalent Compounds

Covalent compounds have very different properties from ionic compounds.

1. Physical State

Covalent compounds can exist in all three states:

• Gas: O₂, N₂, CH₄, CO₂, NH₃, HCl
• Liquid: H₂O, ethanol (C₂H₅OH), petrol
• Solid: sugar (C₁₂H₂₂O₁₁), naphthalene

2. Melting and Boiling Points

Generally low (compared to ionic compounds).

Examples:

• $CH_4$ melting point: -182°C
• $H_2O$ boiling point: 100°C
• $NH_3$ boiling point: -33°C
• $C_2H_5OH$ boiling point: 78°C

Reason: Weak intermolecular forces (van der Waals).

Compared to ionic NaCl (801°C) — covalent compounds are much softer.

3. Electrical Conductivity

Covalent compounds are poor conductors of electricity.

Why?

• No free ions.
• No free electrons.
• Bonds are localised between specific atoms.

Exceptions:

• Acids (HCl, H₂SO₄) — ionise in water → conduct.
• Graphite — has free electrons → conducts.

4. Solubility

Generally:

• Soluble in non-polar solvents (kerosene, benzene).
• Insoluble in water (mostly).

Exceptions:

• Polar covalent (HCl, NH₃, sugar) — soluble in water.
• Ionic-like behaviour (acids) — soluble in water.

5. Bond Strength

Very strong within a molecule (intramolecular). But weak between molecules (intermolecular).

This is why:

• Covalent compounds are difficult to break chemically.
• But easy to melt or boil physically.

Comparison Table — Ionic vs Covalent

Property	Ionic	Covalent
Bond formation	Electron transfer	Electron sharing
Elements	Metal + Non-metal	Non-metal + Non-metal
Physical state	Solid (mostly)	Solid/Liquid/Gas
Melting point	High	Low
Boiling point	High	Low
Solubility in water	Soluble	Mostly insoluble
Solubility in kerosene	Insoluble	Soluble
Conductivity (in water)	Yes	No (mostly)
Examples	NaCl, MgO	H₂O, CO₂, CH₄

Carbon's Special Tetravalency

Carbon's electronic configuration: 2, 4.

Carbon needs 4 more electrons to complete its octet.

It does this by sharing 4 electrons with other atoms.

This means carbon forms exactly 4 covalent bonds.

Tetravalency in Action

Carbon makes 4 bonds — these can be:

1. Four single bonds: Methane (CH₄): 4 C-H single bonds. Ethane (C₂H₆): 1 C-C single + 6 C-H single bonds.

2. Two single + one double bond: Ethene (C₂H₄): 1 C=C double + 4 C-H single bonds.

3. One single + one triple bond: Ethyne (C₂H₂): 1 C≡C triple + 2 C-H single bonds.

4. Two double bonds: Carbon dioxide (CO₂): 2 C=O double bonds.

This flexibility is what makes carbon so versatile.

Bond Lengths and Strengths

Bond	Length (pm)	Strength (kJ/mol)
C — C (single)	154	348
C = C (double)	134	614
C ≡ C (triple)	120	839
C — H	109	413
C — O	143	358
C = O	122	745

More bonds → shorter and stronger.

Why Are Carbon Compounds Special?

Combining tetravalency + ability to bond with itself + small size:

• C can form long chains.
• C can form rings.
• C can have branches.
• C can have multiple bonds.
• C can bond with H, O, N, S, halogens.

Result: millions of compounds.

This is what makes life — carbon-based.

A Glimpse of What's Coming

In the next sections, we'll see:

• Allotropes of carbon (diamond, graphite, fullerene).
• Hydrocarbons (CH compounds).
• Functional groups.
• Important compounds — ethanol, ethanoic acid.
• Soaps and detergents.

All — possible because of carbon's covalent bonding ability.

🧠 Memory Capsule

A quick glance just before the board exam.

1. Carbon's Configuration

C: atomic number 6, configuration 2, 4. Valency = 4 (tetravalency).

2. Why Covalent and Not Ionic?

Losing 4 electrons → too much energy. Gaining 4 electrons → unstable nucleus. Solution: share 4 electrons → covalent bonds.

3. Covalent Bond Definition

Bond formed by sharing of electrons between two atoms — so that both attain stable octet.

4. Types of Covalent Bonds

Type	Shared pairs	Example
Single	1	H-H, C-C, H-Cl
Double	2	O=O, C=C, C=O
Triple	3	N≡N, C≡C

5. Bond Length / Strength

Triple > Double > Single (strength) Single > Double > Triple (length)

6. Properties of Covalent Compounds

1. Low melting/boiling points.
2. Insoluble in water (mostly).
3. Soluble in organic solvents (kerosene).
4. Poor conductors of electricity.
5. All physical states (gas, liquid, solid).

7. Common Covalent Examples

• H₂, O₂, N₂ (diatomic)
• H₂O, NH₃, CO₂ (small molecules)
• CH₄, C₂H₆, C₂H₄, C₂H₂ (hydrocarbons)

8. Ionic vs Covalent — Quick Recall

Feature	Ionic	Covalent
Bond	Transfer	Sharing
Melting point	High	Low
Water solubility	Yes	No
Conductivity	Yes (in water)	No

9. Carbon's 4 Bonds — Possible Combinations

• 4 single bonds (CH₄, C₂H₆)
• 2 single + 1 double (CH₂=CH₂)
• 1 single + 1 triple (HC≡CH)
• 2 double bonds (CO₂)

10. Board's 'Golden' Questions

1. Why does carbon form covalent bonds?
2. Properties of covalent compounds.
3. Lewis dot structure of $H_2$, $O_2$, $N_2$.
4. Tetravalency of carbon.
5. Difference between ionic and covalent.

Final Formula: Carbon shares 4 electrons → covalent bonds → millions of compounds → life itself.

Solved Examples

Example 1: NCERT — Why Carbon Forms Covalent Bonds

Why does carbon form compounds mainly by covalent bonding?

Solution:

Carbon's electronic configuration: 2, 4

Carbon has 4 valence electrons. To achieve stable octet, it could:

Option 1: Lose 4 electrons → C⁴⁺

• Requires very high energy (4 successive ionisations).
• Too unstable.

Option 2: Gain 4 electrons → C⁴⁻

• Carbon nucleus has only 6 protons.
• Cannot hold 10 electrons stably.
• Highly unstable ion.

Option 3: Share electrons (Covalent bonding)

• Each C atom shares 4 electrons with other atoms.
• Both atoms achieve octet.
• No need for energetically expensive transfer.
• Stable.

Hence, carbon prefers covalent bonding.

Examples:

• $CH_4$: C shares 4 electrons with 4 H atoms.
• $CO_2$: C shares 4 electrons with 2 O atoms (double bonds).
• $C_2H_6$: 2 C atoms share with each other and with H.

[NCERT — fundamental question]

Example 2: NCERT — Lewis Dot Structures

Draw the Lewis dot structures (electron-dot structures) for: (a) $H_2$ (b) $O_2$ (c) $N_2$ (d) $CH_4$ (e) $H_2O$ (f) $CO_2$

Solution:

(a) H₂

Each H has 1 electron. Share to form 1 pair.

H : H or H — H (single bond)

(b) O₂

Each O has 6 valence electrons. Each needs 2 more for octet. Share 2 pairs.

: O :: O : or O = O (double bond)

(c) N₂

Each N has 5 valence electrons. Each needs 3 more. Share 3 pairs.

: N ::: N : or N ≡ N (triple bond)

(d) CH₄

C has 4 valence electrons. 4 H atoms each share 1 electron with C.

Structure:

      H
      |
  H : C : H    or    H — C — H
      |                  |
      H                  H

4 single bonds.

(e) H₂O

O has 6 valence electrons. 2 H atoms share 1 each. O has 2 lone pairs (4 unshared electrons).

Structure: H : O : H or H — O — H

(f) CO₂

C shares 2 pairs with each O.

Structure: O :: C :: O or O = C = O

2 double bonds.

[NCERT textbook — every year]

Example 3: NCERT — Properties of Covalent Compounds

List the main physical properties of covalent compounds with examples.

Solution:

1. Physical State

Covalent compounds exist in all three states:

• Gas: H₂, O₂, CO₂, NH₃, CH₄
• Liquid: H₂O, C₂H₅OH (ethanol), petrol
• Solid: sugar, naphthalene, urea

2. Low Melting and Boiling Points

Examples:

• CH₄: -182°C melting
• H₂O: 0°C melting, 100°C boiling
• C₂H₅OH: -114°C melting, 78°C boiling

Reason: weak intermolecular forces (van der Waals).

3. Poor Electrical Conductivity

Generally non-conductors.

Reason: no free ions or electrons.

Exceptions:

• HCl in water (ionises)
• Graphite (delocalised electrons)

4. Solubility

Soluble in non-polar solvents (kerosene, benzene). Insoluble in water (mostly).

Exceptions:

• HCl, NH₃, sugar — soluble in water (polar).
• HCl in water → H⁺ + Cl⁻ (ionic-like behaviour).

5. Bond Strength

Bonds within molecule — strong. Forces between molecules — weak.

Examples Compared

Compound	State	M.P. (°C)	Water sol.
H₂O	Liquid	0	—
CH₄	Gas	-182	No
Sugar	Solid	186	Yes
HCl	Gas	-114	Yes (ionises)
Naphthalene	Solid	80	No

[NCERT — every year]

Example 4: NCERT — Ionic vs Covalent

Distinguish between ionic and covalent compounds with examples.

Solution:

Comparison Table

Property	Ionic	Covalent
Bond formation	Electron transfer	Electron sharing
Elements	Metal + Non-metal	Non-metal + Non-metal
State	Mostly solid	Solid/Liquid/Gas
Melting point	High (>800°C)	Low (<200°C usually)
Boiling point	High	Low
Solubility in water	Soluble	Mostly insoluble
Solubility in kerosene	Insoluble	Soluble
Conductivity (solid)	No	No
Conductivity (in water)	Yes	No (mostly)
Brittleness	Brittle	Variable

Examples

Ionic:

• NaCl (M.P. 801°C, conducts in water)
• MgO (M.P. 2852°C)
• CaCl₂ (M.P. 772°C)

Covalent:

• H₂O (M.P. 0°C)
• CH₄ (M.P. -182°C)
• CO₂ (M.P. -78°C, sublimes)
• C₆H₁₂O₆ (glucose, M.P. 146°C)

Why the Difference?

Ionic:

• Strong electrostatic forces (cation-anion).
• Crystal lattice — must break entirely.
• Hence high M.P., B.P.

Covalent:

• Strong intramolecular bonds.
• But weak intermolecular forces.
• Hence low M.P., B.P.

Key Insight: Type of bond determines properties.

[NCERT — important]

Example 5: NCERT — Tetravalency of Carbon

Explain the concept of tetravalency of carbon. Give examples.

Solution:

Definition

'Tetravalency' = property of carbon to form 4 covalent bonds with other atoms.

From Latin: 'tetra' = 4, 'valens' = strength/valency.

Origin of Tetravalency

Carbon: atomic number 6, configuration 2, 4.

4 valence electrons → needs 4 more for octet. Shares 4 electrons → forms 4 bonds.

Examples of Tetravalency

1. Methane ($CH_4$) — 4 single bonds

4 H atoms, each forming a single bond with C.

        H
        |
    H — C — H
        |
        H

4 C-H single bonds.

2. Ethane ($C_2H_6$) — 7 single bonds

2 C atoms + 6 H atoms. 1 C-C single bond + 6 C-H single bonds.

3. Ethene ($C_2H_4$) — 1 double + 4 single

Each C uses 4 bonds:

• 1 C=C double bond (counts as 2)
• 2 C-H single bonds (each C connects 2 H atoms)

4. Ethyne ($C_2H_2$) — 1 triple + 2 single

Each C uses 4 bonds:

• 1 C≡C triple bond (counts as 3)
• 1 C-H single bond

5. Carbon Dioxide ($CO_2$) — 2 double bonds

1 C atom + 2 O atoms. 2 C=O double bonds (each counts as 2 bonds). Total: 4 bonds.

Verification — Counting Bonds

Each carbon must have exactly 4 bonds:

Compound	C-C	C-H	C-O	Total
CH₄	0	4	0	4 ✓
C₂H₆	1	3 (per C)	0	4 (per C) ✓
C₂H₄	2 (double)	2 (per C)	0	4 ✓
C₂H₂	3 (triple)	1 (per C)	0	4 ✓
CO₂	0	0	4 (2 double)	4 ✓

Importance of Tetravalency

This unique tetravalency allows carbon to:

• Form long chains.
• Form branched chains.
• Form rings.
• Form double, triple bonds.
• Connect to many different atoms.

Result: millions of carbon compounds — including all of life!

[NCERT — important]

Example 6: An Interesting — Single, Double, Triple Bonds

What are single, double, and triple bonds? Give examples.

Solution:

Single Bond

One pair of electrons shared between two atoms. Represented by one line (—).

Examples:

• H — H ($H_2$)
• H — Cl ($HCl$)
• C — C (in ethane)
• C — H (in methane)

Bond strength: ~348 kJ/mol (C-C). Bond length: longer (154 pm for C-C).

Double Bond

Two pairs of electrons shared. Represented by two lines (=).

Examples:

• O = O ($O_2$)
• C = O (in CO₂)
• C = C (in ethene)

Bond strength: ~614 kJ/mol (C=C). Bond length: shorter (134 pm for C=C).

Triple Bond

Three pairs of electrons shared. Represented by three lines (≡).

Examples:

• N ≡ N ($N_2$)
• C ≡ C (in ethyne / acetylene)
• C ≡ N (in cyanide)

Bond strength: ~839 kJ/mol (C≡C). Bond length: shortest (120 pm for C≡C).

Comparison Table

Bond	Pairs	Length	Strength	Example
Single	1	154 pm	348 kJ/mol	C-C
Double	2	134 pm	614 kJ/mol	C=C
Triple	3	120 pm	839 kJ/mol	C≡C

Why Multiple Bonds?

To achieve octet:

In O₂: each O has 6 valence electrons. Sharing 2 pairs gives each O an octet (8 electrons).

In N₂: each N has 5 valence electrons. Sharing 3 pairs gives each N an octet.

In ethene: each C has 4 valence electrons. Sharing 2 pairs between C-C and sharing 1 pair each with 2 H atoms — gives each C an octet.

Practical Insight

Multiple bonds (especially double bonds) make compounds more reactive.

Example: ethene ($C_2H_4$) is more reactive than ethane ($C_2H_6$).

The double/triple bonds can break to add other atoms (addition reactions).

[Board: 5-mark]

Example 7: NCERT — Why Are Covalent Compounds Poor Conductors?

Explain why covalent compounds are generally poor conductors of electricity.

Solution:

Basic Principle

Electrical conduction requires free charge carriers (ions or electrons).

Covalent Compounds

In covalent bonds:

• Electrons are shared between specific atoms.
• They are localised in the bonds.
• No free electrons.
• No free ions (no transfer of electrons).

Hence: no charge carriers → no conduction.

Examples of Non-Conductors

• $CH_4$ (methane) — gas, no conduction.
• $H_2O$ (pure water) — almost no conduction.
• Sugar in water — molecules don't ionise → no conduction.
• Petrol, kerosene — non-conductors.

Important Exceptions

1. Acids (like HCl, H₂SO₄)

Although covalent in pure state, in water they ionise:

$$HCl + H_2O \rightarrow H_3O^+ + Cl^-$$

Now free ions present → conducts electricity.

This is why aqueous acids are conductors.

2. Graphite

Although carbon (covalent), graphite has a special structure:

• Layers of C atoms.
• Each C bonded to 3 others.
• 4th electron is delocalised (free).
• These free electrons → conduction.

Hence graphite is a unique conductor among covalent solids.

Graphite is used in:

• Electrodes.
• Pencils.
• Lubricants.

Comparison

Substance	Conducts electricity?	Reason
NaCl (solid)	No	Ions fixed
NaCl (in water)	Yes	Ions free
Sugar (in water)	No	Molecules, not ions
HCl (in water)	Yes	Ionises
Diamond	No	All bonds localised
Graphite	Yes	Free electrons
H₂O (pure)	No	Almost no ions

Summary

Covalent compounds:

• Pure: poor conductors.
• In water: still poor (unless they ionise).
• Acids and graphite are notable exceptions.

[NCERT — 3-mark]

Example 8: NCERT — Identifying Covalent Bonds

Which of the following compounds are covalent? $NaCl$, $H_2O$, $MgO$, $CO_2$, $CH_4$, $K_2S$, $CCl_4$, $NH_3$, $CaCl_2$

Solution:

Rule

Ionic = Metal + Non-metal. Covalent = Non-metal + Non-metal.

Analysis

Compound	Elements	Type
NaCl	Na (metal) + Cl (non-metal)	Ionic
H₂O	H + O (both non-metal)	Covalent
MgO	Mg (metal) + O (non-metal)	Ionic
CO₂	C + O (both non-metal)	Covalent
CH₄	C + H (both non-metal)	Covalent
K₂S	K (metal) + S (non-metal)	Ionic
CCl₄	C + Cl (both non-metal)	Covalent
NH₃	N + H (both non-metal)	Covalent
CaCl₂	Ca (metal) + Cl (non-metal)	Ionic

Covalent Compounds (5)

✓ H₂O, CO₂, CH₄, CCl₄, NH₃

Ionic Compounds (4)

✓ NaCl, MgO, K₂S, CaCl₂

Verification by Properties

Covalent compounds:

• Lower melting points.
• Don't dissolve in water (mostly).
• Don't conduct electricity.

Ionic compounds:

• High melting points.
• Soluble in water.
• Conduct in water/molten.

Practical Test

If unsure, check:

1. What's the melting point? (high → ionic; low → covalent)
2. Does it conduct electricity in water? (yes → ionic; no → covalent)
3. Is it gas/liquid at room temperature? (often → covalent)

[NCERT — fundamental]

Example 9: NCERT — Comparing $H_2$ and $HCl$

Compare the bonding in $H_2$ and $HCl$.

Solution:

H₂ Molecule

Atoms: 2 hydrogen atoms. Each H has 1 electron.

Bond formation: Both H atoms share 1 electron each — total 1 pair. Each H now has 2 electrons (helium config).

Structure: H — H

Type: Pure (non-polar) covalent bond.

Both atoms are identical → equal sharing of electrons → no charge separation.

HCl Molecule

Atoms: 1 hydrogen + 1 chlorine. H has 1 electron; Cl has 7 valence electrons.

Bond formation: H shares 1 electron with Cl. Cl shares 1 electron with H. Total 1 pair shared.

H has 2 electrons (helium config). Cl has 8 electrons (argon config).

Structure: H — Cl

Type: Polar covalent bond.

Cl is more electronegative than H → electrons are pulled towards Cl. Result: Cl has slight negative charge ($\delta^-$); H has slight positive charge ($\delta^+$).

Key Differences

Property	H₂	HCl
Bond type	Pure covalent	Polar covalent
Sharing	Equal	Unequal
Charge separation	None	Yes ($\delta^+ \delta^-$)
Dipole	No	Yes
In water	Insoluble	Highly soluble (ionises)
Conductivity	No	Yes (in water)

Why HCl Behaves Like Acid in Water?

Although HCl is covalent in pure state, in water:

$$HCl + H_2O \rightarrow H_3O^+ + Cl^-$$

Water's polarity helps ionise HCl.

Free H⁺ (or H₃O⁺) → makes the solution acidic. Free Cl⁻ → balances the charge.

This gives HCl its ionic-like behaviour in water.

Hence: HCl gas — covalent; HCl in water — acts ionic.

Key Insight: Polarity of bonds → polarity of molecules → behaviour in water.

[Board: 3-5 mark]

Example 10: A Concluding Question

(a) Why does carbon prefer covalent bonding? (b) Draw electron-dot structure of methane. (c) Compare: $H_2O$ and $CO_2$. (d) Explain tetravalency with one example.

Solution:

(a) Why Carbon Prefers Covalent Bonding

Carbon configuration: 2, 4. Has 4 valence electrons.

Three options:

1. Lose 4 electrons → C⁴⁺ (too much energy needed).
2. Gain 4 electrons → C⁴⁻ (small nucleus can't hold 10 electrons stably).
3. Share 4 electrons → covalent bonds → both stable.

Hence carbon prefers covalent bonding.

(b) Electron-Dot Structure of Methane (CH₄)

C has 4 valence electrons (4 dots). Each H has 1 electron (1 dot). Each H shares 1 electron with C.

        H (1 dot)
        :
        :
H : : C : : H
        :
        :
        H (1 dot)

Or simplified:

      H
      |
  H — C — H
      |
      H

4 single bonds; C has 8 electrons in outer shell; each H has 2.

(c) Comparison: H₂O and CO₂

Property	H₂O	CO₂
Bond	2 single (O-H)	2 double (C=O)
Structure	Bent (V-shape)	Linear
State at 25°C	Liquid	Gas
Polarity	Polar	Non-polar (overall)
Bond polarity	O-H polar	C=O polar
In water	Universal solvent	Slightly soluble

H₂O is bent because of lone pairs on O. CO₂ is linear because no lone pairs on C.

(d) Tetravalency with One Example

Tetravalency: Carbon's ability to form 4 covalent bonds.

Example: Methane ($CH_4$)

      H
      |
  H — C — H
      |
      H

C has 4 single bonds with 4 H atoms. Each H provides 1 electron; C provides 1 electron per bond. Total: 4 shared pairs.

Each C-H bond:

• 2 electrons (1 from each atom).
• C ends with 8 electrons (octet).
• H ends with 2 electrons (helium config).

All 4 bonds are equal — methane is tetrahedral in shape.

This is the simplest example of carbon's tetravalency.

[Board — 5-mark mixed]

Example 11: A Numerical — Counting Electrons

In $CH_4$, count the total number of: (a) Valence electrons of C. (b) Valence electrons of H atoms. (c) Total shared electrons. (d) Total bonds formed.

Solution:

Atoms in CH₄:

• 1 C atom
• 4 H atoms

(a) Valence electrons of C

C: configuration 2, 4. Valence electrons = 4.

(b) Valence electrons of H atoms

Each H has 1 valence electron. 4 H atoms × 1 = 4 valence electrons.

(c) Total shared electrons

Each C-H bond shares 2 electrons (1 from C, 1 from H). 4 bonds × 2 electrons = 8 shared electrons.

Or: C contributes 4, H atoms contribute 4 → total 8 shared.

(d) Total bonds formed

4 C-H single bonds. Total bonds = 4.

Verification

Octet rule check:

• C: 8 electrons (4 bonds × 2 each) ✓
• Each H: 2 electrons (1 bond × 2) ✓

All atoms satisfied.

Bond Distribution

         H
         |
     H — C — H     ← 4 single bonds
         |
         H

Tetrahedral structure — 109.5° bond angle.

[Board: 3-mark numerical]

Example 12: NCERT — Why Diamond is Hard, Graphite is Soft

Both diamond and graphite are covalent forms of carbon. Explain why diamond is the hardest substance and graphite is soft.

Solution:

Diamond — Structure

Each C atom is bonded to 4 other C atoms in a tetrahedral arrangement. All 4 valence electrons are used in covalent bonds. Forms a giant 3D network.

Properties due to structure:

• All bonds are equal and strong.
• 3D rigid network.
• No layers to slide.
• Hardest natural substance.
• High melting point (~3500°C).
• Non-conductor (no free electrons).
• High density (3.5 g/cm³).

Graphite — Structure

Each C atom is bonded to only 3 other C atoms in a planar (flat) hexagonal arrangement. One electron per C atom is delocalised (free). Forms layers of hexagonal sheets. Layers held together by weak van der Waals forces.

Properties due to structure:

• Layers can slide over each other → soft and slippery.
• Free electrons → good conductor of electricity.
• Lower density (2.2 g/cm³).
• Used in pencils (slides off on paper).
• Used as a lubricant.
• Used in electrodes (conducts).

Comparison

Property	Diamond	Graphite
Bonds per C	4	3
Structure	3D network	2D layers
Free electrons	None	Yes
Hardness	Hardest	Soft
Conductivity	None	Good
Density	3.5 g/cm³	2.2 g/cm³
Use	Cutting, jewellery	Pencils, electrodes

Key Insight

Same element (C), different structure → drastically different properties.

This shows the importance of arrangement in determining properties.

Both are 'allotropes' — different forms of the same element.

(More on allotropes in Section 2!)

[NCERT — important]

Example 13: Polar vs Non-Polar Covalent Bonds

Classify the following bonds as polar or non-polar covalent: (a) H — H (b) C — H (c) O — H (d) C = O (e) N — H (f) Cl — Cl

Solution:

Principle

Polarity depends on electronegativity difference:

• Pure covalent (non-polar): same atoms (no difference).
• Polar covalent: different atoms (some difference).

Electronegativity values:

• H: 2.1
• C: 2.5
• N: 3.0
• O: 3.5
• Cl: 3.0

Analysis

(a) H — H

Same atoms; difference = 0. ✗ Non-polar.

(b) C — H

Difference: 2.5 - 2.1 = 0.4 (small). ✗ Almost non-polar (slightly polar).

(c) O — H

Difference: 3.5 - 2.1 = 1.4 (significant). ✓ Polar covalent. O has slight δ⁻; H has slight δ⁺.

(d) C = O

Difference: 3.5 - 2.5 = 1.0. ✓ Polar covalent. O has δ⁻; C has δ⁺.

(e) N — H

Difference: 3.0 - 2.1 = 0.9. ✓ Polar covalent. N has δ⁻; H has δ⁺.

(f) Cl — Cl

Same atoms; difference = 0. ✗ Non-polar.

Summary

Bond	Difference	Type
H — H	0	Non-polar
C — H	0.4	Almost non-polar
O — H	1.4	Polar
C = O	1.0	Polar
N — H	0.9	Polar
Cl — Cl	0	Non-polar

Practical Significance

Polar bonds → polar molecules → solubility in water. Non-polar bonds → non-polar molecules → solubility in oil.

Example:

• H₂O (polar) — dissolves NaCl, sugar.
• CCl₄ (non-polar) — dissolves naphthalene, oil.
• Petrol (non-polar) — won't mix with water.

'Like dissolves like' — fundamental rule.

[Board: 3-mark]

Example 14: Summary Question

(a) Define covalent bond. (b) Why is carbon's valency 4? (c) Draw structure of $C_2H_4$ (ethene). (d) Compare melting point of NaCl vs CH₄. (e) Mention 3 properties of covalent compounds.

Solution:

(a) Covalent Bond

A chemical bond formed by mutual sharing of electron pairs between two atoms — so that both achieve a stable noble gas configuration (octet).

Examples: H-H, C-C, O-H, etc.

(b) Carbon's Valency = 4

Reason: Configuration 2, 4 → 4 valence electrons.

To complete octet (8 electrons), C needs 4 more. By sharing 4 electrons → 4 covalent bonds.

Hence valency = 4 (tetravalency).

(c) Structure of Ethene ($C_2H_4$)

2 C atoms + 4 H atoms. 1 C=C double bond + 4 C-H single bonds.

    H       H
     \     /
      C = C
     /     \
    H       H

Bonds:

• Each C: 1 double bond (counts as 2) + 2 single (with H) = 4 bonds ✓
• Each H: 1 single bond ✓

Total: 6 bonds (1 C=C + 4 C-H + considering double counts as 2).

(d) NaCl vs CH₄ Melting Point

Compound	M.P.	Reason
NaCl	801°C	Strong ionic bonds (electrostatic)
CH₄	-182°C	Weak intermolecular forces

Difference: ~983°C!

Reason:

• NaCl: ions held strongly in 3D crystal lattice.
• CH₄: small non-polar molecules, only van der Waals forces.

(e) 3 Properties of Covalent Compounds

1. Low melting and boiling points. Examples: H₂O (B.P. 100°C), CH₄ (B.P. -161°C).

2. Insoluble in water (mostly). Examples: oil, petrol, naphthalene. Exceptions: HCl, sugar (polar).

3. Poor conductors of electricity. Reason: no free ions/electrons. Exceptions: acids (in water), graphite.

[Board: 5-mark mixed]

Example 15: An Interesting — Lone Pairs

What are lone pairs? Identify lone pairs in $H_2O$, $NH_3$, and $HCl$.

Solution:

Definition

'Lone pair' = a pair of valence electrons that is not shared in any bond.

These belong to one atom only — no bonding partner.

H₂O — 2 Lone Pairs

O has 6 valence electrons. 2 are shared with H atoms (in 2 O-H bonds). 4 remain as lone pairs (2 pairs).

        ..
    H — O — H
        ..

The two dots above and below O represent 2 lone pairs.

NH₃ — 1 Lone Pair

N has 5 valence electrons. 3 are shared with H atoms (in 3 N-H bonds). 2 remain as 1 lone pair.

        ..
    H — N — H
        |
        H

The two dots on N represent 1 lone pair.

HCl — 3 Lone Pairs

Cl has 7 valence electrons. 1 is shared with H (in 1 H-Cl bond). 6 remain as 3 lone pairs.

        : :
    H — Cl
        : :

Cl has 3 lone pairs.

Why Lone Pairs Matter

1. They affect molecular shape.

• $H_2O$: 2 lone pairs → V-shaped (bent).
• $NH_3$: 1 lone pair → pyramidal.
• $CH_4$: 0 lone pairs → tetrahedral.

2. They make molecules polar.

• Lone pairs create asymmetry → dipole moment.
• Hence H₂O and NH₃ are polar.

3. They allow further bonding.

• Lone pairs can form 'coordinate' or 'dative' bonds.
• Example: $NH_4^+$ (ammonium) — N's lone pair shared with H⁺.

Summary

Molecule	Lone pairs	Shape	Polarity
H₂O	2	Bent	Polar
NH₃	1	Pyramidal	Polar
CH₄	0	Tetrahedral	Non-polar
HCl	3 (on Cl)	Linear	Polar

Lone pairs — invisible but important!

[Board: 3-mark]

Example 16: An Application — Bond Energy and Stability

Which bond is most stable: C-C, C=C, or C≡C? Why?

Solution:

Comparison Table

Bond	Shared Pairs	Length	Energy (kJ/mol)
C — C	1	154 pm	348
C = C	2	134 pm	614
C ≡ C	3	120 pm	839

Most Stable: C ≡ C (Triple Bond)

Triple bond has highest bond energy (839 kJ/mol). Hence requires most energy to break. Most stable in terms of energy.

But — Most Reactive Different Question

'Stable' (high bond energy) ≠ 'unreactive'.

Triple bonds are more reactive in many ways:

• Can undergo addition reactions.
• Can break partially (3 → 2 → 1).
• More electron density → easier to attack.

Single bonds are less reactive:

• All electrons are bonded.
• No multiple bond to break.
• Stable to addition.

Trade-off — Stability vs Reactivity

Triple bond:

• High energy (stable).
• But can react more (lose 1 pair to become double).

Single bond:

• Lower energy (less stable in terms of energy).
• But less reactive (no extra pairs to lose).

Practical Examples

Ethyne ($C_2H_2$):

• Triple bond.
• Reacts easily with H₂, Br₂.
• Used in oxyacetylene welding.

Ethene ($C_2H_4$):

• Double bond.
• Reacts with H₂, Br₂ (decolourises bromine water — a test).

Ethane ($C_2H_6$):

• Single bond.
• Less reactive (only substitution reactions).
• Stable, used as fuel.

Industrial Applications

Cracking heavy hydrocarbons:

• Long single-bond chains broken.
• Form smaller molecules with double bonds.
• These are more reactive — used to make plastics, etc.

Polymerisation:

• Many ethene molecules join via double bond.
• Form polyethylene plastic.

Conclusion

Most stable: C ≡ C (highest bond energy). Most reactive: C ≡ C (multiple bonds).

A paradox: same bond can be both stable (energetically) and reactive (chemically)!

[Board: 5-mark]