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Study Guide · Chapter 4

Chemical Bonding & Molecular Structure

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Introduction: Why Atoms Bond

Imagine atoms as lonely individuals seeking companionship. Some are willing to give electrons entirely (like sodium), others are desperate to receive them (like chlorine), and still others prefer to share electrons fairly (like carbon). These different personalities of atoms lead to different types of chemical bonds—and those bonds determine everything about a molecule's properties: whether it dissolves in water, its melting point, its electrical conductivity, its reactivity.

In this chapter, we'll explore how atoms combine through ionic bonding, covalent bonding, and metallic bonding. We'll decode molecular shapes using VSEPR theory and learn why water is bent (not straight) and why carbon dioxide is linear (not bent). Finally, we'll understand hybridization—a quantum mechanical concept that explains how atoms reshape their electron orbitals to form multiple bonds.

Part 1: Three Main Types of Chemical Bonding

Ionic Bonding: Electron Transfer

Definition: Formation of bonds between a metal and a nonmetal through complete transfer of electrons from the metal to the nonmetal.

Mechanism:

Metal atom loses electrons (becomes a cation: e.g., Na → Na⁺ + e⁻)
Nonmetal atom gains electrons (becomes an anion: e.g., Cl + e⁻ → Cl⁻)
Electrostatic attraction between oppositely charged ions holds them together

Example: Sodium Chloride Formation

Na atom (11 protons, 11 electrons)
Configuration: 1s² 2s² 2p⁶ 3s¹
               
               Loses 1 electron from 3s¹
               ↓
               Na⁺ (11 protons, 10 electrons)
               Configuration: 1s² 2s² 2p⁶
               
Cl atom (17 protons, 17 electrons)
Configuration: 1s² 2s² 2p⁶ 3s² 3p⁵
               
               Gains 1 electron to fill 3p⁶
               ↓
               Cl⁻ (17 protons, 18 electrons)
               Configuration: 1s² 2s² 2p⁶ 3s² 3p⁶

Result: Na⁺ and Cl⁻ are electrostatically attracted
Formula: NaCl (sodium chloride, or table salt)

Conditions for Ionic Bonding:

Electronegativity difference > 1.7 (using Pauling scale)
Metal (low EN) loses electrons to nonmetal (high EN)

[Memory Hook] Ionic = electron transfer; metal loses, nonmetal gains; electrostatic attraction

Properties of Ionic Compounds

Property	Explanation
Solid state at room temp	Strong electrostatic forces hold ions in rigid lattice
High melting/boiling points	Requires breaking strong ionic bonds
Soluble in polar solvents (water)	Water's polar nature dissolves ion lattice
Conduct electricity when molten/dissolved	Free ions can move as electric charge carriers
Brittle	Shifting layers brings like-charges together, causing repulsion

Example: Calcium Carbonate (CaCO₃)

Limestone, chalk, seashells
High melting point (1612°C) due to ionic bonding
Insoluble in water (but dissolves in acids)

Covalent Bonding: Electron Sharing

Definition: Formation of bonds through sharing of electrons between nonmetal atoms.

Mechanism:

Two atoms bring unpaired electrons close together
Electron clouds overlap (overlap region)
Shared electrons attract both nuclei simultaneously, holding atoms together

Example: Hydrogen Molecule (H₂)

H atom: 1 electron in 1s orbital (•)
H atom: 1 electron in 1s orbital (•)

Bonding:
    •                    ••          
    H    +    H    →    H—H    (H₂)
    •            •      

The two electrons (in overlap region) are attracted to both nuclei,
creating a bond. The bond is represented by a line (—) or dots (••).

Electronegativity and Bond Type:

EN Difference	Bond Type	Electron Distribution
< 0.4	Nonpolar covalent	Electrons shared equally
0.4 – 1.7	Polar covalent	Electrons shared unequally; electron density shifts toward more EN atom
> 1.7	Ionic	Complete electron transfer

Example: Water (H₂O)

EN(O) ≈ 3.4, EN(H) ≈ 2.1
Difference: 3.4 - 2.1 = 1.3 → Polar covalent
Oxygen attracts electrons more strongly; electrons spend more time near oxygen
Result: Oxygen becomes partially negative (δ-), hydrogens become partially positive (δ+)

     δ+    O    δ+
    H  \  //  /  H
         δ-

The bent shape (105°) creates a net dipole moment pointing toward oxygen.

[Memory Hook] Covalent = electron sharing; EN difference < 1.7; creates molecules

Properties of Covalent Compounds

Property	Explanation
Gases, liquids, or solids at room temp	Weak intermolecular forces (van der Waals) between molecules
Low melting/boiling points	Only intermolecular forces, not covalent bonds, break
Poor electrical conductivity	No free ions; electrons are localized
Many soluble in nonpolar solvents	Like-dissolves-like principle
Often exist as discrete molecules	Atoms bonded within molecules; weak inter-molecular attraction

Example: Sugar (Sucrose, C₁₂H₂₂O₁₁)

Melts at 186°C (much lower than ionic NaCl at 801°C)
Dissolves in water (polar solvent) but not in oil (nonpolar)
Non-conducting as pure solid or liquid

Single, Double, and Triple Bonds

Single Bond (—): One pair of electrons shared

Example: H—H (hydrogen), C—H (in methane), Cl—Cl

Double Bond (=): Two pairs of electrons shared

Example: O=O (oxygen gas), C=O (carbon dioxide), C=C (in ethene)

Atoms are pulled closer together in double bonds → shorter bond lengths

Triple Bond (≡): Three pairs of electrons shared

Example: N≡N (nitrogen gas), C≡C (in acetylene)

Strongest bonds; shortest bond lengths; hardest to break

[Memory Hook] Single < double < triple; increasing bond strength and shortness

[Exam Trap] Students confuse bond length. Triple bonds are shortest, not longest. More electron pairs = stronger attraction = closer atoms.

Metallic Bonding

Definition: Bonding in metals through a "sea of electrons" model: valence electrons are delocalized across all atoms.

Mechanism:

Metal atoms lose valence electrons into a common "electron sea"
These electrons move freely throughout the metal structure
Metal cations are held together by attraction to the electron sea

Metallic Bonding (Iron, Fe):

    Fe²⁺  Fe²⁺  Fe²⁺  Fe²⁺
     •     •     •     •
    ////////// electron sea //////////
     •     •     •     •
    Fe²⁺  Fe²⁺  Fe²⁺  Fe²⁺

The sea of electrons binds all cations together.

Properties of Metals:

Property	Explanation
High electrical/thermal conductivity	Delocalized electrons move freely
Malleable and ductile	Electrons adjust to new positions when atoms shift
High melting points	Strong metallic bonds
Luster (shine)	Delocalized electrons absorb and emit light easily
Insoluble in polar or nonpolar solvents	Metallic structure doesn't dissolve

[Memory Hook] Metallic bonding = sea of electrons; explains conductivity and malleability

Part 2: Electronegativity and Bond Polarity

Electronegativity Difference Decision Tree

Electronegativity Difference
            |
            ├─ < 0.4 → Nonpolar Covalent (e.g., Cl—Cl, C—H)
            |
            ├─ 0.4 to 1.7 → Polar Covalent (e.g., H—Cl, H—O—H)
            |
            └─ > 1.7 → Ionic (e.g., Na—Cl, Ca—O)

Polar Covalent Bonds and Dipoles

In a polar covalent bond, the electron pair is pulled toward the more electronegative atom, creating a dipole—a separation of partial charges.

Example: Hydrogen Chloride (HCl)

EN(Cl) = 3.0, EN(H) = 2.1
Difference: 0.9 → Polar covalent
Cl is more electronegative; electron density shifts toward Cl

H—Cl    becomes    H^δ+—Cl^δ-
                    ← dipole arrow points toward negative end

This is a polar molecule: it has a dipole moment (μ)

Dipole Moment (μ): Measures the magnitude of the separated charge

Units: Debye (D)
HCl has μ ≈ 1.08 D
H₂O has μ ≈ 1.85 D (stronger dipole due to bent geometry)
Nonpolar molecules: μ ≈ 0

[Memory Hook] Polar = electrons not shared equally; electronegativity difference creates dipole

Part 3: VSEPR Theory (Valence Shell Electron Pair Repulsion)

VSEPR predicts molecular shapes based on a simple idea: electron pairs (both bonding and lone pairs) repel each other and arrange to be as far apart as possible.

Steps to Predict Molecular Shape

Write Lewis structure (dots around symbol showing valence electrons)
Count electron groups on central atom:
- Bonding pairs (single, double, triple bonds count as 1 group)
- Lone pairs
Determine electron geometry (arrangement of all electron groups)
Determine molecular geometry (arrangement of atoms only, ignoring lone pairs)
Predict bond angles and 3D shape

Electron Geometries and Molecular Shapes

2 Electron Groups: Linear

Electron geometry: Linear
Molecular geometry: Linear (if no lone pairs), Bent (if 1 lone pair)

Example: CO₂ (Carbon dioxide)

Central atom: Carbon (4 valence electrons)
Bonding: O=C=O (two double bonds, 2 electron groups)
Shape: Linear, 180° bond angle

O=C=O     (linear; nonpolar due to symmetry)

3 Electron Groups: Trigonal Planar

Electron geometry: Trigonal planar
Molecular geometry: Trigonal planar (3 bonds, 0 lone pairs), Bent (2 bonds, 1 lone pair)

Example: BF₃ (Boron trifluoride)

Central atom: Boron (3 valence electrons)
Bonding: Three B—F single bonds (3 electron groups)
Shape: Trigonal planar, 120° bond angles

        F
        |
    F—B—F
    (top view)

    Electron geometry: Trigonal planar
    Molecular geometry: Trigonal planar

4 Electron Groups: Tetrahedral

Electron geometry: Tetrahedral (109.5° angles)
Molecular geometry depends on lone pairs:
- 4 bonds, 0 lone pairs: Tetrahedral (e.g., CH₄, CF₄)
- 3 bonds, 1 lone pair: Trigonal pyramidal (e.g., NH₃)
- 2 bonds, 2 lone pairs: Bent (e.g., H₂O)

Example: Water (H₂O)

Central atom: Oxygen (6 valence electrons)
Bonding: Two O—H single bonds (2 bonding electron groups)
Lone pairs: 2 (filled p² on oxygen)
Total electron groups: 4
Electron geometry: Tetrahedral
Molecular geometry: Bent (or angular)
Bond angle: ~105° (less than 109.5° due to lone pair repulsion)

         H
         |
        O   (bent shape; dashed line = lone pair)
        |  ..
        H

Dipole moment: μ ≈ 1.85 D (polar molecule)
Polarity: The two O—H dipoles don't cancel; net dipole points between H atoms

[Memory Hook] VSEPR = electron pairs repel; arrange maximum distance apart

Example: Methane (CH₄)

Central atom: Carbon (4 valence electrons)
Bonding: Four C—H single bonds
Shape: Tetrahedral, ~109.5° bond angles
Nonpolar (even though C—H bonds are slightly polar, symmetry cancels dipoles)

           H
           |
       H—C—H    (tetrahedral; perspective drawing with wedge/dash bonds)
           |
           H

5 Electron Groups: Trigonal Bipyramidal

Axial bonds (top/bottom): 180° apart
Equatorial bonds (middle): 120° apart
Lone pairs preferentially occupy equatorial positions (more room)

Example: PCl₅ (Phosphorus pentachloride)

Central atom: Phosphorus (5 valence electrons)
Bonding: Five P—Cl single bonds
Shape: Trigonal bipyramidal

         Cl (axial)
         |
    Cl—P—Cl  (trigonal bipyramidal)
         |
        Cl (axial)
      (1 Cl in equatorial plane)

3 equatorial Cl atoms in plane; 2 axial Cl atoms above/below

6 Electron Groups: Octahedral

All electron groups arranged 90° apart
Forms an octahedron (8-sided polyhedron)

Example: SF₆ (Sulfur hexafluoride)

Central atom: Sulfur (6 valence electrons)
Bonding: Six S—F single bonds
Shape: Octahedral, 90° bond angles

          F (top)
          |
      F—S—F
      |   |
      F—S—F
          |
          F (bottom)
    (simplified 2D view; actually 3D octahedron)

Part 4: Hybridization

Hybridization is the mixing of atomic orbitals to create new hybrid orbitals suitable for bonding.

Why Hybridization?

The Bohr/Schrödinger model says carbon has electrons in 2s and 2p orbitals:

2s is spherical
2p orbitals are dumbbell-shaped

Yet methane (CH₄) has four identical C—H bonds at 109.5° tetrahedral angles. How can spherical s and dumbbell p orbitals create four identical bonds?

Answer: The 2s and 2p orbitals hybridize into four identical sp³ hybrid orbitals, each oriented toward a corner of a tetrahedron.

Types of Hybridization

sp Hybridization (Linear)

One s orbital + one p orbital mix
Creates 2 hybrid orbitals at 180° angles
Associated with: single and triple bonds

Example: Acetylene (HC≡CH)

     sp hybridized carbon atoms
     
H—C≡C—H

Each carbon: 2 sp hybrid orbitals (one for σ bond to other C, one for σ bond to H)
           2 unhybridized p orbitals (form π bond with other carbon)
           
Result: Linear molecule, 180° bond angle

sp² Hybridization (Trigonal Planar)

One s orbital + two p orbitals mix
Creates 3 hybrid orbitals at 120° angles
Associated with: double bonds

Example: Ethene (C₂H₄, or H₂C=CH₂)

           H    H
            \  /
             C=C
            /  \
           H    H

Each carbon: 3 sp² hybrid orbitals (two for σ bonds, one participates in π bond)
           1 unhybridized p orbital (forms π bond with other carbon)
           
Result: Planar molecule, 120° bond angles, can't rotate around double bond

sp³ Hybridization (Tetrahedral)

One s orbital + three p orbitals mix
Creates 4 hybrid orbitals at 109.5° angles
Associated with: single bonds

Example: Methane (CH₄)

Each carbon: 4 sp³ hybrid orbitals (one for each C—H σ bond)
           No unhybridized p orbitals
           
Result: Tetrahedral, 109.5° bond angles, can freely rotate

sp³d Hybridization (Trigonal Bipyramidal)

One s orbital + three p orbitals + one d orbital mix
Creates 5 hybrid orbitals
Associated with: compounds of period 3+ elements

Example: PCl₅

Central P: 5 sp³d hybrid orbitals (five P—Cl σ bonds)
Result: Trigonal bipyramidal

sp³d² Hybridization (Octahedral)

One s orbital + three p orbitals + two d orbitals mix
Creates 6 hybrid orbitals
Associated with: transition metals and heavier elements

Example: SF₆

Central S: 6 sp³d² hybrid orbitals (six S—F σ bonds)
Result: Octahedral

[Memory Hook] sp (linear, 180°) < sp² (planar, 120°) < sp³ (tetrahedral, 109.5°)

Part 5: Resonance Structures

Resonance occurs when a molecule can be represented by multiple valid Lewis structures that differ only in electron placement (not atom positions).

Example: Ozone (O₃)

Structure 1:          Structure 2:
     O                      O
    //                     \\
   O—O                     O—O
    
Both structures are equally valid!
The real structure is an average (hybrid) of both.

The ozone molecule is not "Structure 1 sometimes and Structure 2 sometimes." Rather, it's a hybrid of both, with the middle O—O bond having partial double bond character.

Result of resonance:

Increased stability (resonance stabilization)
Shorter bonds than single bonds but longer than double bonds
The true structure can't be represented by a single Lewis structure

[Memory Hook] Resonance = molecule is average of multiple Lewis structures; increased stability

Part 6: Hydrogen Bonding (Special Intermolecular Force)

Hydrogen bonding is an unusually strong intermolecular force (not a true chemical bond) that occurs when hydrogen is bonded to highly electronegative atoms: fluorine (F), oxygen (O), or nitrogen (N).

Mechanism

Molecule 1 (H-bonded to F/O/N):     Molecule 2 (F/O/N with lone pair):
    δ+ δ-                                  :
   H—F        ←hydrogen bond→          H—O—H
   
The δ+ hydrogen of molecule 1 is attracted to the lone pair (δ-) on molecule 2.

Why Hydrogen Bonding is Strong

Electronegativity: F, O, N are highly electronegative; pull electron density away from H
Size: H is tiny; when bonded to F/O/N, it's very positively polarized
Geometry: Lone pairs on F/O/N are favorably oriented for bonding

Consequences of Hydrogen Bonding

Water (H₂O):

Boiling point: 100°C
Without hydrogen bonding, predicted boiling point: ~-80°C (similar to H₂S)
Density maximum at 4°C (unusual; most liquids contract upon cooling)
Ice is less dense than water (ice floats due to H-bonding structure)

Ammonia (NH₃):

Boiling point: -33°C
Without hydrogen bonding, predicted: ~-70°C

Hydrogen Fluoride (HF):

Boiling point: 19.5°C
Without hydrogen bonding, predicted: ~-110°C

Biological Importance:

Proteins: Hydrogen bonds between carbonyl oxygen and amine hydrogen stabilize secondary structures (α-helices, β-sheets)
DNA: Hydrogen bonds between base pairs (A—T has 2 H-bonds; G—C has 3) hold the double helix together and allow separation for replication

[Memory Hook] Hydrogen bonding = H bonded to F/O/N; creates unusual properties (high boiling points, ice floats)

[Exam Trap] Hydrogen bonding is intermolecular (between molecules), not intramolecular (within molecules). It's much weaker than covalent bonds.

Conclusion

Understanding chemical bonding—from ionic to covalent to metallic—is the gateway to predicting molecular properties. VSEPR theory and hybridization let you visualize how atoms arrange themselves in 3D space. Master these concepts, and you can explain why sodium reacts violently with water, why methane is a gas while silicon carbide is a hard solid, and why water has such extraordinary properties.

23 MCQ Questions

Q1: Which electronegativity difference indicates an ionic bond?

A) < 0.4
B) 0.4 – 1.7
C) > 1.7
D) Cannot be determined from EN difference

Q2: In the formation of sodium chloride (NaCl), what happens to the electron from sodium?

A) It is completely transferred to chlorine
B) It is shared equally with chlorine
C) It remains with sodium
D) It is lost to the environment

Q3: Which property is characteristic of ionic compounds?

A) Poor electrical conductivity in solid state
B) Low melting points
C) Solubility in nonpolar solvents like oil
D) Soft and malleable

Q4: A compound is formed between two nonmetal atoms with an electronegativity difference of 1.2. What type of bond is this?

A) Ionic bond
B) Nonpolar covalent bond
C) Polar covalent bond
D) Metallic bond

Q5: In a triple bond (≡), how many electron pairs are shared?

A) 1
B) 2
C) 3
D) 4

Q6: What is the main feature of metallic bonding?

A) Complete transfer of electrons
B) Sharing of electrons between two atoms
C) Delocalized electrons moving freely in a "sea"
D) Hydrogen bonding between atoms

Q7: According to VSEPR theory, electron pairs around a central atom:

A) Attract each other
B) Repel each other and arrange to maximize distance
C) Have no effect on molecular shape
D) Form covalent bonds with the central atom

Q8: Water (H₂O) has a bent molecular geometry because:

A) It has 3 electron groups (2 bonding, 1 lone pair)
B) It has 4 electron groups (2 bonding, 2 lone pairs)
C) Hydrogen is more electronegative than oxygen
D) It forms hydrogen bonds

Q9: What is the molecular geometry of methane (CH₄)?

A) Linear
B) Trigonal planar
C) Tetrahedral
D) Trigonal bipyramidal

Q10: The bond angle in water (H₂O) is approximately:

A) 90°
B) 104.5°
C) 109.5°
D) 120°

Q11: In carbon dioxide (CO₂), which type of hybridization is the central carbon atom undergoing?

A) sp
B) sp²
C) sp³
D) sp³d

Q12: Hybridization of sp² creates how many hybrid orbitals and at what angles?

A) 2 orbitals at 180°
B) 3 orbitals at 120°
C) 4 orbitals at 109.5°
D) 5 orbitals at 90°

Q13: In ethane (C₂H₆), what type of hybridization is present in the carbon atoms?

A) sp
B) sp²
C) sp³
D) sp³d

Q14: Resonance structures in a molecule:

A) Represent different molecules
B) Represent average configurations where electrons aren't fixed
C) Have different atom positions
D) Occur only in ionic compounds

Q15: Which molecules exhibit hydrogen bonding?

A) CH₄ (methane)
B) H₂O (water), HF (hydrogen fluoride), NH₃ (ammonia)
C) CO₂ (carbon dioxide)
D) All of the above

Q16: Hydrogen bonding requires hydrogen bonded to:

A) Any nonmetal
B) Highly electronegative atoms: F, O, or N
C) Only fluorine
D) Carbon or nitrogen only

Q17: Why is the boiling point of water (100°C) much higher than expected?

A) High molecular weight
B) Strong covalent bonds within molecules
C) Extensive hydrogen bonding between molecules
D) Ionic bonding in water

Q18: In phosphorus pentachloride (PCl₅), the hybridization of the central phosphorus is:

A) sp
B) sp²
C) sp³
D) sp³d

Q19: Sulfur hexafluoride (SF₆) has an octahedral shape, indicating which hybridization?

A) sp³d
B) sp³d²
C) d²sp³
D) dsp³

Q20: A polar covalent bond is created when:

A) Atoms have identical electronegativities
B) Atoms have significantly different electronegativities (0.4–1.7 difference)
C) A metal bonds with a nonmetal
D) Hydrogen bonds to F, O, or N

Q21: In the molecule CO₂, the carbon atom has:

A) 2 electron groups and sp hybridization
B) 3 electron groups and sp² hybridization
C) 4 electron groups and sp³ hybridization
D) 5 electron groups and sp³d hybridization

Q22: Which of the following molecules is nonpolar despite having polar bonds?

A) HCl
B) H₂O
C) CO₂
D) NH₃

Q23: The dipole moment of a molecule:

A) Depends only on electronegativities of atoms
B) Depends on bond polarity and molecular geometry
C) Is always present if any polar bonds exist
D) Cannot be predicted from Lewis structures

Answer Key: 1-C, 2-A, 3-A, 4-C, 5-C, 6-C, 7-B, 8-B, 9-C, 10-B, 11-A, 12-B, 13-C, 14-B, 15-B, 16-B, 17-C, 18-D, 19-B, 20-B, 21-A, 22-C, 23-B

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