Example: NaCl (sodium chloride)

$Na \rightarrow Na+ + e-$ (sodium loses 1 electron) $Cl + e- \rightarrow Cl-$ (chlorine gains 1 electron) $Na+ + Cl- \rightarrow NaCl$ (ionic bond forms)

Electronegativity:

• Na: 0.9
• Cl: 3.0
• ΔEN = 2.1 (large difference → ionic)

2. Covalent Bonds

Definition: Sharing of electron pairs between atoms

Formation:

• Electrons are shared between atoms (not transferred)
• Both atoms contribute to shared electron pair
• Occurs between nonmetals
• Can be single, double, or triple bonds

Types of Covalent Bonds:

a) Nonpolar Covalent:

• Equal sharing of electrons
• ΔEN < 0.4
• Symmetrical electron distribution
• Example: H₂, O₂, N₂, Cl₂

b) Polar Covalent:

• Unequal sharing of electrons
• 0.4 < ΔEN < 1.7
• More electronegative atom attracts electrons more
• Creates partial charges (δ+ and δ-)
• Example: H₂O, HCl, NH₃

Characteristics of Covalent Compounds:

• Lower melting/boiling points than ionic (generally)
• Do not conduct electricity (no charged particles)
• Can be gases, liquids, or solids at room temperature
• Soft and flexible (molecular compounds)

Example: H₂O (water)

Each H shares electrons with O:

• O has 6 valence electrons
• Each H has 1 valence electron
• 2 covalent bonds form (O-H bonds)
• O achieves octet, H achieves duet

Electronegativity:

• H: 2.1
• O: 3.5
• ΔEN = 1.4 (moderate difference → polar covalent)

3. Metallic Bonds

Definition: Attraction between metal cations and delocalized electrons

Formation:

• Metal atoms release valence electrons
• Electrons form "sea" of mobile electrons
• Metal cations are embedded in electron sea
• Electrons are delocalized (shared among all atoms)

Characteristics:

• Occur in pure metals and alloys
• Conduct electricity in solid state (mobile electrons)
• Conduct heat well
• Malleable and ductile (layers can slide)
• Lustrous (shiny)
• Variable melting points (depends on metal)

Example: Sodium metal (Na)

• Each Na atom contributes 1 valence electron to "sea"
• Na⁺ ions arrange in crystal lattice
• Electrons move freely throughout structure

Visualization:

$[Na+][Na+][Na+]$ $\text{Electron sea: } e^- e^- e^- e^- e^-$

Electronegativity and Bond Type

Electronegativity (EN): Ability to attract bonding electrons

Pauling Scale: 0.7 (Cs) to 4.0 (F)

Predicting Bond Type

Use difference in electronegativity (ΔEN):

Note: These are guidelines; bonding is a spectrum, not discrete categories.

Example Calculations:

1. H-Cl bond:

• H: 2.1, Cl: 3.0
• ΔEN = |3.0 - 2.1| = 0.9
• Polar covalent

2. Na-Cl bond:

• Na: 0.9, Cl: 3.0
• ΔEN = |3.0 - 0.9| = 2.1
• Ionic

3. C-H bond:

• C: 2.5, H: 2.1
• ΔEN = |2.5 - 2.1| = 0.4
• Nonpolar covalent (borderline)

Bond Polarity

Polar bonds have unequal electron distribution:

• More electronegative atom has partial negative charge (δ-)
• Less electronegative atom has partial positive charge (δ+)
• Creates dipole moment

Representation:

$H-Cl$

With electronegativity values:

• H (2.1) → δ+
• Cl (3.0) → δ-

Dipole arrow: Points from δ+ to δ-

$H \xrightarrow[\delta^-]{ \delta^+} Cl$

or

$H \xrightarrow[-]{+} Cl$

Bond Strength

Factors Affecting Bond Strength

1. Bond order: Triple > Double > Single
2. Atomic size: Smaller atoms form stronger bonds
3. Electronegativity difference: Greater difference can mean stronger bonds (to a point)

Bond Length vs. Bond Energy

Inverse relationship:

• Shorter bond → Stronger bond → Higher bond energy
• Longer bond → Weaker bond → Lower bond energy

Example: Carbon bonds

As bond order increases: length decreases, energy increases

Lewis Structures and Bonding

Lewis structures show:

• Valence electrons as dots
• Bonds as lines (2 electrons per line)
• Lone pairs as pairs of dots

Steps to draw:

1. Count total valence electrons
2. Connect atoms with single bonds
3. Complete octets (start with outer atoms)
4. Place remaining electrons on central atom
5. Form multiple bonds if needed

Example: CO₂

Total valence electrons: 4 (C) + 6×2 (O) = 16

$O=C=O$

• C has 4 bonds (8 electrons) ✓
• Each O has 2 bonds + 2 lone pairs (8 electrons) ✓
• Total: 16 electrons ✓

Comparing Properties

Exceptions and Special Cases

1. Network Covalent Solids

• Giant covalent structures (diamond, SiO₂)
• Very high melting points
• Very hard
• Do not conduct electricity

2. Polar vs. Nonpolar Molecules

Bond polarity ≠ Molecular polarity

• Molecule can have polar bonds but be nonpolar overall (CO₂, CCl₄)
• Symmetry cancels out individual bond dipoles
• Will discuss more in VSEPR theory

3. Transition Between Bond Types

Bonding is a continuum, not discrete:

• No bond is 100% ionic (even NaCl has some covalent character)
• No bond is 100% covalent (even H₂ has tiny dipole from quantum effects)

Summary

1. Ionic bonds: Transfer of electrons, metal + nonmetal, ΔEN > 1.7
2. Covalent bonds: Sharing of electrons, nonmetal + nonmetal, ΔEN < 1.7
   • Nonpolar: ΔEN < 0.4
   • Polar: 0.4 < ΔEN < 1.7
3. Metallic bonds: Delocalized electrons, metal + metal
4. Electronegativity difference predicts bond type
5. Bond strength: Depends on bond order, atomic size, and electronegativity

(b) C-O: $\Delta EN = |3.5 - 2.5| = 1.0$

(c) Cl-Cl: $\Delta EN = |3.0 - 3.0| = 0.0$

Step 2: Apply classification rules

Step 3: Classify each bond

(a) K-Cl: ΔEN = 2.2

• ΔEN > 1.7 → Ionic bond
• K is metal, Cl is nonmetal ✓
• K loses electron → K⁺
• Cl gains electron → Cl⁻

(b) C-O: ΔEN = 1.0

• 0.4 < ΔEN < 1.7 → Polar covalent bond
• Both C and O are nonmetals ✓
• O is more electronegative → partial negative charge (δ-)
• C has partial positive charge (δ+)

(c) Cl-Cl: ΔEN = 0.0

• ΔEN = 0 → Nonpolar covalent bond
• Identical atoms share electrons equally
• No charge separation
• Example: Cl₂ molecule

Answer:

• (a) K-Cl: Ionic
• (b) C-O: Polar covalent
• (c) Cl-Cl: Nonpolar covalent

Summary: Large ΔEN (>1.7) indicates electron transfer (ionic), moderate ΔEN (0.4-1.7) indicates unequal sharing (polar covalent), and small ΔEN (<0.4) indicates equal sharing (nonpolar covalent).