Topic 2.1 Notes – Types of Chemical Bonds

1. Why Atoms Bond and What Controls Bond Type

Atoms bond because the attractions between positive nuclei and negative electrons lower potential energy. When those attractions outweigh repulsions, a stable bond forms.

The electrons involved are the valence electrons:

• Outer-shell s and p electrons
• The only electrons that participate in bonding
• The ones that determine chemical behavior

Whether electrons are transferred, shared equally, or shared unequally depends mostly on electronegativity (EN).

Electronegativity and Periodic Trends

Electronegativity is an atom’s ability to attract electrons in a bond.

Here’s the trend you need locked in.

• Increases left → right
• Decreases down a group

Highest values are found near the top right corner, which is why fluorine is the most electronegative element.

Why this happens comes straight from Coulomb’s Law:

$$F \propto \frac{q_{1} q_{2}}{r^{2}}$$

• More protons (larger nuclear charge) → stronger attraction
• Greater distance between nucleus and valence electrons → weaker attraction

Across a period:

• Proton count increases
• Valence shell stays the same
• Attraction gets stronger → EN increases

Down a group:

• Valence electrons are farther from nucleus
• Shielding increases
• Attraction weakens → EN decreases

This explains why fluorine is extremely electronegative and cesium is not.

Bond strength and bond type both depend on this same charge-and-distance idea.

2. The Three Types of Chemical Bonds

Ionic Bonding

Forms when electrons are transferred, usually from a metal to a nonmetal.

• Metal loses electron → cation (+)
• Nonmetal gains electron → anion (−)
• Attraction is purely electrostatic

Structure matters here.

Sodium chloride crystal lattice

Instead of existing as separate $\ce{Na+}$ and $\ce{Cl-}$ pairs, ionic compounds form extended 3D networks of alternating positive and negative ions.

That repeating 3D structure is a crystal lattice.

Properties

• High melting and boiling points (strong electrostatic attractions)
• Brittle
• Conduct electricity when molten or dissolved
• Do not conduct as solids (ions can’t move)

What makes one ionic compound stronger than another?

1. Charge magnitude
   • $\ce{CaO}$ (2+ and 2−) > $\ce{NaCl}$ (1+ and 1−)
2. Ion size
   • Smaller ions → shorter distance → stronger attraction

On exams, compare charge first, then size. Charge has the bigger effect.

Covalent Bonding

Forms between nonmetals that share valence electrons.

Nonpolar Covalent

• Similar EN
• Electrons shared equally
• No partial charges
• Example: C-H is effectively nonpolar

Polar Covalent

• Different EN
• Unequal sharing
• More electronegative atom becomes δ−
• Less electronegative atom becomes δ+

Nonpolar, polar, and ionic bonding comparison

The three panels show the progression from equal sharing (Cl-Cl), to unequal sharing with partial charges (H-Cl), to full electron transfer (Na⁺ and Cl⁻).

Greater EN difference → larger bond dipole.

Important idea: Bonding is a continuum.

• All polar bonds have some ionic character.
• There is no sharp cutoff between ionic and covalent.

General guideline (not a rule):

• 0-0.4 → nonpolar covalent
• 0.4-1.7 → polar covalent
• >1.7 → mostly ionic

But AP questions often expect you to confirm using properties, not just EN difference.

Molecular (covalent) substances

• Lower melting points
• Poor electrical conductivity in all states

Metallic Bonding

Occurs between metal atoms.

Valence electrons are delocalized. They form a mobile “sea” around positive metal cations.

This explains metallic properties:

• Conduct electricity as solids (mobile electrons)
• Malleable and ductile (layers slide without breaking bonds)
• Shiny (electrons absorb and re-emit light)
• Good thermal conductors

Stronger metallic bonding occurs when:

• More valence electrons are contributed
• Metal atoms are smaller

3. How to Determine Bond Type from Elements and Properties

When you’re given a compound, think in layers.

1. Element types

• Metal + nonmetal → usually ionic
• Nonmetal + nonmetal → covalent
• Metal + metal → metallic

2. Electronegativity difference

• Larger difference → more ionic character

3. Observed properties (most reliable)

AP questions love giving you melting point and conductivity data instead of saying the bond type directly. You’re expected to connect structure to properties.

4. The Big Relationship to Lock In

Everything connects like this:

Periodic trend → electronegativity → electron distribution → bond type → structure → properties

If you can explain that chain clearly, you’re thinking the way AP graders want.