AP Chemistry 2.1: Types of chemical bonds

Atoms interact because of attractions and repulsions between charged particles.

• Remember:

• protons are positive

• electrons are negative

• opposite charges attract

• like charges repel

When two atoms approach each other, several interactions occur:

• each nucleus attracts the other atom's electrons

• the two nuclei repel each other

• the electrons repel each other

A bond forms when the attractive interactions produce a lower-energy, more stable arrangement than the separated atoms.

• This idea— lower energy = greater stability —will appear throughout chemistry.

For AP Chemistry, the major categories are:

• Ionic bonding

• Covalent bonding

• Metallic bonding

• These categories are useful models, but real bonds can have properties between the categories.

An ionic bond is the electrostatic attraction between oppositely charged ions.

Ionic compounds generally form between:

• metal + nonmetal

• Example:

• NaCl

Sodium tends to lose one electron:

• Na → Na⁺ + e⁻

Chlorine tends to gain one:

• Cl + e⁻ → Cl⁻

The resulting ions attract:

• Na⁺ ⋯ Cl⁻

The attraction between the positive and negative ions holds the solid together.

It is common to imagine one Na⁺ attached to one Cl⁻.

• That is not really what solid NaCl looks like.

• Instead, ionic compounds form large, repeating crystal lattices .

Each positive ion is surrounded by negative ions, and each negative ion is surrounded by positive ions.

• So an ionic solid is held together by many electrostatic attractions throughout the entire structure.

Ionic compounds do not usually exist as individual molecules.

• Therefore, we describe the simplest ratio of ions using a formula unit .

NaCl means:

• Na⁺ : Cl⁻ = 1 : 1

MgCl₂ means:

• Mg²⁺ : Cl⁻ = 1 : 2

Al₂O₃ means:

• Al³⁺ : O²⁻ = 2 : 3

A covalent bond forms when atoms share electrons.

Covalent bonding usually occurs between:

• nonmetal + nonmetal

• Example:

• H₂

Each hydrogen contributes one electron.

The two atoms share the pair:

• H—H

The shared electrons are attracted to both nuclei.

A single covalent bond contains:

• 1 shared electron pair

• represented by:

• —

A double bond contains:

• 2 shared electron pairs

• represented by:

• =

A triple bond contains:

• 3 shared electron pairs

• represented by:

• ≡

• Examples:

H₂:

• H—H

O₂:

• O=O

N₂:

• N≡N

Bond order tells you how many bonding pairs connect two atoms.

Single bond:

• bond order = 1

Double bond:

• bond order = 2

Triple bond:

• bond order = 3

Bond order strongly affects bond strength and length.

Generally:

• higher bond order → stronger bond

• and

• higher bond order → shorter bond

• So:

triple bond = shortest and strongest

• double = intermediate

• single = longest and weakest

• when comparing bonds between the same types of atoms.

Covalent electrons are not always shared equally.

If one atom attracts the shared electrons more strongly, the electrons spend more time near that atom.

• This produces a polar covalent bond .

The more electronegative atom gains a partial negative charge:

• δ−

The less electronegative atom gains a partial positive charge:

• δ+

The symbol δ means partial , not a full ionic charge.

Electronegativity measures an atom's attraction for shared electrons in a bond.

General trend:

• increases → across a period

• increases ↑ up a group

Fluorine has the greatest electronegativity.

When two atoms have very similar electronegativities:

• electrons are shared fairly equally.

• → mostly nonpolar covalent

As the difference increases:

• → more polar covalent

With a very large difference:

• → more ionic character

Khan Academy's Unit 2 practice specifically asks students to compare bonds based on their amount of ionic character and electronegativity differences.

Do not think there is a magical boundary where a bond suddenly becomes completely ionic.

• Instead, think of bonding as a spectrum:

• nonpolar covalent ← polar covalent → ionic

As electronegativity difference increases:

• ionic character generally increases

Compare:

• N—O

• C—F

• Mg—O

• Which would probably have the greatest ionic character?

Mg is a metal and O is a highly electronegative nonmetal.

The electronegativity difference is large.

• Therefore:

Mg—O has strong ionic character.

N and O have more similar electronegativities, so N—O is much more covalent.

Metal atoms also bond differently from typical covalent molecules.

In a metal:

• metal nuclei/core ions form an organized structure

valence electrons are relatively mobile

• electrons are delocalized instead of belonging to one specific bond

You may hear this described as:

• positive metal ions surrounded by a "sea" of mobile electrons

Because the valence electrons are mobile, they can move through the solid when an electric field is applied.

• Therefore metals are usually:

• good electrical conductors

Malleable means a material can be hammered or shaped without easily shattering.

Metal atoms can shift position while the delocalized electrons continue holding the structure together.

• Therefore metal layers can slide without immediately breaking the bonding structure.

Be able to distinguish:

And remember:

• larger electronegativity difference → greater ionic character