Types of Chemical Bonds

Full Notes

Chemical bonds form when atoms interact through their valence electrons. The nature of a bond – ionic, polar covalent, nonpolar covalent, or metallic – depends on the electronegativities of the atoms involved and their positions in the periodic table.

Recap - Electronegativity and Periodic Trends

Electronegativity is a measure of how strongly an atom attracts electrons in a bond.

• Increases left to right across a period (due to increased nuclear charge)
• Decreases down a group (due to increased distance and shielding)
• Fluorine is the most electronegative element; cesium and francium are among the least.

These trends can be explained using:

• Coulomb’s law – attraction between nucleus and electrons
• Shell model – more shells = more shielding = lower attraction to shared electron pair
• Effective nuclear charge (Z_eff) – higher Z_eff = stronger pull on electrons

What is ionic bonding?

Ionic bonding is the electrostatic attraction between oppositely charged ions. Usually, ionic bonding occurs between metals (which form positively charged ions) and non-metals (which usually formed negatively charged ions).

What is a covalent bond?

Covalent bonding is the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons.

• Each atom shares one or more electrons to achieve a full outer shell.
• The shared electrons are electrostatically attracted to both nuclei, creating a strong bond.

Sometimes, two atoms will share more than one pair of electrons, forming double and triple bonds.

• Double bonds = two shared pairs of electrons (e.g. O=O, CO₂)
• Triple bonds = three shared pairs of electrons (e.g. N≡N)

These bonds are stronger and shorter than single bonds.

Types of Covalent Bonds

Nonpolar covalent bonds

• Atoms have the same or similar electronegativities
• Electrons in the bond are shared equally between the two atoms

Polar covalent bonds

A polar covalent bond occurs when there is a difference in electronegativity between the two bonded atoms. The bonding electrons are pulled closer to the more electronegative atom, making the electrons unevenly shared.

This creates partial charges (δ⁺ and δ⁻) at either end of the bond:

• The more electronegative atom becomes δ⁻
• The less electronegative atom becomes δ⁺

• Chlorine (Cl) is more electronegative than Hydrogen (H).
• The bonding electrons are pulled closer to Cl, giving it a partial negative charge (δ⁻).
• H loses electron density, giving it a partial positive charge (δ⁺).

The Bonding Continuum: Ionic to Covalent

Bonding continuum

Ionic and covalent bonding are the extremes of a continuum of bonding type. In reality, most bonds exist between the two extremes - the greater the difference in electronegativity, the more ionic character the bond has and the smaller the difference in electronegativity, the more covalent character the bond has.

Example:

• NaCl: large difference in Pauling Electronegativity values → ionic
• HCl: moderate difference → polar covalent
• Cl₂: no difference → non-polar covalent

All polar covalent bonds have some ionic character, and even ionic bonds have covalent aspects. Bond type should be viewed as a sliding scale, not an either/or.

Determining Bond Type in Practice

While electronegativity differences give a good starting point, the best way to classify bonding is to examine the properties of the compound:

• Ionic compounds:
  • High melting points
  • Solid at room temperature
  • Conduct electricity when dissolved in water or molten
  • Crystalline solids
• Covalent compounds:
  • Lower melting points
  • Often liquids or gases
  • Do not conduct electricity in solution
  • Molecules, not ions

Metallic Bonding

In metallic solids, valence electrons are delocalized, meaning they are not bound to any one atom.

• Electrons move freely through the structure (“sea of electrons”)
• This explains key properties of metals:
  • Electrical conductivity
  • Malleability and ductility
  • Lustrous appearance

Example: Structure of Sodium (Na)

• Each sodium atom loses one outer electron, forming Na⁺ ions.
• The lost electrons become delocalised, forming an electron cloud.

There is strong attraction between Na⁺ ions and the delocalised electrons, which holds the metal together.

Matt’s exam tip

Don’t just memorize bond types — understand the cause. Look at the periodic table and compare electronegativity values. Also, practice identifying which atom will carry a partial charge, and consider how this affects bond polarity and interactions.

Summary

• The type of chemical bond depends on the electronegativity of the atoms involved and their position in the periodic table.
• Bonds range from nonpolar covalent to ionic, with polar covalent bonds in between.
• Electronegativity trends explain this behavior, along with key concepts like Coulomb’s law, effective nuclear charge, and electron shielding.
• Metals exhibit metallic bonding, where electrons are delocalized.
• Key points to remember:
  • Electronegativity increases across a period, decreases down a group
  • Nonpolar covalent = similar EN; Polar covalent = moderate difference; Ionic = large difference
  • Polar bonds have dipoles and partial charges
  • Bonding is a spectrum — not a set of fixed categories
  • Properties (melting point, conductivity) help identify bond type