Chemical bonds form when atoms interact in ways that lower their total energy - almost always by achieving a more stable electron configuration. The three main bond types are ionic (electron transfer between metals and nonmetals), covalent (electron sharing between nonmetals), and metallic (electrons delocalized across a sea of metal cations). Each type produces distinct physical and chemical properties.
Electronegativity difference determines bond type: a large difference (ΔEN > 1.7) produces ionic character; a small difference (ΔEN < 0.5) means nonpolar covalent; intermediate values produce polar covalent bonds. Real bonds exist on a spectrum from pure covalent to pure ionic, not in discrete categories.
Understanding Chemical Bonding
Bonding is best understood by tracking valence electrons and the forces between charged particles. In a sodium chloride crystal, each sodium ion attracts several chloride ions around it, rather than forming one separate sodium chloride pair. This repeating arrangement is called a lattice.
It explains why a crystal resists being pulled apart. It also explains brittleness. When layers in the lattice are pushed out of position, ions with the same charge can end up close together.
Their strong repulsion splits the crystal. A solid ionic compound cannot usually carry electric current because its ions are locked in place. When it melts or dissolves in water, the ions can move and carry charge.
A covalent bond is only part of the story for a molecular substance. The shape of a whole molecule decides whether its bond polarities cancel or combine. Carbon dioxide has polar carbon oxygen bonds, but its straight, symmetrical shape makes the molecule overall nonpolar.
Water has polar oxygen hydrogen bonds and a bent shape, so it has one more negative side near oxygen and more positive areas near hydrogen. Water molecules attract one another strongly through hydrogen bonding. This helps explain water’s unusually high boiling point for such a small molecule.
It is important to separate bonds within a molecule from forces between molecules. Breaking covalent bonds changes the substance chemically. Separating molecules during melting or boiling usually overcomes weaker intermolecular forces.
Metal properties come from the way metal atoms are packed in layers. Their outer electrons can move through the solid, while the positive metal ions remain in a regular arrangement. A voltage can direct these mobile electrons, producing an electric current.
The same electrons transfer thermal energy quickly, which is why many metals feel cold at room temperature. When a metal is bent or hammered, its ion layers can slide past one another without creating the strong same charge repulsions that shatter an ionic crystal. Alloys are mixtures of metals, sometimes with small amounts of other elements.
Different sized atoms disturb the layers and make sliding harder. This is why steel is generally harder than pure iron.
When drawing Lewis structures, begin by counting all available valence electrons. Make a simple skeleton, place single bonds, then complete outer atoms before placing remaining electrons on the central atom. Check that the total number of dots and bonding electrons matches the starting count.
Formal charge can help choose between possible structures because arrangements with smaller charges are often more reasonable. The octet rule is a useful starting pattern, not an unbreakable law. Hydrogen needs only two electrons.
Boron can have fewer than eight, while elements in the third period or below can sometimes have more. After drawing the electron arrangement, use VSEPR to predict the three dimensional shape.
Lone pairs take up more space than bonding pairs, so they often compress bond angles. Careful diagrams link electron arrangement, shape, polarity, and observable properties.
Key Facts
- Ionic bonds form between metals and nonmetals; electrons are transferred. Produce crystalline solids with high melting points.
- Covalent bonds form between nonmetals; electrons are shared. Produce molecules with lower melting points.
- Metallic bonds form between metal atoms; electrons are delocalized. Produce conductive, malleable solids.
- Polar covalent: unequal sharing (one atom pulls harder). Creates a partial negative charge (δ-) on the more electronegative atom.
- Lewis structures show valence electrons as dots around atoms and bond pairs as lines.
- VSEPR theory predicts molecular shape based on repulsion of electron pairs.
Vocabulary
- Ionic bond
- A bond formed by the electrostatic attraction between oppositely charged ions (one atom transfers electrons to another).
- Covalent bond
- A bond formed by the sharing of electron pairs between atoms.
- Metallic bond
- A bond in metals formed by a 'sea' of delocalized electrons surrounding positively charged metal cations.
- Electronegativity
- Measure of an atom's ability to attract bonding electrons; drives polarity.
- Polar covalent bond
- A covalent bond in which electrons are shared unequally, creating partial charges on each atom.
Common Mistakes to Avoid
- Thinking ionic compounds are always soluble in water. Many ionic compounds are insoluble - solubility depends on specific lattice and hydration energies.
- Drawing Lewis structures with electrons for inner-shell electrons. Only valence electrons are shown in Lewis structures.
- Assuming all covalent compounds don't conduct electricity. Some polar molecules in solution ionize and conduct. The rule is for pure covalent molecular substances.
- Confusing bond polarity with molecular polarity. A molecule can have polar bonds but be nonpolar overall if the dipoles cancel by symmetry (e.g. CO₂ is linear and nonpolar).
Practice Questions
- 1 Predict the type of bond (ionic, polar covalent, or nonpolar covalent) between: Na and Cl, H and O, C and H.
- 2 Draw the Lewis structure for H₂O. How many bonding pairs and lone pairs does the oxygen atom have?
- 3 Explain why NaCl has a much higher melting point than H₂O using bond type reasoning.