Bonding
Chemical Bonds
Forces holding atoms together — ionic, covalent, metallic, and beyond
A chemical bond is an attractive force holding atoms together in molecules or compounds. Three main types: (1) Ionic — electrostatic attraction between cations and anions (NaCl). (2) Covalent — sharing of electron pairs between atoms (H₂O, CO₂). (3) Metallic — electron sea among positive ions (metals). Bond character is a continuum based on electronegativity difference. Strong bonds (>200 kJ/mol). Weaker intermolecular forces (hydrogen bonds, van der Waals) hold molecules to each other but aren't true chemical bonds.
- Ionic bondElectrostatic attraction; ΔEN > 1.7
- Covalent bondShared electron pair; ΔEN < 1.7
- Metallic bondElectron sea around cations
- Bond energy range~150-1000 kJ/mol (single to triple)
- Triple bond~3× single bond energy (e.g., N₂)
- Bond lengthShorter = stronger; ~1-2 Å typical
Interactive visualization
Press play, or step through manually. The visualization is yours to drive — try it before reading on.
Watch the 60-second explainer
A condensed visual walkthrough — narrated, captioned, under a minute.
Why bonds matter
- Molecular structure. Bond geometry defines shape.
- Reactivity. Weak bonds break first.
- Material properties. Hardness, conductivity.
- Thermodynamics. Bond energies drive reactions.
- Drug design. Drug-receptor binding.
- Biochemistry. Proteins, DNA held by bonds.
- Energy storage. Bond energy = chemical fuel.
Common misconceptions
- Bonds are physical objects. Force, not material.
- Ionic vs covalent are absolute. Continuum of bond character.
- All bonds break at same temperature. Energy varies widely.
- Octet rule is universal. Many exceptions (H, Be, B, expanded octets).
- Hydrogen bonds are chemical bonds. Intermolecular force; weaker.
- Bond breaking releases energy. Bond breaking REQUIRES energy.
Frequently asked questions
What's the difference between ionic and covalent?
Ionic: complete electron transfer from one atom to another, creating positive and negative ions held by electrostatic force. Common between metal + nonmetal (NaCl). Covalent: electrons shared between atoms in a bond. Common between nonmetals (H₂O, CO₂). Determined by electronegativity difference (ΔEN). Pure ionic > 1.7; pure covalent < 0.5; polar covalent in between.
How does the octet rule work?
Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (full s and p subshells), like noble gases. Drives much of bonding. Examples: Na (1 valence) loses one → Na⁺ (octet). Cl (7 valence) gains one → Cl⁻ (octet). Carbon (4) shares 4 → octet via covalent bonding. Exceptions: H (2 electrons), Be (4), B (6), period 3+ elements (expanded octet — e.g., SF₆).
What's a metallic bond?
Metal atoms in a lattice donate their valence electrons to a "sea" surrounding all atoms. Positive ions held in place by electrostatic attraction to mobile electrons. Explains: conductivity (mobile electrons), ductility (electrons reorganize to keep bonding), shine (electron interaction with light). Different from covalent (localized) and ionic (rigid).
How strong are bonds?
Bond energy = energy to break bond. C-H: 413 kJ/mol. C-C single: 348. C=C double: 614. C≡C triple: 839. N≡N: 945 (very strong; why N₂ inert). H-F: 568 (strongest single bond involving H). Bond strength depends on bond order, electronegativity, atom size. Stronger bonds → harder to break → less reactive.
What determines bond polarity?
Electronegativity difference. ΔEN ≈ 0: nonpolar (e.g., H-H, C-H). ΔEN 0.5-1.7: polar covalent (e.g., O-H, ΔEN = 1.4). ΔEN > 1.7: ionic (e.g., Na-Cl, ΔEN = 2.1). Polar bonds make polar molecules (H₂O) — affects boiling point, solubility. Symmetry can cancel polarity: CO₂ has polar bonds but linear → nonpolar overall.
Can a single bond be longer than a double?
No — double bonds are shorter and stronger. C-C: 154 pm. C=C: 134 pm. C≡C: 120 pm. More electron sharing → atoms pulled closer. Bond strength inversely related to length. Quick rule: triple < double < single bond length; triple > double > single bond strength.
What about coordinate bonds?
Special covalent bond where both shared electrons come from one atom (vs each atom contributing one). Example: NH₃ + H⁺ → NH₄⁺. Nitrogen donates lone pair to form fourth N-H bond. Once formed, indistinguishable from regular covalent bond. Common in coordination chemistry (transition metal complexes).