Chemical Bonding
Atoms bond to reach a lower energy state. How they bond — by transferring or sharing electrons — determines everything from a substance's melting point to how it reacts with acids.
Primary sources: RSC Periodic Table · NCBI Biochemistry — Water chapter
01 · Why atoms bond
Lower energy means greater stability.
Isolated atoms are not always in their lowest energy state. By sharing or transferring electrons, atoms can fill their outer electron shells and achieve a more stable configuration — usually matching the electron arrangement of a noble gas.
The octet rule states that most atoms are most stable with eight electrons in their outer shell. Hydrogen is an exception (it seeks two, matching helium). The rule is a useful approximation, not an absolute law — elements like sulfur and phosphorus can exceed eight.
Understanding why atoms bond requires knowing their electron configuration. See our atomic structure guide for shells, subshells and the Aufbau principle.
02 · Bond types
Four ways atoms interact.
The type of bond formed depends primarily on the atoms involved and the difference in their electronegativity.
03 · Polarity and molecular shape
Uneven sharing creates partial charges.
Electronegativity (Pauling scale) measures how strongly an atom attracts shared electrons. Fluorine is the most electronegative element (4.0); caesium among the least (0.79). When bonded atoms differ in electronegativity, the shared electrons sit closer to the more electronegative atom — creating a polar covalent bond.
A molecule's overall polarity depends on both its bond polarities and its shape. CO₂ has two polar bonds but is non-polar overall because its linear shape causes the bond dipoles to cancel. Water (H₂O) is polar because its bent shape means the dipoles do not cancel.
VSEPR theory predicts molecular shape by minimising repulsion between electron pairs. A central atom with four bonding pairs and no lone pairs is tetrahedral (109.5°). Each lone pair pushes bonding pairs closer together, reducing the bond angle.
Polarity drives solubility ("like dissolves like") and underpins acid–base chemistry — for more, see our acids and bases guide.
04 · Metallic bonding
A sea of delocalised electrons.
In a metal, atoms release their outer electrons into a shared "sea" of delocalised electrons that can move freely through the lattice. The positive metal ions are attracted to this sea, holding the lattice together.
This model explains the key properties of metals:
- Electrical conductivity — delocalised electrons carry charge.
- Thermal conductivity — electrons transfer kinetic energy rapidly.
- Malleability and ductility — layers of ions can slide without breaking bonds, because the electron sea moves with them.
- Lustre — electrons at the surface absorb and re-emit light of many frequencies.
Higher charge on metal ions and smaller ionic radii produce stronger metallic bonds. This is why magnesium has a higher melting point than sodium, and tungsten (used in light-bulb filaments) has the highest melting point of any metal.
Questions about chemical bonding
- Ionic bonds involve the complete transfer of one or more electrons from a metal to a non-metal, creating oppositely charged ions that attract. Covalent bonds involve the sharing of electron pairs between non-metal atoms. See the Royal Society of Chemistry periodic table for electronegativity data that predicts bond type.
- Electronegativity measures how strongly an atom attracts shared electrons. When two bonded atoms differ in electronegativity by roughly 0.4–1.7 (Pauling scale), the bond is polar covalent: the more electronegative atom carries a partial negative charge (δ−). Water (H₂O) is a classic example. See our acids and bases guide for how polarity drives water's behaviour as a solvent.
- Hydrogen bonds are not true chemical bonds — they are intermolecular forces. They are strong enough to give water its unusually high boiling point and to hold the two strands of DNA together. The NCBI Biochemistry chapter on water explains why hydrogen bonding is central to life.
- VSEPR works by counting all electron pairs (bonding and lone pairs) around the central atom. Lone pairs repel more strongly than bonding pairs, distorting ideal geometries. For example, water has a tetrahedral electron geometry but a bent molecular shape due to two lone pairs. See the Khan Academy VSEPR article.
Bonding drives acid–base chemistry
Polarity and hydrogen bonding explain why water acts as a solvent and how acids donate protons. Read the acids guide next.