Skip to main content
ScienceVerse
ChemistryDifficulty 2-4

Chemical Bonds

A chemical bond is the attraction that holds atoms together in a substance. In an ionic bond, oppositely charged ions attract after electrons have moved from one atom to another; in a covalent bond, two atoms share a pair of electrons; in a metal, outer electrons are shared across the whole structure. Bonding lowers a system's energy, so breaking a bond always costs energy and forming one releases it.

Sign in to save this concept.

Potential energy and bond formation

As two hydrogen atoms approach, attraction between each nucleus and the shared electrons lowers the potential energy until nuclear repulsion takes over; the bond length is the separation at the energy minimum. Separating the atoms completely again takes the bond energy: 436 kJ mol⁻¹ for H–H in the gas phase.

Electronegativity as a derived scale

Pauling derived electronegativity from bond dissociation energies, producing a dimensionless relative scale; it is a calculated, not directly measured, property. Values rise across a period and fall down a group.

Common misconception: Ionic and covalent bonding are idealised limits. Real bonds have mixed character, which is why Δχ thresholds misclassify compounds such as MnI₂.
Full explanation — the complete reference version every reading depth is based on

Three ways atoms hold together

  • Ionic: a metal atom loses electrons to a non-metal atom; the resulting positive and negative ions attract electrostatically (for example sodium chloride).
  • Covalent: two atoms, usually non-metals, share a pair of electrons that both nuclei attract (for example the H–H bond in hydrogen gas).
  • Metallic: outer electrons are delocalised over many metal atoms as a 'sea of electrons', which is why metals conduct electricity and can be hammered into shape.

Sharing fairly or unfairly: electronegativity

Electronegativity measures how strongly an atom pulls on the electrons in a bond. When two identical atoms bond, they share equally (a non-polar covalent bond). When the atoms differ, the electrons spend more time near the more electronegative atom, which gains a partial negative charge (δ−) while its partner becomes partially positive (δ+). On the Pauling scale, fluorine is 3.98, chlorine 3.16 and sodium 0.93.

Δχ=∣χA−χB∣\Delta\chi = |\chi_A - \chi_B|

Electronegativity difference between bonded atoms A and B: small → non-polar, moderate → polar covalent, large → ionic.

Worked example (ScienceVerse calculation): for sodium and chlorine, Δχ = 3.16 − 0.93 = 2.23 — a large difference, consistent with sodium chloride being ionic. For an O–H bond in water, Δχ = 3.44 − 2.20 = 1.24 — a polar covalent bond with oxygen partially negative. For two chlorine atoms in Cl₂, Δχ = 0, so the bond is non-polar covalent.

Common misconception: It is tempting to treat 'ionic' and 'covalent' as two separate boxes with a sharp cut-off. Bonding is a continuum: hydrogen fluoride (Δχ = 3.98 − 2.20 = 1.78) is polar covalent, yet manganese(II) iodide (Δχ = 2.66 − 1.55 = 1.11) is ionic. The difference is a guide, not a rule.

Bonds and energy

A covalent bond forms because the attraction of both nuclei for the shared electrons outweighs the repulsion between the nuclei, lowering the energy. So breaking a bond always needs energy: separating one mole of hydrogen molecules into atoms takes 436 kJ, and the same amount is released when the bonds form.

Common misconception: Many people think breaking bonds releases energy (for example when fuel burns). Breaking a bond always absorbs energy; a reaction releases energy overall only when the new bonds formed are stronger than the old ones broken.

How we know, and why it matters

In the 1930s Linus Pauling was among the pioneers who used quantum mechanics to describe chemical bonding, and he received the 1954 Nobel Prize in Chemistry for his research into the nature of the chemical bond; he developed electronegativity values by comparing the energies needed to break different bonds. Bond type explains everyday properties: ionic solids such as salt are hard, brittle and high-melting, and conduct electricity only once melted or dissolved, because only then can their ions move.

Ask ScienceVerse

Still curious about Chemical Bonds? Ask a question, get hints, take a short lesson or try a challenge. The tutor answers only from this concept's approved sources, and says so when it has none.

Ask the tutor about this concept on the full tutor page.

Connections

Prerequisites

Understand these first:

Guided learning path

See everything to learn before this, in order, with your progress:

Leads to

This concept is a building block for:

Related concepts

Try the experiment

Put this concept into practice with a hands-on activity (each shows its supervision requirement first):

Check your understanding

Take a quick check of two to five questions, with an explanation for every answer:

See the neighbourhood of Chemical Bonds in the Knowledge Galaxy

Sources and methodology

Claims marked “awaiting scientific review” cite the sources listed but have not yet been signed off by a scientific reviewer.

Content status: published 1 October 2026.

  • Scientific review: this version has not yet been signed off by a scientific reviewer.
  • The Advanced explanation has not yet been reviewed for age suitability.