Double and triple bonds are appropriate in a Lewis structure when single bonds and lone pairs cannot give key atoms the right number of electrons (usually an octet) and reasonable formal charges. In practice, you “upgrade” single bonds to double or triple bonds to fix incomplete octets or bad formal charges on typical multiple‑bond formers like C, N, O, S, and P.

Core rule: fix incomplete octets

The main trigger for forming a double or triple bond is that the central atom still has less than eight electrons after you’ve placed all valence electrons as single bonds and lone pairs.

  • Draw a skeleton with single bonds from the central atom to surrounding atoms.
  • Distribute remaining electrons as lone pairs to give octets to terminal atoms first.
  • If the central atom still has fewer than eight electrons and no more electrons remain, convert a lone pair from a neighboring atom into a bonding pair, creating a double or triple bond.

This is explicitly taught as: if atoms still have incomplete octets but all electrons are used, “borrow” lone pairs from surrounding atoms to make double or triple bonds.

Which atoms commonly form multiple bonds?

Not every atom likes to share more than one pair of electrons. General guidance focuses on a few usual suspects.

  • Common multiple‑bond atoms: C, N, O, P, and S frequently form double or triple bonds in Lewis structures.
  • Poor multiple‑bond formers: terminal halogens such as F and Cl are normally kept as single‑bonded with three lone pairs.
  • Elements in the third period and beyond (like S, P, Cl) can have expanded valence shells, so they can participate in double bonds, but they do not require them just to reach stability.

So if something must form a double bond, look first at bonds between C, N, O, S, and P.

When double/triple bonds are appropriate

Once you have a draft Lewis structure, double and triple bonds become appropriate in a few recurring situations.

  • The central atom lacks an octet even though all valence electrons are placed.
  • The formal charges can be significantly improved by converting a lone pair on a neighboring atom into a bonding pair (for example, turning O−\text{O}^-O− and C+\text{C}^+C+ into neutral O=C).
  • The molecule or ion is known to contain multiple bonds (for instance CO2_22​, SO2_22​, NO3−_3^-3−​), and the octet and formal charge checks support that pattern.

Some practice sources phrase this as: create double or triple bonds by changing a lone pair into a bonding pair if the central atom still has less than eight valence electrons.

When double/triple bonds are not required

Learners are sometimes told rules like “any bond to oxygen is double” or “always use multiple bonds when the central atom has an expanded octet”; these overstatements are not generally correct.

  • Multiple bonds are not automatically required just because the central atom can expand its octet; such atoms can often satisfy bonding with only single bonds and extra lone pairs.
  • A lone pair on a terminal atom does not always have to be converted into a double bond; this is done only when needed for octets or formal charge correction.
  • Oxygen can form single bonds (e.g., in alcohols and peroxides), so a bond to oxygen is not automatically double.

The better rule is: improve octets and formal charges, not follow a memorized “always double‑bond oxygen” shortcut.

Practical step‑by‑step checklist

Many teaching resources combine these ideas into a repeatable decision process for when to introduce multiple bonds in Lewis structures.

  1. Count total valence electrons and adjust for charge.
  1. Draw the skeletal structure with single bonds from the central atom.
  1. Give terminal atoms octets with lone pairs; place any leftover electrons on the central atom.
  1. Check octets.
    • If the central atom (C, N, O, P, S, etc.) has fewer than eight electrons and no electrons remain, convert a neighboring terminal lone pair into a second (or third) bond.
  1. Check formal charges.
    • If large or awkward formal charges exist, see whether adding a double or triple bond between suitable atoms reduces them while preserving octets.

In forum‑style discussions, helpers often summarize it this way: once electrons are all placed, use double or triple bonds only when they are necessary to give central atoms full octets and sensible formal charges, primarily between atoms that commonly form multiple bonds (C, N, O, P, S).

Information gathered from public forums or data available on the internet and portrayed here.