Q.Which of the following is electron deficient?

The octet rule -- every atom ends up with eight valence electrons -- works cleanly for many second-period-centred molecules, but it has three recognised families of exception.

1. Electron-deficient central atoms. Some central atoms end up with FEWER than eight electrons around them because they simply do not have enough valence electrons (or enough bonding partners) to reach an octet. Boron trifluoride, BF₃, is the classic case: boron has only three valence electrons, so even after forming three B-F single bonds and giving each fluorine a complete octet, boron itself sits with just six electrons. A tempting fix -- moving a fluorine lone pair to form a B=F double bond -- is actually LESS favourable, because it forces a positive formal charge onto fluorine, the most electronegative atom present; so BF₃ genuinely exists with boron's octet incomplete. BCl₃ and BeCl₂ behave the same way.

2. Odd-electron molecules. A handful of molecules have a central atom contributing an ODD total number of valence electrons, so no Lewis structure can pair every single electron — one electron is inevitably left unpaired. Nitric oxide (NO) and nitrogen dioxide (NO₂) are the standard examples.

3. Expanded valence shells. Central atoms from period 3 onward have empty, energetically-accessible d orbitals available, and can use them to accommodate MORE than eight electrons. Sulphur hexafluoride (SF₆, sulphur surrounded by twelve electrons across six bonds) and phosphorus pentachloride (PCl₅, phosphorus surrounded by ten electrons across five bonds) are the textbook cases; the hybridisation that makes this possible (sp³d, sp³d²) is covered under Hybridisation of Atomic Orbitals.

Recognising which family a molecule falls into -- and, for family 1, correctly picking the LOWER-formal-charge (incomplete-octet) structure over the tempting-but-wrong completed-octet alternative -- is a frequently tested skill.