Covalent Bonding, Lewis Structures & Formal Charge Why do atoms form covalent bonds? Non-metals tend to share electrons to reach a stable configuration (octet for most, duet for H). When two atoms share a pair of electrons, a covalent bond forms. Think of each shared pair as a handshake where both partners contribute one hand. Multiple handshakes between the same two atoms give double or triple bonds. The first bond formed along the internuclear axis is a strong, end-on overlap called a sigma (σ) bond; any extra bonds are side-on overlaps called pi (π) bonds. Single (H2), double (O2), and triple (N2) covalent bonds: electron sharing shown as overlapping clouds and matching Lewis dot structures. Bond formed by sharing one or more pairs of electrons between atoms (usually non-metals). Covalent bond One shared electron pair (1 σ bond). bond pair = 1 Single bond Two shared pairs (1 σ + 1 π). Double bond Three shared pairs (1 σ + 2 π). Triple bond End-on overlap along the internuclear axis; strongest first bond. Sigma (σ) bond Side-on overlap above/below the axis; additional bond(s) after σ. Pi (π) bond Covalent bond where both electrons come from the same atom (donor → acceptor). Coordinate (dative) bond Non-bonding pair of electrons localized on one atom. Lone pair A pair of electrons shared in a covalent bond. Bond pair Element symbol with dots showing valence electrons (outer-shell electrons). Lewis dot symbol Main-group atoms tend to achieve 8 electrons in their valence shell (duet for H). Octet rule Book-keeping charge on an atom in a Lewis structure assuming equal sharing of bonding electrons. Formal charge When a single Lewis structure cannot depict electron distribution; several valid forms exist. Resonance The real structure; weighted average of all significant resonance forms. Resonance hybrid Central atom accommodates more than 8 electrons (common for period-3 and beyond). Expanded octet Molecules/ions with an odd total electron count; at least one atom cannot have an octet. Odd-electron species Core terms (learn these now) Lewis dot symbols and covalent bonds Lewis showed that valence electrons can be represented as dots around the symbol: H(1), C(4), N(5), O(6), F(7), Ne(8). Atoms with 1–3 dots tend to share/lose; 5–7 tend to gain/share. Examples with SMILES to visualize connectivity: H2 ([H][H]), F2 (F-F), O2 (O=O), N2 (N N), HCl ([H]Cl), CH4 (C), H2O (O), NH3 (N), CO2 (O=C=O), C6H6 (benzene; c1ccccc1). How to draw a Lewis structure (always do this sequence) Algorithm Write the skeleton: central atom is usually the least electronegative (H is never central). Count total valence electrons (add for anions, subtract for cations). Place single bonds (2 e− per bond) between the central atom and surrounding atoms. Complete octets on outer atoms with lone pairs. Put remaining electrons on the central atom as lone pairs. If the central atom lacks an octet, convert an adjacent lone pair into a multiple bond (double/triple) as needed. For ions, enclose the structure in square brackets with the charge superscript. Finally, compute formal charges and choose the structure with minimum formal charge and negative charge on more electronegative atoms. Lewis-structure workflow for CO2 and SO4 2− : count electrons → connect with single bonds → complete octets → add multiple bonds if needed → check formal charges. H2O H–O–H; O has 2 lone pairs NH3 H–N(H)–H; N has 1 lone pair CH4 C with four H single bonds CO2 O=C=O (linear connectivity) NO3− [O–N(=O)–O]−; 3 equivalent resonance forms SO4 2− [O–S(=O)2–O] 2−; resonance with two S=O and two S–O− PCl5 P with five P–Cl single bonds (expanded octet) SF6 S with six S–F single bonds (expanded octet) XeF4 Xe with four Xe–F bonds; Xe also has 2 lone pairs (expanded octet) Lewis structures: NEET-standard species at a glance Species Formula Lewis dot sketch (text) bonds around central atom lone pairs on central atom Central atom selection hack: Often the atom that appears once (e.g., C in HCN) or the least electronegative (except H) sits in the center. tip Formal charge: quick, safe structure checking Formal Charge is computed using this expression for each atom. Key relation Formal charge helps you compare valid Lewis structures. Lower magnitude formal charges are preferred. Negative charges should reside on more electronegative atoms (like O, F). You may also see the same relation written with L for lone-pair electrons: FC = V − L − B/2 (here L = N). Example: nitrate (NO3−), 24 e− total Propose one form: O=N–O with another O− (i.e., one N=O and two N–O−). Formal charges: N (V=5, N=0, B=8) → 5−0−4 = +1; doubly bonded O (V=6, N=4, B=4) → 6−4−2 = 0; singly bonded O− (V=6, N=6, B=2) → 6−6−1 = −1. All three positions are equivalent by resonance, so the actual hybrid has three equal N–O bonds with bond order 1.33. Common molecules vs expanded-octet species: CH4, H2O, NH3 obey octet; PCl5 shows expanded octet. Inset illustrates a sample formal-charge calculation. Coordinate (dative) bond — both electrons from the donor Sometimes, one atom donates a lone pair to an electron-deficient atom or ion. The bond formed is still covalent, but the donation is one-sided initially. Examples you must know: NH3 + H+ → NH4+ (ammonium; N donates its lone pair to H+) and BF3·NH3 adduct (NH3 donates to electron-deficient BF3). We show this with an arrow: NH3 → BF3. Ammonia donates its lone pair to a proton to form NH4+. lone-pair donor (Lewis base) azane ammonia proton hydrogen cation acceptor (Lewis acid) product ammonium ion ammonium Formation of ammonium, NH4+ (coordinate bond step) H+ approaches the lone pair on N of NH3. The lone pair from N forms a new N–H bond (coordinate bond). In NH4+, all four N–H bonds are equivalent; the bond formed by donation becomes indistinguishable from the others. Electron-deficient B (6 e−) in BF3 attracts a donor. NH3 donates its lone pair to B, forming a B←N coordinate bond. B now attains an octet in the adduct. Adduct formation: BF3·NH3 Lewis acid (acceptor) boron trifluoride trifluoroborane Lewis base (donor) azane ammonia Lewis adduct BF3·NH3 ammonia–boron trifluoride adduct Electron-deficient BF3 accepts a lone pair from NH3 to complete B’s octet. Coordinate bond diagram: two panels on white. Left: NH3 donates lone pair to H+ to form NH4+; show electron-pair arrow from N to H+. Right: BF3 accepts lone pair from NH3 (B←N). Label donor/acceptor, lone pairs, and final ions/adduct. Clean vector chemistry style, red curved arrows, no internal captions. 2026-05-26T17:04:17.741Z gpt-image-2 Coordinate bonding illustrated: NH3 + H+ → NH4+ and NH3 → BF3 to form BF3·NH3. Arrow shows donation from donor to acceptor. Resonance: when one Lewis structure isn’t enough If multiple valid Lewis structures differ only in placement of electrons (not atom positions), the actual molecule is a resonance hybrid. Classic cases: CO3 2− , NO3−, NO2−, SO2, SO3, O3, and benzene (C6H6). Resonance equalizes bond lengths and spreads charge, often stabilizing the species. Resonance in CO3 2− : three equivalent structures with the double bond at different O positions; the hybrid shows three equal C–O bonds (bond order 1.33). gpt-image-2 2026-05-26T17:04:17.687Z Three-panel carbonate resonance: draw O=C–O− with arrows to its two equivalents (double bond rotates). Below, a dashed-bond resonance hybrid circle indicating all C–O bonds equal. Include formal charges on O and C. Vector, black atoms, red arrows. CO3 2− 3 equivalent forms with one C=O and two C–O− each 1.33 for each C–O NO3− 3 equivalent forms with one N=O and two N–O− 1.33 for each N–O NO2− 2 equivalent forms with one N=O and one N–O− 1.5 for each N–O SO2 2 forms with S=O at alternate O; charge-separated ≈1.5 for each S–O SO3 3 forms; S=O at alternate positions; charge-separated 1.33 for each S–O O3 2 forms with O=O at alternate positions 1.5 for each O–O C6H6 (benzene) 2 Kekulé forms; delocalized π ring 1.5 for each C–C Species Resonance structures (description) Hybrid bond order (typical) Resonance and average bond order Resonance structures flip back and forth in time. They do not alternate. The molecule is a resonance hybrid — a single, blended structure with delocalized electrons. Octet rule — and its three classes of exceptions Octet is a guiding rule, not a law. Exceptions fall into three clear buckets: incomplete octet (electron-deficient centers like Be and B), expanded octet (hypervalent centers in period 3 or beyond like P, S, Xe), and odd-electron species (NO, NO2, ClO2). Know examples cold for NEET. Category Species Electrons around central atom Why exception / remark Octet exceptions you must recognize Incomplete octet BeCl2 Be has only two bonds (4 e−) — electron-deficient Incomplete octet BF3 B has three bonds (6 e−) — strong Lewis acid Incomplete octet LiCl 2 (if viewed covalently as Li–Cl) Mainly ionic in solid; octet concept for covalent Lewis forms is not directly applicable — included as a textbook electron-deficient illustration Expanded octet PF5 10 Hypervalent P (period 3), more than 8 e− around P Expanded octet SF6 12 Hypervalent S with six S–F bonds Expanded octet PCl5 10 Hypervalent P Expanded octet IF7 14 Hypervalent I (period 5) Expanded octet XeF2 10 Hypervalent Xe (noble-gas compound) Expanded octet XeF4 12 Hypervalent Xe; also has 2 lone pairs Expanded octet XeF6 14 Hypervalent Xe Odd-electron NO Odd total (11 e−) — one unpaired electron Cannot give all atoms an octet Odd-electron NO2 Odd total (17 e−) — one unpaired electron Central N lacks full octet Odd-electron ClO2 Odd total (19 e−) — one unpaired electron Octet not possible for all atoms gpt-image-2 Expanded-octet examples: PF5 (10 e− around P) and SF6 (12 e− around S) drawn as Lewis structures with bond pairs only — no shapes discussed here. 2026-05-26T17:04:17.892Z Side-by-side Lewis structures for PF5 and SF6: central P with five single P–F bonds; central S with six S–F bonds. Indicate electron counts (10e− and 12e−) near the central atom. Clean vector style, neutral palette, no 3D geometry. All atoms in a molecule must always satisfy the octet rule. Hydrogen has a duet; Be and B often have incomplete octets; many third-period and heavier elements (P, S, Xe, I) can have expanded octets. Odd-electron species (NO, NO2, ClO2) also break the octet. Lewis structures show connectivity and electron pairs. Actual shapes require VSEPR analysis (next unit). Lewis structures alone determine the precise 3D geometry of a molecule. A double bond is exactly twice as strong as a single bond. A double bond has one σ and one π bond. π bonds are weaker than σ bonds, so a double bond is stronger than a single but not double its strength. Water’s (H2O) covalent bonds make it a superb solvent and buffer for body temperature. Most biomolecules and drugs are covalently bonded — their functions depend on how electrons are shared. remember neet-alert Common traps: (1) Forgetting to adjust electron count for ionic charges, (2) Putting H at the center, (3) Not checking formal charges, (4) Missing octet exceptions like BF3, SF6, NO, and noble-gas fluorides (XeF2/4/6). Here L denotes lone-pair (non-bonding) electrons; L = N. Alternative notation (same meaning)