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Concept 01 / 07Chemical Bonding and Molecular Structure

Kossel-Lewis Approach: Octet Rule, Lewis Structures and Formal Charge

Valence electrons and Lewis symbols, ionic and covalent bond formation by the octet rule, drawing Lewis structures step by step, formal charge and choosing the best structure, coordinate bonds, and the limitations of the octet rule (incomplete octet, odd-electron molecules, expanded octet).

Figure 1
Kossel and Lewis (1916) proposed that atoms combine so as to attain the stable eight-electron configuration of a noble gas (the octet rule), either by transferring electrons (ionic bonds, Kossel) or by sharing them (covalent bonds, Lewis). Hydrogen, lithium and beryllium aim for the two-electron duet of helium. Lewis symbols show valence electrons as dots around the element symbol; the number of valence electrons equals the group number for groups 11 and 22 and the group number minus 1010 for groups 1313 to 1818. Drawing a Lewis structure. 1. Count the total valence electrons; add one per negative charge and subtract one per positive charge. 2. Choose the skeleton: usually the least electronegative atom (not H) is central; H and F are always terminal. 3. Join atoms with single bonds (two electrons each). 4. Complete octets on the outer atoms with lone pairs, then place any remaining electrons on the central atom. 5. If the central atom lacks an octet, convert lone pairs on neighbours into multiple bonds. For CO2\mathrm{CO_{2}}: 4+2×6=164 + 2 \times 6 = 16 electrons, giving O=C=O with two lone pairs on each oxygen (figure). Formal charge. To decide between possible structures and locate charges, assign each atom a formal charge: FC=(valence electrons)−(non-bonding electrons)−12(bonding electrons)\text{FC} = (\text{valence electrons}) - (\text{non-bonding electrons}) - \frac{1}{2}(\text{bonding electrons}). The formal charges add up to the overall charge. The preferred structure has the smallest formal charges, with any negative formal charge on the more electronegative atom. In the nitrate ion (figure), N has +1+1, the double-bonded O has 00, and each single-bonded O has −1-1: total −1-1. Formal charge is a bookkeeping device, not the actual charge distribution. Coordinate (dative) bond: both shared electrons come from one atom, as in NH4+\mathrm{NH_{4}^{+}} (from the lone pair of NH3\mathrm{NH_{3}} to H+\mathrm{H^{+}}) or H3N→BF3\mathrm{H_{3}N \rightarrow BF_{3}}. Once formed, it is indistinguishable from other covalent bonds (all four N-H bonds of NH4+\mathrm{NH_{4}^{+}} are identical). Limitations of the octet rule. Incomplete octet: BeCl2\mathrm{BeCl_{2}} (4 electrons around Be), BF3\mathrm{BF_{3}}, AlCl3\mathrm{AlCl_{3}} (6 around B or Al). These are electron-deficient and act as Lewis acids. Odd-electron molecules: NO, NO2\mathrm{NO_{2}}, ClO2\mathrm{ClO_{2}} cannot give every atom an octet; they are paramagnetic free radicals. Expanded octet: third-period and heavier elements can hold more than eight electrons (PCl5\mathrm{PCl_{5}}: 1010; SF6\mathrm{SF_{6}}: 1212; IF7\mathrm{IF_{7}}: 1414), traditionally attributed to vacant d orbitals. The rule also says nothing about molecular shape or bond energy, and it cannot explain the existence of noble gas compounds such as XeF2\mathrm{XeF_{2}}, XeF4\mathrm{XeF_{4}}. Key Takeaways: - Total valence electrons: add for negative charge, subtract for positive charge. - Least electronegative atom central; H and F terminal. - FC=V−N−B2\text{FC} = V - N - \frac{B}{2}; minimise formal charges; negative FC on the more electronegative atom. - Coordinate bonds become equivalent to ordinary covalent bonds once formed. - Exceptions: incomplete octet (B, Be, Al), odd electrons (NO, NO2\mathrm{NO_{2}}), expanded octet (P, S, Cl, Xe). - Pitfall: forgetting to add the ionic charge when counting electrons for an ion. Brain-Teaser 1: Draw the Lewis structure of SO42−\mathrm{SO_{4}^{2-}} minimising formal charges. Solution: Electrons: 6+4×6+2=326 + 4 \times 6 + 2 = 32. With four S-O single bonds and three lone pairs on each O, S has formal charge +2+2 and each O −1-1. Converting two lone pairs into S=O double bonds (S expands its octet to 1212) gives S 00, two double-bonded O 00 and two single-bonded O −1-1: total −2-2, with smaller formal charges. (Modern analyses suggest the real bonding is closer to the charge-separated picture with ionic contributions, which is one reason formal charge should be used as a guide, not as physical truth.) Brain-Teaser 2: Which structure of N2O\mathrm{N_{2}O} is better: N=N=O or N=O=N? Solution: 1616 valence electrons. For N=N=O (with two lone pairs on each end atom): formal charges are −1-1 on terminal N, +1+1 on central N, 00 on O. For N=O=N: terminal N atoms −1-1 each, central O +2+2. The first has smaller formal charges and places no positive charge on the most electronegative atom, so it is preferred, and experiment confirms the N-N-O connectivity. Brain-Teaser 3: Identify which of these violate the octet rule and how: BCl3\mathrm{BCl_{3}}, ClO2\mathrm{ClO_{2}}, XeF4\mathrm{XeF_{4}}, CO\mathrm{CO}. Solution: BCl3\mathrm{BCl_{3}}: incomplete octet (6 electrons on B). ClO2\mathrm{ClO_{2}}: odd electron count (1919), so one atom has an unpaired electron. XeF4\mathrm{XeF_{4}}: expanded octet (1212 electrons on Xe: four bonds and two lone pairs). CO\mathrm{CO}: obeys the octet rule with a triple bond (: ⁣C≡O ⁣::\!\mathrm{C}\equiv\mathrm{O}\!:), though with formal charges −1-1 on C and +1+1 on O, which explains its small dipole moment and why it binds metals through carbon.

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Kossel-Lewis Approach: Octet Rule, Lewis Structures and Formal Charge — DPP

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