Lewis structures Knowledge-Based Study

Study notes

1. Valence Electrons and the Octet Rule

   Lewis dot symbols represent the valence electrons of an atom as dots placed around the element's chemical symbol. Valence electrons are the outermost electrons involved in chemical bonding.

   Key principles include:

   • Valence electron count: Main-group elements have valence electron counts equal to their group number (in the 1–8 numbering scheme). For example, carbon has 4 valence electrons, nitrogen has 5, and oxygen has 6.
   • The Octet Rule: Nonmetal atoms tend to share electrons until they achieve a stable noble gas configuration, typically surrounded by 8 valence electrons (an octet). Hydrogen is an exception, requiring only 2 electrons to achieve stability (a duet).
   • Bonding vs. Nonbonding Pairs: Shared pairs of electrons between two nuclei form covalent bonds, while unshared pairs localized on a single atom are called lone pairs.

   Systematic Approach to Covalent Molecules

   Drawing an accurate Lewis structure requires a systematic accounting of all valence electrons:

   1. Count total valence electrons: Sum the valence electrons from all contributing atoms. For polyatomic cations, subtract one electron for each positive charge; for polyatomic anions, add one electron for each negative charge.
   2. Identify the central atom: Choose the least electronegative element (excluding hydrogen, which only forms one bond) as the central atom.
   3. Form skeletal single bonds: Connect each surrounding atom to the central atom using a single covalent bond. Each single bond accounts for 2 valence electrons.
   4. Satisfy terminal atom octets: Distribute the remaining electrons as lone pairs to satisfy the octets (or duet for hydrogen) of all terminal atoms first.
   5. Assign remaining electrons to the central atom: Place any leftover electrons as lone pairs on the central atom.
   6. Form multiple bonds if necessary: If the central atom lacks an octet after distributing all available electrons, convert one or more lone pairs from surrounding atoms into double or triple bonds.

   Evaluating Plausible Structures

   When multiple arrangements or bonding patterns are possible, formal charge helps determine the most stable Lewis structure.

   Formal charge is calculated for each atom using the following relation:

   Formal Charge = (Valence electrons of free atom) - (Nonbonding electrons) - (Number of covalent bonds)

   Guidelines for selecting the preferred Lewis structure:

   • A neutral molecule should have formal charges that sum to zero; an ion must have formal charges summing to the overall ionic charge.
   • Structures with formal charges closest to zero are generally preferred.
   • Negative formal charges should reside on the more electronegative elements.
   • Adjacent atoms bearing like formal charges indicate an unfavorable electron distribution.

   Important Octet Exceptions in Covalent Molecules

   While the octet rule applies broadly to second-row nonmetals like carbon, nitrogen, oxygen, and fluorine, three primary classes of exceptions exist:

   • Incomplete Octets: Elements such as beryllium and boron often form stable compounds with fewer than eight valence electrons around the central atom (for example, $BF_3$ has six valence electrons around boron).
   • Odd-Electron Molecules: Molecules containing an odd number of total valence electrons, such as nitric oxide ($NO$) and nitrogen dioxide ($NO_2$), cannot satisfy the octet rule for every atom and are classified as free radicals.
   • Expanded Valence Shells: Elements in Period 3 and beyond have available $d$ orbitals and can accommodate more than eight electrons around the central atom (for example, $PCl_5$ with 10 valence electrons and $SF_6$ with 12 valence electrons). Period 2 elements (such as C, N, O, and F) never expand their valence shell.