1. Writing Lewis Formulas Chem 1061 Updated November 2009 ©Lance S. Lund Anoka-Ramsey Community College

2. Steps in Writing Lewis-dot Structures Calculate the total number of valence electrons. Write the skeleton structure, connecting every bonded pair of atoms with a pair of dots or a dash. Distribute electrons to the atoms surrounding the central atom (or atoms), satisfying the octet rule. Distribute remaining electrons as pairs to the central atom (or atoms). Shift pairs of electrons from the surrounding atoms to the central atom, if multiple bonding is necessary for the octet rule.

3. Rules for Predicting the Skeleton Structure Many small molecules or ions consist of a central atom around which are bonded atoms of greater electronegativity. Molecules or polyatomic ions with symmetrical formulas often have symmetrical structures. Oxoacids have oxygen atoms (and possibly other electronegative atoms) bonded to a central atom, with one or more hydrogen atoms usually bonded to the oxygen atoms. Of several possible structural formulas, the one in which the atoms have their usual number of covalent bonds is generally preferred.

4. Resonance Delocalized bonding is a type of bonding in which a bonding pair of electrons is spread over a number of atoms rather than localized between two atoms. A resonance description describes the electron structure of a molecule having delocalized bonding by writing all possible electron-dot formulas. Whenever it is possible to write several plausible electron- dot formulas – differing merely in the allocation of single and double bonds to the same kinds of atoms – delocalized bonding is expected. Do not misinterpret resonance as flipping back and forth between forms.

5. Exceptions to the Octet Rule Molecules and ions with an odd number of electrons in their Lewis structures are called free radicals (or simply radicals). Examples: NO and NO 2 Be, B, Al, Ga, Ge, In, Sn usually form less than a complete octet. P, S, Cl, Kr, and all of the elements below them often form more than a complete octet.

6. Formal Charge (F.C.) Formal Charge is the hypothetical charge obtained by assuming that bonding electrons are shared equally between bonded atoms and that lone pairs of electrons belong entirely to one atom. F.C. on an atom = (valence e’s on free atom) – (lone pair e’s) – ½ (bonding e’s) In selecting between Lewis formulas: – Small (or zero) formal charges on individual atoms are better than large ones. – When F.C. cannot be avoided, the most electronegative atom should be assigned a negative formal charge.

7. Bond Length and Bond Order The distance between the nuclei in a bond is called the bond length. An approximate bond length for atoms joined by single bonds may be calculated by summing the covalent radii of the two atoms involved in the bond. The number of pairs of electrons in a bond is called the bond order. Bond length depends on bond order. As bond order increases, bond strength increases and the nuclei are pulled closer together, which decreases the bond length.

8. Bond Energy The average enthalpy change for the breaking of a bond in a molecule in the gaseous phase is called the bond energy. Energy is required to break bonds (endothermic) and is released when bonds are reformed (exothermic). As bond order increases, bond energy increases. The enthalpy of reaction (  H) is approximately equal to the sum of the bond energies for the bonds broken minus the sum of the bond energies for the bonds formed.