Based on “Inorganic Chemistry”, Miessler and Tarr,

Simple bonding theory

Based on “Inorganic Chemistry”, Miessler and Tarr, 4th edition, 2011, Pearson Prentice Hall

Images from Miessler and Tarr “Inorganic Chemistry” 2011 obtained from Pearson Education, Inc.

Lewis electron-dot diagrams� It is somewhat oversimplified

� It is a good starting point for analyzing bonding

� Bonds result as two atoms share one (or more) pair of electrons

Other pairs of electrons are non-bonding (lone pairs) – they � Other pairs of electrons are non-bonding (lone pairs) – they contribute to shape and reactivity, but do not hold atoms together

� Shared electrons satisfy requirements for both atoms involved

� Most of these Lewis structures are based on the “octet rule” – s and p orbitals (eight electrons) as a noble gas core

Brief review on how to make Lewis

structures� Count valence electrons for all atoms in the molecule; if it is an anion, add one electron per negative charge; if it is a cation, substract one electron per positive charge.

� Arrange atoms placing the least electronegative atom in the center (some exceptions apply)(some exceptions apply)

� Bond atoms together with a single bond

� Satisfy octet rule for all possible atoms starting with terminal atoms and moving to the center atom (do not use more electrons than those you counted at the beginning)

� If you still have electrons, place them in pairs in the center atom

� If you are out of electrons and still need to satisfy the octet rule for some atom, make multiple bonds by moving lone pairs.

Resonance structures� Multiple Lewis structures that differ only in the arrangement of lone pairs or multiple bonds

� All structures must be drawn

� Real structure is a “hybrid” of all structures you draw.

� Overall electronic energy is lowered, making it more stable.� Overall electronic energy is lowered, making it more stable.

� Example: carbonate and nitrate anions, sulfur trioxide

Expanded shells� Sometimes it is impossible to draw a structure consistent with the octet rule for all atoms (more than eight electrons in the central atom)

� For elements at or beyond the 3rd period, use of d orbitals will allow the expansionwill allow the expansion

� The term hypervalent is often used to describe central atoms that have electron counts greater than the usual

Expanded shells beyond 12e–

� IF7 (14e-)

� TaF83- (16e-)

� XeF82- (18e-)

Formal charge� Apparent electronic charge of each atom in a molecule, based on the electron-dot structure.

� All electrons in a bond are assumed equally shared (100% covalency)

� Formal charge = (valence e-) – (e- in Lewis structure)� Formal charge = (valence e-) – (e- in Lewis structure)� Lone pairs are assigned to the atom

� Shared pairs are equally shared (one e- per atom)

� Charge on molecule (zero) or ion = sum of all formal charges

� Formal charges closer to zero are preferred

� Negative charges on electronegative elements are preferred

VSEPR theory� Based on electrostatic repulsion between electrons

VSEPR theory

Lone pair repulsion� Repulsions follow this order lp-lp > lp-bp > bp-bp

� Bonding pairs (bp) are more confined in space than lone pairs (lp)

Lone pair repulsion (AX3E2)

Effect of multiple bonds� VSEPR considers multiple bonds to have a slightly repulsive effect than a single bond because of the repulsive effect of πelectrons

Lone pair vs double bond

� They will compete for the space available

Electronegativity and atomic size effects

Electronegativity and bond angles� If the central atom is the same, as the electronegativity of the halogen increases, it exerts a stronger pull on the bonded electron pairs, reducing the “concentration” of electrons near the central atom.

� If the outer atoms remain the same, as the electronegativity � If the outer atoms remain the same, as the electronegativity of the central atom increases, the bond angle will increase

Molecule+ X-P-X angle Molecule X-S-X angle

PF3 97.8 H2O 104.2

PCl3 100.3 H2S 92.1

PBr3 101 H2Se 90.6

Molecular polarity

Hydrogen bonding� Particularly strong dipole-dipole force ocurring between H (attached to O, N or F) and other O, N or F atom