CH110B

Chapter 9 Notes

As you remember, a chemical reaction is the formation and/or breaking of chemical bonds

Chemical bonds are the forces that hold atoms together in compounds

Bonds are electrostatic forces; attractions between opposite charges and repulsions between like charges (see Fig. 9.1)

When 2 atoms are brought together, there is an optimum distance where the unfavorable repulsions and the favorable attractions are maximized (see the potential energy diagram in Fig. 9.2)

Bonds are responsible for such physical properties as melting point and boiling point

Bonds (and the type and number of atoms present) ultimately determine the shape of a molecule, and structure determines function

Lewis theory

Valence electrons participate in bonding

Metals and nonmetals combine by transferring electrons, forming cations and anions involved in ionic bonds

Nonmetals combine with each other by sharing electron pairs (overlapping orbitals), forming covalent bonds

Complete transfer is a 100% ionic "bond"

Equal sharing is a 100% covalent bond

The above are two extremes

When atoms gain, lose, or share electrons, atoms tend to acquire the electron configuration of a noble gas ("noble gas configuration")

Remember, electrons are transferred, NOT protons, so the atom becomes an ion of the original element, not the noble gas itself....noble gas configurations are hyper-stable

H, Li, and Be follow the duet rule (He configuration), while all other elements (except transition metals) follow the octet rule (Ne, Ar, Kr, Xe, Rn configuration with eight valence electrons)

A Lewis symbol is the element's symbol with "dots" on 4 sides of it to represent the valence electrons

Electrons are shown "unpaired" whenever possible, though the idea of electron spin had not been developed when Lewis' made his theories

Reactions can be drawn showing the transfer of electrons

Looking at tables of ionization energy and electron affinity along with the energies of some changes in state, lattice energies (see Fig. 9.4) and bond energies, the energy of a reaction can be calculated (Ex.- the enthalpy of formation of NaCl in Fig. 9.5, the Born-Haber cycle)

Coordinate covalent bonds occur when a single atom "donates" BOTH of the electrons in a covalent bond

Lone pairs vs. bonded pairs

Multiple bonds

Sharing is not always "equal"

See the Pauling scale shown in Fig. 9.8

ELECTRONEGATIVITY is a measure of an atom's ability to attract bonding electrons to itself (Fr is the least electronegative element, and F is the most electronegative...related to IE and EA)

The greater the difference in electronegativity the greater the ionic character.....the more ionic character to the bond, the more polar it is

A diatomic molecule is completely nonpolar

Polarity is drawn using an arrow with a cross through the tail (see Fig. 9.11)

"Partially positive" atoms are denoted by a "delta-plus", partially negative atoms are denoted by a "delta-minus"

Lewis structures of molecules (see Fig. 9.6)

Drawing Lewis structures

Lewis structures have terminal atoms drawn around a central atom....think of each atom as a four-sided box that must obey the octet (duet) rule.....draw dashes for single bonds, then add lone pairs of electrons to the terminal atoms to get an octet...then add lone pairs and/or form multiple bonds to central atoms as needed to account for the total number of valence electrons

We will discuss how to determine the actual shape of molecules from the Lewis structure in Chapter 10

Central atoms normally have low electronegativity and terminal atoms have high electronegativity

C, N, O, S are often double-bonded, C and N can be triple bonded (see Example 9.7 for mustard gas)

Lewis structures are not always "right"

Formal charge is the difference between the number of valence electrons in the free atom and the number assigned to it in a Lewis structure

Lone pairs are assigned to the atom they surround, and bonded pairs are considered to be shared equally by the two atoms involved in the bond

The smaller the formal charge on each atom, the more likely the Lewis structure

Formal charges for a molecule must sum to zero....sum must be equal to ion charge for a polyatomic ion

Resonance

When there are two or more equivalent Lewis structures which can be drawn for a compound, the bonds are said to be delocalized (Ex.- ozone)

Exceptions to the octet rule

Odd numbers of valence electrons.....free radicals

Incomplete octets

Expanded valence shells

Bond order

Bond length

Bond energy

Unsaturated hydrocarbons

Alkenes

Alkynes

Fats and oils and hydrogenation (oh, my!!!)

Polymers (see Fig. 9.17)