An ionic bond results from the transfer of electrons from one atom (the cation) to another (the anion) and the subsequent electrostatic attraction between these oppositely charged particles which results.  This can only occur when metals and non-metals bind to one another.

In contrast, covalent bonds result from sharing of electrons and this occurs when two non-metals interact.  If the two non-metals are identical in all respects (identity + chemical environment) then the sharing is equal and the bond is a pure covalent bond.  For the interaction of any two unlike atoms there is necessarily some disparity in the sharing of electrons and the result is a polar covalent bond.  Each of the atoms in the bond takes on some partial charge (as previously seen for water, below) denoted with a d⁺ or d^-.
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Periodic trends for electronegativity:
    - across a period, left to right, electronegativity INCREASES
    - down a column (family), electronegativity DECREASES

In all three types of bonds, ionic, polar covalent, and covalent, it is the valence electrons which are most important in the chemical bonding, and spin pairing results from these interactions (since paramagnetic atoms form diamagnetic compounds).

The Lewis structure for any particular atom in a molecule or ionic substance is a way of representing the atom, it's electronic structure, and it's bonding capabilities, in the compound.  For ionic substances coulombic attraction between oppositely charged ions holds the solid together (quite strongly).  In these cases, no sharing of electrons occurs to a first approximation and the Lewis representation of each ion consists of the element symbol, and the valence electrons arranged in pairs, with dots representing each electron.

For the cations this is most often just the chemical symbol itself, since the cation attains the stable noble gas configuration when it loses all of it's valence electrons.  For the anions, the representation typical consists of the element symbol and four pairs of dots (representing the 8 electrons in the s+p subshells) which allows this anion to attain the next highest closed shell electronic configuration.

If one is dealing with main group elements all d/f electrons can be ignored since these act as "core" electrons, being too low in energy to significantly participate in bonding even when they are present.

Procedure for drawing Lewis structures in ionic substances:

1) assign oxidation states for both the cation and the anion
2) count up the valence electrons in the ions
3) draw the structure as described above

Examples: NaCl, Al₂S₃, CaBr₂

Na = +1, no valence electrons,
Cl = -1, 8 valence electrons
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Lewis noticed that with this set of guidelines, the main-group anions almost invariably contained 8 electrons around them, while the main-group cations contained none (or the 8 s+p electrons from the previous period).  Stability was associated with this configuration which became known as the "octet rule."

Procedure for drawing Lewis Structures

    1) Determine the total # of valence e- in the molecule
            (sum for all atoms), but ignore d-electrons for main group elements
    2) Determine the central atom (often lowest electronegativity)
            - never F
            - never H
            - if formula is something like AB₃ or AB₄,
                then A is the central atom.  (e.g CH₄, NH₃, ClO₄^-)
    3) Draw a skeleton structure:
            - draw lines between all bonded atoms
            - each line represents one covalent bond = 2 e^-
    4) Appropriate the remaining electrons in in such a way as to
        accomplish the following:

            * maximize the number of bonds:
                additional lines may be drawn between atoms and
                these signify multiple bonds.
            * minimize the formal charge = Q_F on the atoms
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        Bonds are denoted by single lines, while two closely spaced dots represent "lone pairs" of electrons belonging to a single atom and which are not directly involved in bonding.

    5) The following "rules" apply to 4)

            * electrons are "counted" for each atom as follows:
                - bonds (lines) count as two electrons for each atom
                in the bond...  (bonds increase the electron count for
                each atom involved in sharing.)
                - All remaining electrons are counted individually.

            * H never gets more than two electrons
                - always just one covalent bond, no lone pairs

            * Li, Be, B, C, N, O, F, Ne never get more than 8 electrons
               - the "real" octet rule! atoms oftenobtain
                    8 valence electrons in the best Lewis structure
                - F always gets one single bond and three lone pairs
                    in the best Lewis Structure
                - terminal halogen atoms behave similarly to fluorine

    Q_F = # of valence electrons in free atom (V) - # of electrons "belonging" to the atom in the molecule.

    * for bond pairs of electrons (lines, BP's) only half the electrons depicted actually "belong" to the atom from the standpoint of assigning formal charges
    * for lone pairs of electrons (dots, LP's) the electrons belong 100% to the atom they are associated with.

Q_F = V - (BP's) - 2 x LP's

Note that the sum of all Q_F's
for all the atoms in a polyatomic ion = the charge on the ion,
for all atoms in a molecule (neutral species) = 0.

SQ_F = charge on the ion
        = 0 for neutral species

    One of the great accomplishments of Lewis Structures is that they allow estimation of the extent of covalent bonding between two atoms in a molecule or molecular ion.

    Single bond: one line in the structure, 2 bonding electrons
    Double bond: two lines in the structure, 4 bonding e-
    Triple bond: three lines, 6 e- between the atoms
    Qaudruple bond (transition metals only): 8 e-, 4 lines

Why is this important?

    * Strength of chemical bonds increase
       with increasing bond order
    * Observed bond distance (see table below),
       the separation between the atoms
       involved in a chemical bond,
       decreases as bond order increases
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bond	length (pm)	bond	length (pm)	bond	length (pm)
C-C	154	N-N	145	C-N	147
C=C	134	N=N	123	C=N	128
CC	120	NN	109.8	CN	116