Covalent bonds occur when two nonmetals share electrons. For example, HCl is a covalently bonded molecule because the atoms share both the single electron from the hydrogen and a single valence electron from the chlorine to fill the valence shells of both atoms.

Covalent bonds can be polar or nonpolar. Polarity in bonds occurs due to differences in electronegativity between the atoms, so electrons can be pulled more towards one atom than the other, resulting in a slightly positive and slightly negative area. The most non-polar covalent bonds occur in the diatomic elements which are two of the same atom bonded together. This most often occurs in the elements Br, I, N, Cl, H, O, F. Many polar covalent bonds occur with fluorine as it is very electronegative.

Ionic bonds are the result of the exchange of electrons between a metal and a nonmetal. For example, NaCl is an ionic compound, because Na wants to lose an electron (low ionization energy) while Cl wants to gain an electron (high electron affinity). Once this occurs, the Na+ ion is attracted to the Cl- ion, and this strong electrostatic attraction creates the ionic bond.

The lattice energy of an ionic compound is the magnitude of the electrostatic attraction within an ionic compound. If the lattice energy is higher, the compound will be harder to break apart, and if the lattice energy is lower, the compound will be easier to break apart. Based on Coulomb’s Law, the lattice energy will be higher in ionic compounds in which the ions are small and have higher magnitudes of charges.

Metallic bonds occur between two or more metal atoms. An alloy occurs when different kinds of metals are bonded together. As seen in the image below, a pure metal has all the same atom in it, an interstitial alloy has a smaller atom (metal or nonmetal) in between the large metal molecules, and a substitutional alloy switches out some pure metal atoms for atoms of a different metal.

A single covalent bond forms between nonmetal atoms when they share an electron pair. When more than one pair is shared, a double or triple bond forms. Double bonds are both stronger and shorter than single bonds, and triple bonds are both stronger and shorter than double bonds.

You will be given a molecule. For example, let's choose ammonia, which is notated as NH3.

Determine the number of valence electrons in the compound using the periodic table. Add or subtract electrons if the molecule happens to be charged

Resonance structures occur when a double or triple bond can be placed anywhere within the Lewis Dot structure without changing the overall shape of the molecule. In the image above, the carbonate ion has a single double bond which can be placed with any of the oxygen molecules without changing the overall structure of the molecule, giving it a resonance structure. On a macroscopic level, resonance structures state that all the resonance bonds in a molecule are actually the same. In the example above, the resonance bond is stronger than a single bond but weaker than a double bond. Common resonance structures include NO2, O3, and benzene (C6H6)

As mentioned in the example above, formal charge is a naming convention that can be used to check your work when creating a Lewis Dot structure. The formal charge of each atom is given by (# of valence electrons)-(# of bonds)-(# of unshared electrons) and should equal zero if the molecule has no charge. If the molecule is charged, the formal charges should add to the overall charge, and the most electronegative atom should have the negative formal charge if there is one.

When identifying a VSEPR model, start with a Lewis Dot structure of the molecule. Count the number of regions of electron density around the central atom in your Lewis Dot structure. Then count how many of these regions are lone electron pairs.

Match the numbers you found out to the numbers in the image above. You may need to memorize this information. Then, follow the matched structure to obtain all of the information that goes with it.

To determine the bond hybridization in a molecule, count the regions of electron density as you would for a VSEPR structure. Remember, lone pairs count once, as do all bonds.

• 2 regions: sp

• 3 regions: sp2

• 4 regions: sp3

• 5 regions: sp3d

• 6 regions: sp3d2

To remember this list, a tip is that the exponents add to the number of regions of electron density, and that the exponents for each letter represent pairs in that orbital, so s cannot exceed 1, p cannot exceed 3, and d cannot exceed 5.

Sigma and Pi bonds are the basis of molecular orbital theory. Sigma bonds are all of the single bonds in a molecular structure as well as the first bond in any double or triple bond. Pi bonds are the second bond in a double bond and the second or third bonds in a triple bond.