Covalent bonds can be single, double, or triple. If only one pair of electrons is shared, a single bond is formed. This single bond is a sigma bond (s), in which the electron density is concentrated along the line that represents the bond joining the two atoms.

 

    However, double and triple bonds occur frequently (especially among carbon, nitrogen, oxygen, phosphorus, and sulfur atoms) and come about when atoms can achieve a complete octet by sharing more than one pair of electrons between them. If two electron pairs are shared between the two atoms, a double bond forms, where one of the bonds is a sigma bond, and the other is a pi bond (p). A pi bond is a bond in which the electron density is concentrated above and below the line that represents the bond joining the two atoms. If three electron pairs are shared between the two nuclei, a triple bond forms. In a triple bond, the first bond to form is a single, sigma bond and the next two to form are both pi.

 

    Multiple bonds increase electron density between two nuclei: they decrease nuclear repulsion while enhancing the nucleus-to-electron density attractions. The nuclei move closer together, which means that double bonds are shorter than single bonds and triple bonds are shortest of all.

 

    Metallic bonds exist only in metals, such as aluminum, gold, copper, and iron. In metals, each atom is bonded to several other metal atoms, and their electrons are free to move throughout the metal structure. This special situation is responsible for the unique properties of metals, such as their high conductivity.

 

 

    Here are some rules to follow when drawing Lewis structures—you should follow these simple steps for every Lewis structure you draw, and soon enough you’ll find that you’ve memorized them. While you will not specifically be asked to draw Lewis structures on the test, you will be asked to predict molecular shapes, and in order to do this you need to be able to draw the Lewis structure—so memorize these rules! To predict arrangement of atoms within the molecule

1. Find the total number of valence electrons by adding up group numbers of the elements. For anions, add the appropriate number of electrons, and for cations, subtract the appropriate number of electrons. Divide by 2 to get the number of electron pairs.
2. Determine which is the central atom—in situations where the central atom has a group of other atoms bonded to it, the central atom is usually written first. For example, in CCl₄, the carbon atom is the central atom. You should also note that the central atom is usually less electronegative than the ones that surround it, so you can use this fact to determine which is the central atom in cases that seem more ambiguous.
3. Place one pair of electrons between each pair of bonded atoms and subtract the number of electrons used for each bond (2) from your total.
4. Place lone pairs about each terminal atom (except H, which can only have two electrons) to satisfy the octet rule. Leftover pairs should be assigned to the central atom. If the central atom is from the third or higher period, it can accommodate more than four electron pairs since it has d orbitals in which to place them.
5. If the central atom is not yet surrounded by four electron pairs, convert one or more terminal atom lone pairs to double bonds. Remember that not all elements form double bonds: only C, N, O, P, and S!