Chapter 4: Covalent Bonding and Chemical Structure Representation The Octet Rule -An atom with 8 electrons (an octet ) in its outer shell has the same number of valence electrons as the noble gas in the same row of the periodic table. This is often referred to as a closed shell. -The fact that atoms in covalently bonded molecules are observed to share their valence electrons to attain a complete octet resulted in the octet rule. To determine whether an atom obeys the octet rule or not: 1. Count the lone pair electrons on the atom. 2. Count all the bonding electrons (twice the number of bonds to the atom). 3. If the sum of #1 and #2 is eight, the atom obeys the octet rule. Formal Charge formal charge the difference between the number of valence electrons and the number of electrons surrounding an atom in a particular Lewis structure formal charge = valence electrons - electrons surrounding atom -Label formal charges using these conventions: a) Show only non-zero formal charges. b) Formal charges are always circled. -The sum of the formal charges adds up to the overall charge.
How to Draw Lewis Structures 1. Determine which atoms are central (bonded to two or more atoms) and terminal (bonded to only one other atom). - Hydrogen atoms are always terminal. - Central atoms are generally (not always) those with the lowest electronegativity. - Atoms with the highest electronegativity are generally (not always) terminal. 2. Write the atomic symbol for the central atom. Distribute the symbols for the terminal atoms around the central atom. 3. Place the valence electrons around each atom as dots. Any electrons in excess of 4 shall be paired. For cations, remove a valence electron from the central atom. For anions, add a valence electron to the mot electronegative terminal atom which has an incomplete octet. Be sure that none of the atoms at this stage have eight valence electrons. 4. Unpaired electrons on the central atom are paired with unpaired electrons on each of the terminal atoms to form single bonds. 5. If unpaired electrons are present on adjacent atoms after step 4, then form additional bonds between these atoms until all the unpaired electrons are paired. 6. If necessary, rearrange electrons so that each atom has an octet of electrons (H will have a duet). This may involve forming multiple bonds using lone pairs to achieve the octet for adjacent atoms. 7. Redraw the structure, remembering to draw lone pairs as pairs of dots and bond pairs as lines. If you are drawing a complex ion, the structure should be surrounded by square brackets ad the charge is indicated outside the top right of the square brackets. 8. Label the formal charges. 9. To double check your Lewis structure, count the overall number of electrons in the structure (each line represents 2 and each dot represents 1). The overall number of electrons should equal the sum of the valence electrons from each atom in the formula, taking into account the overall charge on the species. 10. If more than one Lewis structure can be drawn, follow these guidelines to determine the best one: - Select the structure in which formal charges are minimized (as close to 0 as possible). - Select the structure in which negative formal charges are on the more electronegative atoms and the positive formal charges are on the less electronegative atoms.
Resonance Structures resonance structure when two or more valid Lewis structures can represent the arrangement of electrons in the same molecule or molecular ion, each individual structure is called a resonance structure. -the positions and connectivity of the atomic nuclei must be identical in all resonance forms -Resonance structures are drawn with a double-headed arrow between them. -The actual electron distribution is an average of the electron distribution in the individual structures. ex. for [CO 3 ] 2- -The location of the nuclei and the connectivity between the oxygen atoms are identical, but the arrangement of the electrons is different. ex. for [C 4 H 7 ] - -These two resonance structures have the carbon and hydrogen atoms in the same place, but the double bond and the lone pairs are in different positions.
Exceptions to the Octet Rule 1. Incomplete octets: -The group 13 elements only have three valence electrons in their outer shell to use in bonding. -If they share these electrons with three other atoms, they can only achieve a maximum of six electrons. ex. boron trifluoride: -Compounds in which an atom has an incomplete octet are referred to as being electron deficient. 2. Unpaired electrons: -Species that contain unpaired electrons are called free radicals or radicals. ex. nitric oxide:
3. Hypervalence (expanded octets): -Elements of the third and later periods of the p-block sometime have more than an octet of electrons in chemical compounds. -When compared to the elements of the second period, the larger size of heavier elements permits them to accommodate more than eight electrons in their valence shell. -In Lewis structures, expanded octets result from transforming a lone pair into two unpaired electrons that subsequently can form two additional bonds. -The octet rule will not be exceeded unless necessary to form bonds with more than four atoms or to minimize formal charge. -The octet rule will not be exceeded if it results in placing a negative formal charge on an atom of lower electronegativity.