Lewis Structure and Bonding Capacity

1. Lewis Structure and Bonding Capacity March 17, 2008

2. Covalent bonds (reviews) • Non-metal atoms achieve stability (octet) through sharing valence electrons • Octet rule: An atom (other than H) tends to form bond until it is surrounded by 8 electrons (mainly applicable to elements with only 2s and 2p subshells) • Single bond: Two atoms are held together by one electron pair • Multiple Bonds: Two atoms are held together by two or more pairs of electrons

3. Electronegativity (reviews) • Electronegativity: relative ability of an atom, when bonded, to attract electrons • Trends: EN increases from left to right across a period; decreases down a group • Atoms of elements with widely different EN tend to form ionic bonds (NaCl, CaO, etc) • Atoms of elements with comparable EN tend to form polar covalent bonds • Only atoms of the same element, which have same EN, can by joined by pure covalent bond • An ionic bond forms when the EN difference between two bonding atoms is 2.0 or more

4. Writing Lewis structures for molecules • Lewis structures do NOT represent the complete picture of covalent bonding, but help to explain the bonding scheme in many compounds and account for properties and reactions of molecules. • For example: Nitric acid (HNO3)

5. Steps for writing Lewis structures for molecules: • Four steps to follow: • Least electronegative atom occupies the central position. Hydrogen & fluorine usually occupy the end positions. • Count the total number of valence electrons present. For polyatomic anions, add the number of negative charges to that total. (Example: CO32- ion  we add two electrons because 2- charge indicates that there are two more electrons than are provided by the neutral atoms.) For polyatomic cations, subtract the number of positive charges from this total. (Example: NH4+ ion  we subtract one electron because 1+ charge indicates a loss of one electron from the group of neutral atoms.)

6. Steps for writing Lewis structures for molecules (continued): • Draw a single covalent bond between the central atom and each of the surrounding atoms. Complete the octets of the atoms bonded to the central atom. (For hydrogen, only two electrons.) Electrons belonging to the central or surrounding atoms must be shown as lone pairs if they do not participate in bonding. The total number of electrons to be used is that determined in step 2. • If octet rule is not satisfied for the central atom, try adding double or triple bonds between the surrounding atoms and the central atom, using the lone pairs from the surrounding atoms.

7. Examples 1: Nitrogen trifluoride (NF3) • Step 1: The N atom is less electronegative than F, so the skeletal structure of NF3 is: F N F F • Step 2: The outer-shell electron configurations of N and F are 2s22p3 and 2s22p5, respectively. Thus there are 5 + (3 x 7), or 26, valence electrons to account for in NF3. • Step 3: We draw a single covalent bond between N and each F, and complete the octets for the F atoms. We place the remaining two electrons on N:

8. Examples 2: Nitric acid (HNO3) • Step 1: The skeletal structure of HNO3: O N O H O • Step 2: The outer-shell electron configurations of N, O and H are 2s22p3 and 2s22p4, and 1s1 respectively. Thus there are 5 + (3 x 6) + 1, or 24, valence electrons to account for in HNO3. • Step 3: We draw a single covalent bond between N and each of the three O atoms, and between one O atom and the H atom. Then we fill in electrons to comply with the octet rule for the O atoms: • Step 4: We see that this structure satisfies the octet rule for all the O atoms but not for the N atom. Therefore we move a lone pair from one of the end ) atoms to form another bond with N. Now the octet rule is also satisfied for the N atom:

9. Examples 3: Carbonate ion (CO32- ion) • Step 1: C is less electronegative than O. Therefore, it is most likely to occupy a central position as follows: O O C O • Step 2: The outer-shell electron configurations of C and O are 2s22p2 and 2s22p4, respectively, and the ion itself has two negative charges. Thus the total number of electrons are 4 + (3 x 6) + 2, or 24. • Step 3: We draw a single covalent bond between C and each O and comply with the octet rule for the O atoms: • Step 4: Although the octet rule is satisfied for the O atoms, it is not for the C atom. Therefore we move a lone pair from one of the O atoms to form another bond with C. Now the octet rule is also satisfied for the C atom:

10. Exceptions to the Octet Rule • First category of exception – Incomplete Octet: • In some compound, the number of electrons surrounding the central atom in a stable molecule is fewer than eight; e.g., BeH2; BF3 • Second category of exception – Odd-Electron Molecules: • Some molecules contain an odd number of electrons. Among them are nitric oxide (NO) and nitrogen dioxide (NO2). • Third category of exception – Expanded Octet • Atoms of the second-period elements cannot have more than eight valence electrons; but atoms in and beyond the third period can have more than eight valence electrons because there are 3s, 3p and 3d subshells. • Examples: SF6. This molecule has 12 bonding electron pairs accommodated in 6 orbitals (one 3s, three 3p, and two of the five 3d orbitals). Sulfur can also form many compounds in which it obeys the octet rule, e.g., SCl2 (which is surrounded only by 8 electrons).

11. Bonding Capacity Activity