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Unformatted text preview: CH101/Ison Name: _________, ____________ last, first Ch 5/6 Worksheet I. Drawing Lewis structures: 1. Add up the number of valence electrons in the molecule. (Add or subtract electrons if the species is an ion) 2. Find the central atom. Usually the first atom listed is the central atom. (H can never be the central atom even if listed first) 3. Connect central atom to the others by one bond each. 4. Count the number of bonding electrons in your Lewis structure. If this number is less than that calculated in (1) start by adding multiple bonds where appropriate. 5. Count the number of bonding electrons again. If this number is less than that calculated in (1) start adding lone pairs to the outer atoms until the octet rule is satisfied for all atoms. 6. Check the formal charge on all atoms. If the species is charged, the negative formal charge should reside on the more electronegative atom. Minimize formal charges by expanding the octet on the central atom if possible (see part III). II. General rules for neutral molecules with atoms following the octet rule: 1. Group IV atoms have 4 valence electrons Æ 4 bonds, 0 lone pairs 2. Group V atoms have 5 valence electrons Æ 3 bonds, 1 lone pair 3. Group VI atoms have 6 valence electrons Æ 2 bonds, 2 lone pairs 4. Group VII atoms have 7 valence electrons Æ 1 bond, 3 lone pairs III. Exceptions to the octet rule: 1. The following atoms never violate the octet rule: C, N, O, F 2. p-block atoms in periods 3 and down can obtain an expanded octet. We will mainly deal with compounds containing the following atoms: Si, P, S, Se, Cl, Br, I 3. Group III atoms are electron deficient and are stable with less then an octet. We will mainly deal with examples of molecules containing B or Al as the central atoms. B and Al both form compounds with only 3 bonds and 0 lone pairs....
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