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lecnotes11

# lecnotes11 - MIT OpenCourseWare http/ocw.mit.edu 5.111...

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MIT OpenCourseWare http://ocw.mit.edu 5.111 Principles of Chemical Science Fall 2008 For information about citing these materials or our Terms of Use, visit: http://ocw.mit.edu/terms .

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_______________________________________________________________________________ ________________________________________________________________________________ 5.111 Lecture Summary #11 Readings for today: Section 2.7 (2.8 in 3 rd ed ) – Resonance, and Section 2.8 (2.9 in 3 rd ed ) – Formal Charge. Read for Lecture #12: Section 2.9 (2.10 in 3 rd ed ) – Radicals and Biradicals, Section 2.10 (2.11 in 3 rd ed ) - Expanded Valence Shells, Section 2.11 (2.12 in 3 rd ed ) - Group 13/III Compounds, Section 2.3 (2.1 in 3 rd ed )- The Energetics of Ionic Bond Formation, Section 2.12 (2.13 in 3 rd ed ) – Electronegativity. Topics: I. Lewis structures II. Formal charge III. Resonance structures I. LEWIS STRUCTURES Lewis structures share the total number of valence electrons between atoms so that each atom achieves a noble gas configuration. EXAMPLE: Hydrogen cyanide (HCN) 1. Draw skeleton structure. H and F are always terminal atoms. The element with the lowest ionization energy goes in the middle (with some exceptions). 2. Count the total # of valence e - s. If there is a negative ion, add the absolute value of total charge to count of valence electrons; if a positive ion, subtract. 3. Calculate the total # of e - s needed for each atom to have a full valence shell. 4. Subtract the number in step 2 (valence electrons) from the number in step 3 (total electrons for full shells). The result is the number of bonding electrons .
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