CHEM 150 1nd Edition Chapter 8: Chemical Bonding and Climate Change Lecture 26Outline of Last Lecture • Chemical Bonds (section 8.1)• Understand the definitions of and differences between ionic bonds, covalent bonds, and metallic bonds.• Understand how figure 8.1 (p 369) relates bond energy and bond length to the attractive and repulsive forces present in the atom.Outline of Current Lecture • Lewis dot symbols: The atomic symbol surrounded by a dot for each valence electron present in the atom.• Octet Rule: Main group elements gain, lose or share electrons to achieve a set of eight valence electrons.• Lewis Structures for covalently bonded (molecular) compounds:• Bonds are shown as lines. Lone pairs of electrons (non-bonding) are shown as pairs of dots.• General Rules:• All atoms must have a set of eight valence e- in the final molecule. [Except: H, Be, B] • An atom will usually form one bond for each electron it requires. [One bond per unpaired e- in the Lewis dot symbol.]Current LectureQ: Draw Lewis dot symbols for the following:O Ca Ge Ar Q: Draw Lewis Structure of F2O, NH3, CCl4, and C3H6A:-Some guidelines for Lewis structures:Use the “central atom” premiss: One atom in the center with the others connected to it.• Molecules with multiple C or N will tend to have multiple “centers”• C is the only element that forms long chains of itself.• Avoid O-O and F-F bonds. Never make chains of O or F.• Do not make triangles or squares!• F does not form double bonds (ever!)• In oxoacids, the “ionizable hydrogen” is not attached to the central atom. The H isbonded to an O that is bonded to the central atom-The “procedure” for drawing Lewis structures are in section 8.2 (p 371-372) 1) skeletal structure2) valence electron count3) octets4) multiple bonds (as necessary)Q: Draw Lewis structures for:
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