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Lecture Comments (9)

0 answers

Post by Zachary McCoy on January 4, 2014

Please note: The electronegativity of N is closer to 3 than 3.5 (even closer to 3 than is Cl)! This becomes especially important when you get to organic chemistry (the significance being that oxygen is more electronegative [i.e. NOT nearly equal to nitrogen] than nitrogen).

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Post by Maimouna Louche on June 5, 2012

Thanks!!!!

0 answers

Post by Imran Khalique on May 11, 2012

bondddddddddddddd

1 answer

Last reply by: Zachary McCoy
Sat Jan 4, 2014 7:11 PM

Post by Zahra Malak on December 18, 2011

Dear Professor Goldwhite,
First of all, thanks for your outstanding approach in teaching chemistry.
Maybe not quite related to this video, but I have this question about Se: Can we build Se2 with double bond or not? If not, does it have any relation with the size of this particular element?May you elaborate on it?

3 answers

Last reply by: Zachary McCoy
Sat Jan 4, 2014 7:10 PM

Post by Regina Kim on November 20, 2011

He made a simple miscalculation. The formal charge of F for the first resonance structure is 7-(4+2)=+1, and from that the formal charge of N is -1.
In his answer got, 7-(4+2)=-1 and therefore incorrectly concluded that the N in the structure was +1.
Not a big deal, but just in case someone gets confused. :)

Related Articles:

Covalent Bond, Lewis Structures, Molecular Orbitals

  • Covalent bond: sharing electron pairs

  • Number of valence electrons for atom = Group number

  • Octet rule – rule of 8; 8 electrons around main group atoms in molecules except H, two electrons

  • Rules for Lewis structures: COUNT valence electrons in molecules; put 2 electrons between bonded atoms; use remaining electrons to get 8 shared by other atoms in lone pairs and/or multiple bonds

  • Exceptions to octet rule: Be, B etc. in Groups 1 and 2; third period elements like S, P, can exceed 8.

  • Electronegativity: power of atom in a bond to attract electrons

  • Electronegativity decreases down a group; increases from left to right along a period

  • Bonds can be fully ionic; polar; non-polar

  • Overlaps of atomic orbitals can give molecular orbital’s

  • Sigma orbitals: cylindrical symmetry around internuclear axis

  • Pi orbitals: plane of symmetry between nuclei

Covalent Bond, Lewis Structures, Molecular Orbitals

Lecture Slides are screen-captured images of important points in the lecture. Students can download and print out these lecture slide images to do practice problems as well as take notes while watching the lecture.

  • Intro 0:00
  • Covalent Bond 1:05
    • Hydrogen Atom
    • Lewis
    • Sharing of Electrons
  • Valence Electrons 4:13
    • Example: Lewis Symbols
  • Lewis Structures of Atoms and Ions 10:05
    • Example: Ions
  • Octet Rule 13:34
    • Noble Gases
  • How to Draw Lewis Structures 17:35
    • Single Bond
    • Double Bond
    • Triple Bond
    • Examples: Compounds
  • Formal Charges 32:44
    • Example: 3 Structures of FCN
  • Exceptions to the Octet Rule 40:17
    • Fewer Than 8 Electrons
    • Example: Boron (Electron Deficiency)
    • More Than 8 Electrons
    • Example: Compounds
  • Electronegativity 48:09
    • Example: Values
    • Ionic or Polar Covalent Bond
  • Molecular Orbitals 54:16
  • Sigma and Pi Bonding 55:56
  • Additional Example 1
  • Additional Example 2