


Study with the several resources on Docsity
Earn points by helping other students or get them with a premium plan
Prepare for your exams
Study with the several resources on Docsity
Earn points to download
Earn points by helping other students or get them with a premium plan
Community
Ask the community for help and clear up your study doubts
Discover the best universities in your country according to Docsity users
Free resources
Download our free guides on studying techniques, anxiety management strategies, and thesis advice from Docsity tutors
The concept of covalent bonds, discussing the factors determining bond length, the nature of different types of covalent bonds, and the use of lewis structures to represent them. Topics include polar covalent, ionic, and nonpolar bonds, as well as the electronegativity difference and bond dipoles.
Typology: Summaries
1 / 4
This page cannot be seen from the preview
Don't miss anything!
The bond length is the distance at which the repulsion of the two nuclei equals the attraction of the valence electrons on one atom and the nucleus of the other.
H-O < H-Cl < H-Br. The order of increasing atom size.
The bond dipole points toward the more electronegative atom in the bond, which is the atom with the lower energy orbital. The magnitude of the dipole increases as the electronegativity difference (orbital energy difference) increases. The W-X bond is not polar because the orbital energies of W and X are identical, and the Y-Z bond in more polar than the U-V bond because the energy separation of the bonding orbitals is
∗
2-
ER 6(8) = 48 2(8) + 4(2) = 24 3(8) + 2 = 26 10(8) = 80 VE 2(5) + 4(6) = 34 4 + 4(1) + 6 = 14 1 + 7 + 2(6) = 20 10(6) + 2 = 62 SP 1 / 2 (48 - 34) = 7 1 / 2 (24 - 14) = 5 1 / 2 (26 - 20) = 3 1 / 2 (80 - 62) = 9
H O C O H H C
O
O H
A B
O N
O
N
O
O
H C C H
H C
H
C C
H
H
H C C C
H
H
H
H C
H
H
C
H
C
H
H F C
O
or F
1-
1+
1-
1-
or O^ N^ O
2- , CHO (^2) 1- , NO (^3) 1- , and NO (^2) 1- all have more than one important resonance form. Although two resonance structures can be drawn for N 3
1- , the one that places a -2 formal charge on a nitrogen is not expected to be important.
Structure A involves charge separation, but structure B does not. Structure B is preferred.
The formal charges are shown in the Lewis structure. Oxygen is more electronegative than nitrogen so all bonding N-O electrons are assigned to oxygen when determining oxidation states. The electron pair in the N-N bond is split between the two nitrogen atoms. Nitrogen is in group 5A, but only the one electron in the N-N bond is assigned to in determining the oxidation state: OXN = 5 - 1= +4. The nitrogen atoms are each in the +4 oxidation state. Each oxygen is assigned eight electrons for an oxidation state of -2. Note that the same results could have been obtained by applying the rules given in Section 4.3.
O C
O
O H C
O
H C O H C
H
H
O
H A B C D
CO bond order
4 / 3 2 3 1 Bond length increases as the bond order decreases: C < B < A < D Bond energy increases as the bond order increases: D < A < B < C