chemical bonding i: basic concepts chapter 9. chapter 9 topics lewis dot structures ionic bond...
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Chemical Bonding I:Basic Concepts
Chapter 9
Chapter 9 Topics
•Lewis dot structures•Ionic bond•Covalent bond•Electronegativity and bond polarity•Writing Lewis structure (Octet Rule) •Exception to the octet rule •Resonance •Formal charge•Bond energy and enthalpy
Bonding: Orbitals
Chemical Bonds
1) Ionic Bond:
Metal + nonmetal Ionic compound
v. Low e affinity High e affinity
v. Low electronegativity High electronegativity
Tend to lose e Tend to gain e
He
Ne
Li (1s2 2s1) Li+ (1s2) + e
F (1s2 2s22p5) + e F- (1s22s22p6)
Li + F Li+ F -
(chemical bond in which electron completely transfer form the metal to the nonmetal)
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2) Covalent Bond
nonmetal + nonmetal Covalent compound
H + Cl H-Cl
(chemical bond in which two or more electrons are shared by two atoms)
F F+
7e- 7e-
F F
8e- 8e-
polar bond and polar molecule
nonpolar bond and nonpolar molecule
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Electronegativity (the attraction that an atom has for the electrons in a bond)
•This is different than electron affinity which is about isolated atom
Same electronegativity (nonpolar bond)•For diatomic it is also nonpolar compound
different electronegativity (polar bond)•For diatomic it is also polar compound
•Any diatomic molecule formed from two element of different electronegativity is a polar molecule and have a dipole moment
H F+ -
H H
Resultant =Resultant = 0
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•The grater the difference between the electronegativity of the atoms in a compound, the greater the polarity of the compound
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•Bond Polarity and Dipole moments
H Cl
2.1 3.0
=0.9
N O
3 3.5
=0.5
C O
2.5 3.5
=1.0
H F
2.1 4.0
=1.9
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Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
Increasing difference in electronegativity
Classification of bonds by difference in electronegativity
Difference Bond Type
0 Covalent
2 Ionic
0 < and <2 Polar Covalent
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Q) Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; andthe NN bond in H2NNH2.
Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic
H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent
N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent
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Sa. Ex. : Order according to polarity:
H-H, O-H, H-Cl, H-S, F-H
H-H < H-S < H-Cl < O-H < F-H
Covalent bond Polar covalent bond
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O = C = O O
HH
Polar bonds but nonpolar molecule Polar bonds and polar molecule
Sa. Ex. : Which ones have a dipole moment:
HCl, Cl2, SO3, CH4, H2S
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•Lewis Structures (Octet Rule)
Octet rule: An atom tend to lose or gain or share electrons until its outer shell (valence shell) contain 8 electrons (nobel gas configuration)
Lewis Symbol:
Help us knowing number of bonds but not number of unpaired electrons
B O3 bonds3 v.e.
2 bonds6 v.e.
Drawing Lewis Structures:
A) Deciding the central atom:-
In general, the central atom is the first atom in the binary moleculese.g. : CO2, CO3
2-, NH3, NO2, NO3-, SO3,
SO42-
exceptions are : H2O and H2S
B) Distributing the valence electrons:-1) Count all the valence electrons
2) Place one pair of electrons in each bond
3) Complete the octets for the non-central atoms
4) Place the remaining electrons on the central atom.
5) If the central atom still has less than an octet, use multiple bonds
Sa. Ex. : Write Lewis structure for :
a) HF b) N2
c) NH3 d) CH4
e) CF4 e) NO+
f) CO2 f) NO3-
g) CO32- g) H2S
Exception of the octet rule
2) BF3 or BCl3
3) SF6
1) BeCl2
4) PCl5
central atom with principal quantum number n > 2
The Incomplete Octet
Odd-Electron Molecules
5) NO
N O
Sa. Ex. : Write Lewis structure for :
a) ClF3 b) XeO3
c) RnCl2 d) BeCl2
e) ICl4-
•Resonance
•A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure because the bond length measurement show that the actual bond length is similar.
O O O+ -
OOO+-
O C O
O
- -O C O
O
-
-
OCO
O
-
-
e.g.
< <
Q) What are the resonance structure of NO2- and NO3
-
•Formal charge (Charge calculated from an atom in the Lewis structure)
•Useful to decide which Lewis structure is preferred when more that one Lewis structure are possible. The less the formal charge the more stable the molecule.
Formal Charge=(no. of Valence e- in the isolated atom) – (no. of bonds) – (no. of unshared e-)
Note that: The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.
Using formal charge to select Lewis structure
1) Construct Lewis structure that obey the octet rule.
2) Calculate the formal charge of this molecule
a)If the formal charges are zero, no need to look for another structure
b)If the formal charges are not zero, find another structure that do not follow the octet rule by moving unshared pair of electrons
c) For Lewis structures having similar distributions of formal charges, the most stable structure is the one in which negative formal chargesare placed on the more electronegative atoms.
Q) Which is the most likely Lewis structure for CH2O?
H C O HH
C OH
H C O HC = 4 ve-
O = 6 ve-
2H – 2x1 ve-
12 ve-
2 single bonds (2x2) = 41 double bond = 4
2 lone pairs (2x2) = 4Total = 12
formal charge on C = 4 -3 - 2 = -1
formal charge on O = 6 -3 - 2 = +1
-1 +1
Formal Charge=(no. of Valence e- in the isolated atom) – (no. of bonds) – (no. of unshared e-)
formal charge on H = 1 -1 - 0 = 0
C = 4 ve-
O = 6 ve-
2H = 2x1 ve-
12 ve-
2 single bonds (2x2) = 41 double bond = 4
2 lone pairs (2x2) = 4Total = 12
HC O
H
formal charge on C = 4 -4 - 0 = 0
formal charge on O = 6 -2 - 4 = 0
0 0
Formal Charge=(no. of Valence e- in the isolated atom) – (no. of bonds) – (no. of unshared e-)
the most likely Lewis structure for CH2O is :
H C O H
-1 +1 HC O
H
0 0
Q) Calculate the formal charge for the atoms in:
1)
2)
3)
4)
Sa. Ex. : Write the possible Lewis structure for XeO3 and show the most appropriate structure to the formal charges