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An in-depth exploration of chemical bonding concepts, including ionic, covalent, and metallic bonds, Lewis structures, electronegativity, and polar bonds. It also discusses the relationship between chemical bonding and the greenhouse effect.
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4.1 Types of Chemical Bonds 4.2 Naming Compounds and Writing Formulas 4.3 Lewis Structures 4.4 Electronegativity, Unequal Sharing, and Polar Bonds 4.5 Vibrating Bonds and the Greenhouse Effect 4.6 Resonance 4.7 Formal Charge: Choosing among Lewis Structures 4.8 Exceptions to the Octet Rule 4.8 The Lengths and Strengths of Covalent Bonds
^ ^
1A ns^11 2A ns^22 3A (^) ns^2 np^13 4A (^) ns^2 np^24 5A (^) ns^2 np^35 6A (^) ns^2 np^46 7A (^) ns^2 np^57 Group e-^ configuration # of valence e- Lewis Dot Symbol Lewis Symbols and the Periodic Table Lewis Symbols and the Periodic Table Unpaired dots = bonding capacity. Main Group Elements: Members of same family have same number of valence electrons, and similar bonding capacities.
[Ne]3s^1 [Ne] Na Na+^ + e
[Ne]3s^2 3p^5 [Ne]3s^2 3p^6 = [Ar] e
Cl Na
Cl
Draw the Lewis symbols of the monatomic ions formed by calcium and oxygen. Then draw the Lewis structure of calcium oxide (CaO).
(updated later on with the concept of “formal charge”) Double bond – two atoms share two pairs of electrons Triple bond – two atoms share three pairs of electrons C 2 H 2
Multiple Bonds – sharing more than one pair of electrons H 2 CO N 2
Lewis Structures of Charged Species ClO
NO 2
Electronegativity, Unequal Sharing, and Polar Bonds Electronegativity ( ):
Polar Covalent Bonds - Unequal sharing of electrons resulting in an uneven distribution of charge. Difference Bond Type Covalent 2 Ionic 0.4 < EN < 2 Polar Covalent
As seen previously, electronegativity increases moving up to the right in the periodic table. (Noble gases not included.) Bond polarity increases as ΔEN increases. ΔEN = 1.9 0.9 0.7 0.
Rank, in order of increasing polarity, the bonds formed between - O and C Cl and Ca N and S O and Si Are any of these bonds considered ionic?
4.1 Types of Chemical Bonds 4.2 Naming Compounds and Writing Formulas 4.3 Lewis Structures 4.4 Electronegativity, Unequal Sharing, and Polar Bonds 4.5 Vibrating Bonds and the Greenhouse Effect 4.6 Resonance 4.7 Formal Charge: Choosing among Lewis Structures 4.8 Exceptions to the Octet Rule 4.8 The Lengths and Strengths of Covalent Bonds 2 8
Sample Exercise 4.13: Drawing Resonance Structures Draw all the resonance structures of the nitrate ion. Resonance in Organic Compounds 3 2 Benzene = C 6 H 6 and it forms a ring
4.1 Types of Chemical Bonds 4.2 Naming Compounds and Writing Formulas 4.3 Lewis Structures 4.4 Electronegativity, Unequal Sharing, and Polar Bonds 4.5 Vibrating Bonds and the Greenhouse Effect 4.6 Resonance 4.7 Formal Charge: Choosing among Lewis Structures 4.8 Exceptions to the Octet Rule 4.8 The Lengths and Strengths of Covalent Bonds 3 3
Formal charge (FC): FC = # valence electrons – [ (# lone pair e-) + ½ (# shared)] 3 4
What is the most likely Lewis structure for CO 2?
A way to decide which Lewis Structure, out of several possibilities, that’s the most stable and therefore the most likely.
4.1 Types of Chemical Bonds 4.2 Naming Compounds and Writing Formulas 4.3 Lewis Structures 4.4 Electronegativity, Unequal Sharing, and Polar Bonds 4.5 Vibrating Bonds and the Greenhouse Effect 4.6 Resonance 4.7 Formal Charge: Choosing among Lewis Structures 4.8 Exceptions to the Octet Rule 4.8 The Lengths and Strengths of Covalent Bonds 3 9
1. The Incomplete Octet (Central atom in Group 3A) e.g. BF 3