Ms. RobinsonMCHS ChemistryCovalent molecules

Molecular shapes: Ch 6

Main ideas
Section 6.1: Electronegativity / Electronegativity= electron tug-of-war bet. 2 nuclei
Electronegativity
Depends on size, position in periodic table /
Ionic bonds
(doggie bones analogy) / Electrons are distributed this way:
Examples
Covalent bonds / Electrons are distributed this way:
Examples
Polar covalent bonds / Electrons are distributed this way:
Examples
Conductivity and bond type / What conductivity measures
How to tell ionic compounds
How to tell covalent molecules
How to tell metallic elements
Main ideas / Notes and practice

Section 6.2

/ Lewis Dot Diagrams
Valence electrons / Definition
Lewis dot structures
Objective: draw Lewis structures to show the arrangement of valence electrons in molecules and polyatomic ions
Chlorine example / Octet rule: aim to fulfill (there are exceptions)
H O F Br I N Cl diatomic elements
Single bonds: simple
HCl, CH4, Iodomethane CH3I,
Double, triple bonds
Oxygen gas
Nitrogen gas
Polyatomic ions
(covalent within the ion)
OH-, SO4-2, CN- ,
hydronium H3O+, chlorate ClO 2-
Oxidation # of atoms (C in case of CN-)? / Add or remove electrons according to charge
Lewis resonance structures
Objective: draw resonance structures for simple molecules and recognize when they are required
(two pairs of electrons are shared by three atoms)
Example: SO2 , NO2
Example: NO3-
Covalent prefixes
Objective: name binary inorganic compounds using prefixes, roots, and suffixes / Prefixes
Suffixes
Difference between ionic and covalent naming

Concept practice: write the Lewis structure of the following molecules (some may be resonance structures)

NO2-

Methylammonium ion CH3 NO3+

SO32-

CO2

Hydrogen cyanide HCN

Main ideas

Section 6.3

/ Molecular shapes
Why shape matters
Electron pairs repel
Bonds = shared pairs
Unbonded pairs =lone pairs
VSEPR model
Valence Shell Electron Pair Repulsion model
Objective: predict the shape of a molecule from its Lewis structure using VSEPR models
Polarity
Objective: associate the polarity of molecules with their shapes
/ Bond polarity + shape  molecular polarity
Carbon tetrachloride CCl4
Ammonia NH3
# Bonds / # Electron lone pairs / Type of hybrid orbital / Angle between bonded atoms / Geometric shape (lone + bonding pairs) / Molecular shape (bonding pairs/atoms) / Example
2 / 0 / 180
3 / 0 / 120
4 / 0 / 109.5 / CH4 Methane
3 / 1 / 107 / NH3 Ammonia
2 / 2 / 104.5 / H2O
Not / Octet / Rule
5 / 0 / Five sp3d / 120/90 / Trigonal bipyramidal
6 / 0 / Six sp3d2 / 90/180 / Octahedral
Main ideas / Detail: See P 812 note-taking tips

Section 6.4

/ Intermolecular forces
=betweeeeeeeeen molecules
Intermolecular forces, or Van der Waal’s forces
Objective: describe the type of forces that exist between molecules
Ionic/covalent
Boiling points
Melting points / Description:
Dipole forces
Objective: relate theboiling point of a molecular substance to the shape and polarity of its molecules / Description:
Hydrogen bonds
Objective: explain how hydrogen bonds differ from other polar-polar forces
Definition of hydrogen
bond:
Hydrogen bonds need these factors to exist:
1
2
3 / Compare H2O with H2S properties:

London forces / Noble gases boiling points
He Ne Ar Kr Xe Rn
Description of London forces:

1