Reviews for Chapter 8, 9 and 12
Exam 4
Must be familiar with concepts in
Chapter 7 to better understand chapter
8 and 9
Chapter 8 Bonding and Molecular Structure
8.1
Chemical Bonds: A Preview
Chemical bonds result from attractive and repulsive forces between electrons and
nuclei of the 2 atoms forming molecules
Ionic Bonds and Ionic Crystals (check eg. CaCl2)
Lewis Structures for Ionic Compounds: Electrons Transferred
8.2
The Lewis Theory of Chemical Bonding: An Overview
Lewis Structure is simply chemical symbol + # valence electrons around it
A. Most atoms "want" eight electrons around them (octet rule)
B. Exception: duet rule for hydrogen and helium
C. A covalent bond is a shared par of electrons to get an octet around each
atom also some have expanded octet
Lewis Structures of Some Simple Molecules (H2, F-F, O=O, N≡N )
.
1.Hydrogen,F atoms are rarely central atoms; they are almost always terminal atoms.
2.The central atom usually is the atom with the lowest Electronegativity (except for H
that often has a relatively low Electronegativity).
3. In oxoacids, H usually is bonded to an oxygen atom.(HNO3, H2SO4 etc)
4.Molecules usually are clusters of atoms, rather than long chains.
5Understand bonding patterns of C, N, O and halogens.
Strategies for Writing Lewis Structures
A
B. Lewis structures are easier to draw if one follows a systematic plan in
constructing them. One such plan is as follows.
1.Sum the valence electrons in the structure.
a.For each atom, the number of valence electrons equals that atom’s group
number.
b. Add electrons if the entire species is an anion, subtract them if the species is a
cation.
c. Take off electrons ( 2 electrons per bond) to form single bonds,
2. Arrange the atoms in the preferred skeletal structure (central atoms is
lowest Electronegativity etc) and place a pair of electrons between
each pair of bonded atoms. Subtract single bon elctrons
3. Place remaining electron pairs around terminal atoms so that each has
an octet.
4. Place LEFT OVER electron pairs on the central atom.
5. Create multiple bonds as necessary to make sure each atom follows the
octet rule.
8.3
Learn to calculate formula charge
sum of formal charges = charge of the species
8.4
Learn to draw plausible resonance structures
8.5
Molecules that Don’t Follow the Octet Rule (SF6, IF5, XeF4, BeCl2, BH3 etc.)
Exceptions to octet rule
8.6
Molecular shapes see shape table, and visualize symmetric shapes
8.7
Electronegativity (EN) and BOND POLARITY (Polar Covalent Bonds )
Electronegativity (tendency to pull shared electron towards it)
1. Increases as you move left to right along a Period HIGHEST is F
2. Decreases as you go down a group
3. Noble gases have very low EN
Difference of Pure covalent vs polar covalent vs ionic bond based on EN difference
Chemical bonds result from attractive and repulsive forces between electrons and nuclei of
the 2 atoms forming molecules: Types of chemical bonds based on electronegativity
difference between the bonding atoms
DEN = 0 :Pure covalent bond;
DEN = 0 to 0.39 Nonpolar Covalent
DEN = 0.4-1.9 :Polar covalent bond;
DEN ≥1.9 or Has a metal: Ionic
8.8
Molecualr polarity (net dipole moment µ)
- A molecule may have polar bonds, BUT all cancelling out due to
symmetry, results in NON POLAR molecule (e.g. CO2, CH4)
- Combination of shapes, bond polarity additions and symmetry
-Lone pair on central atom: Usually Polar (except XeF4, SeF2) or linear shapes
- Unsymmetrical molecule is Polar (e.g. CH3Cl, CH2Cl2)


H — F or H — F


8.9 Bond order, Bond length and Bond energy/strength for Covalent Bonds
Bond order = # of bonds for a pair of atoms
e.g. Triple bond has Bond order =3
Bond length is inversely relationship to bond order
Also depends on atomic sizes of atoms making up the bonds
Bond strength also called as bond energy or bond dissociation energy is
Highest for triple bond and least for single bond
Bond energy is defined as the energy required to break covalent bonds holding up 2 atoms
Calculate D H rxn =∑ Bond energies broken - ∑ Bond energies formed
8, 9
Bonding Theory and Molecular Structure
8.6 The Valence-Shell Electron-Pair Repulsion (VSEPR) Method
-minimize electron –electron repulsion between bonded electrons pairs
-Predicts shapes, bond angles and dipole moments
-MEMORIZE Shape table, bond angles
-REMEMBER we care only about electron groups on the central atoms and then
-Predict Electron geometry and MOLECULAR GEOMETRY
-9.1
Atomic Orbital Overlap
This region results from the overlap of the two atomic orbitals, one on each atom.
The greater the degree of overlap, the stronger is the bond.
Most often, the two atomic orbitals that overlap to form the bond are each half-filled
9.2
Hybridization of Atomic Orbitals
A.
We hybridize not just to explain the bonding, but also to create the correct number
of orbitals on the central atom, pointing in the correct direction, to create the molecule.
9.2
number of
valence shell
electron groups
electron group
geometry
central atom
hybridization
orbitals needed for
combination
a.
2
linear
sp
s p
b.
3
trigonal planar
sp 2
s p p
c.
4
sp 3
s p p p
d.
5
tetrahedral
trigonal
bipyramidal
dsp 3
d s p p p
e.
6
octahedral
d 2 sp 3
d d s p p p
[ f.
4(special case)
square planar
dsp 2
d  s  p  p]
Hybrid Orbitals and Multiple Covalent Bonds
B.
Sigma bonds
a.
b.
c.
d.
H C C H
F
F
General distinctions between sigma and Pi
( )
Result from the end-to-end overlap of two orbitals.
Form between all varieties of orbitals.
Are the first bond that forms in a multiple bond, the only bond that forms in a single bond.
Produce the most effective overlap, the strongest bond, at a given internuclear distance.
( )
Pi bonds (ALSo results in geometric isomerism)
a.
Result from the side-to-side overlap of p orbitals.
b. Only form between two p orbitals
c.
Are the second and third bonds of multiple bonds
d. Produce weak overlap and restrict rotation around the bond.
Learn to recognize sigma
and pi in a molecule
Chapter 12
12.1 Properties of Solids and Liquids
A. Properties of liquids
1. High densities compared to gases
2. Indefinite shape, fluid
3. Definite volume
B. Properties of solids
1. High densities compared to gases
2. Definite shape
3. Definite volume
4. May be crystalline or amorphous
12.2 Interaction between ions
• Energy Changes in Ionic Compound, Coloumb’s law
• Ionic bond energy is directly proportional to Magnitude of Charges
• Ionic bond energy is inversly proportional to distance separting the ions
• Recollect Ion sizes
12.3-12.5 Interactions between Molecules
A. Intermolecular forces, ion dipole forces, solvation energy etc
B. Thermal energy = energy of motion
12.4 Types of Intermolecular Forces: Dispersion, Dipole-Dipole, and
Hydrogen Bonding
A. Dispersion force, also known as London forces
1. Weakest intermolecular force
2. Instantaneous dipole
3. Involved in every intermolecular interaction
4. Increases with size and molar mass of compounds/atoms
B. Dipole-dipole force- ONLY IN POLAR MOLECULE
1. Stronger force than dispersion forces
2. Strength a function of dipole moment and structure
C. Hydrogen bonding
1. Strongest intermolecular forces
2. Only involving F, O, or N bonded to a hydrogen atom within a molecule
Intermolecular forces affect state, boiling and melting points
EXAM 4- 100 POINTS
Part 1 Multiple Choice –Show calculations for partial/full credit
10 questions
Part 2 1) VSEPR, Lewis structure predict MOLECULAR geometry,
Dipole moment, hybridization around central atom, resonance
2) Sigma and pi bonds
eg. N≡N how many sigma and how many pi bonds?
3) Enthalpy of reaction from bond energies
STOP HERE for EXAM 4 review
SLIDES ARE BEING CONTINUED for finals review
Intermolecular Forces in Action: Surface Tension and Viscosity
A. Surface tension – why some things more dense than water can still float
B. Viscosity – resistance of liquid to flow
12.5 Evaporation and Condensation
A. Evaporation and vaporization – liquid to gas
B. Condensation – gas to liquid
C. Dynamic equilibrium
D. Boiling
E. Energy change = heat of vaporization
12.6 Melting, Freezing, and Sublimation
A. Melting – solid to liquid
B. Freezing – liquid to solid
C. Energy change = heat of fusion
D. Sublimation – solid to gas
Chapter 13 Phase diagrams
Finals
REVIEW CHAPTERS 1 through 9, 11-13
Finals on Tuesday Dec 20th 9:30-11:30am
Room T-109
40 multiple choice questions
Good Luck and Happy Holidays!!
Chapter 1
Conversion e.g. cm-> meter->kilometer, liters-> ml etc.
Density calculations, density lab, pay attention to sig fig!!
Chapter 2
No. protons, no. neutrons, no. electrons, charge problem
Isotopes, writing in the 14C convention, atomic mass Calculation
Correct formulas vs names for chemicals
Names of compounds molecular and ionic compounds
Chapter 3
how many moles, number of atoms/molecules ?
Empirical formula- easy one; Molecular formulas- Ratio
Balancing equation- write the coefficients?-Sum
Calculations in chemical reaction how much reactant will
produce how much product, moles grams etc
% yield, mass %
Chapter 4
Molarity, Dilution, Acid base titration
net ionic equation(single-double lab)
precipitation reaction stoichiometry
Chapter 11 Gases
Boyle’s law relation of P and V
Charles law relation of V and T
Avogadro's law relation of V and N
combined gas law PVT relationship
PV =nRT, gas laws
Gas reactions-stoichiometry-simple one like e.g. in book
Molar mass, density, partial pressure-Exam 2 e.g.
Chapter
Exothermic? Endothermic?- Exam 2!
Specific heat –Simple
Moles calculation with DH reaction-see quiz 6
Hess’s law
Chapter 6 Quantum
Bohr’s H-atom -Lab
Allowed quantum number
orbital, how many electrons can sub-shells have? S-2, p-6 etc.
Chapter 7 Atomic structure
electronic configuration of atoms and ions, unpaired electrons
valence electrons, periodic properties V. important (IE, size)
Chapter 8, 9
Bonding , Bond energy calculations
Correct Lewis structure
Molecular geometry –important-2 questions!
Type of bonding? Ionic, Polar covalent, non polar covalent etc.
Polar- non polar??, Hybridization-EASY, resonance, VSEPR
Chapter 12,13
Vapor pressure- High vapor pressure means low boiling!!
Critical tempertaure
Triple point, CO2 and H2O phase diagram
Polar dissolves polar and non polar dissolves non polar
heats of vaporization high implies high boiling point and
low vapor pressure etc
H-bonding INTER molecular bond between H and O or F or N
Electrolytes and metallic bonding; London dispersion forces
• 6.023x1023 ATOMS = 1 mole OF ATOMS
weighs At. Mass in g
• 6.023x1023 MOLECULES = 1 mole OF MOLECULES
weighs Molar mass in g
• Suppose the molecule has formula A2B
then
• 1 mole of A2B consists of 6.023x1023 MOLECULES of A2B
and
• 2 X 6.023x1023 ATOMS of A
weighs 2 X At. Mass of A in g
• 1 X 6.023x1023 ATOMS of B
weighs 1 X At. Mass of B in g
Important Formulas to memorize
• Make a list!!