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Chapter 3- Molecular Shape and Structure
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What we are going to learn
How the number of bonds and lone pairs affect the geometry of molecules. (VSEPR)
How the geometries work with polarity to make a dipole moment.
Two theories of bondingValence Bond Theory: Atomic orbital hybridization.
Molecular orbital theory.
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VSEPR:Valence Shell Electron-Pair Repulsion
Basic Principle: Electrons are negatively charged, they want to stay as far away from each other as possible.
Electron pairs show _________________
Single bonds can be treated the same as
______________________________________________
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Some DefinitionsSteric Number: _________________________
Coordination Number: ______________________
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Some (more) DefinitionsElectron geometry: geometry including if you “saw” electron pairs
Molecular geometry: geometry where you don’t “see” electron pairs
Bond angle: angle between bonds.
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Pull out your worksheet, candy and toothpicks!
Thanks to http://www.chemmybear.com/shapes.html for all animations on following slides!
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Steric Number 2
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Steric Number 3
Examples:SO3, BF3, CO3
2-Examples:SO2, CCl2
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Steric Number 4
Examples:CH4, SiH4, PO4
Examples:NH3, PI3
Examples:H2O, OF2
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Steric Number 5
Examples:PCl5, SbF5
Examples:TeCl4, SF4
Examples:ClF3, SeO3
2-Examples:I3
-, XeF2
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Steric Number 6
Examples:SF6, Mo(CO)6
Examples:IF5, BrF5
Examples:XeF4, ClF4
-
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Going back to the Lewis Structures we did:
N2
CH2OBH3NH3
XeF4
N2O
SF6
BH3
H2SO4
POCl3ClF4
-
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Dipole MomentDipole moment: Vector addition of the magnetic moment of polar bonds
And that means…….
In one bond (2 atom) molecule if the bond is a polar the molecule is polar (aka has a dipole moment)
In multi-bond atoms you need to look at the ________________
Look at _____________________________Bonds in opposite directions _____________. Bonds in same direction _______________.
HF
Total Dipole Moment is…
ClF3
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Another Example: XeF2Cl2Two Possible Arrangements
Cl
F
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CF4+ CF4
-
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Going back to the Lewis Structures we did:
N2
CH2OBH3NH3
XeF4
N2O
SF6
BH3
H2SO4
POCl3ClF4
-
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Greenhouse Gases
Responsible for global warming.
Trap heat by absorbing IR radiation
Has permanent or induced dipole
Which of the following are greenhouse gases?N2, O2, CO, NO2, N2O
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Greenhouse Gases?
N2 O2
CO
NO2
N2O
C
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Two Theories of Bonding
Valence Bond Theory: Hybridization
Molecular Orbital Theory
Each atom retains its ___________
New atomic orbitals are formed from mix of ___________
Bonds form when _________________________
Atomic orbitals combine to make _______________
Molecular orbitals ______________
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H2 Bonding
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Valence Bond Theory: Types of Bondss bonds
“________________________”- symmetrical along bondAll single covalent bonds
p bondsOne p bond atom
Nodal Plane- Internuclear axis_________________“locks” rotation, _______________
Two p bond atom
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Sigma Bonds
Videos are available on the webpage
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p Bonds
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p Bonds
p
bond
s bonds bond
s
bond
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p bonds
C2H2
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Atomic Orbital HybridizationCombine the valence orbitals (______________) to make new orbitals
the orbitals are called ____, _______ and __________ orbitals are made from an ________________orbital
____ orbitals are made from an _______________orbitals
____ orbitals are made from an _______________orbitals
The new orbitals give the VSEPR geometry
p orbitals that are not hybridized still exist
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Atomic orbital hybridization sp
s
p p p
sp sp
p p
Things to notice:Combine two orbitals, get two new orbitalsCombine an s and a p get sp orbitalssp is the “name” of the orbital, just like s and p wereEnergy of sp orbital is between that of the s and the p orbital
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two sp orbitals and one left over p orbital
Atomic Orbital: sp another looks orbital
three p orbitalspx, py, pz
sp orbitals expanded out
linear
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Atomic orbital hybridization sp2
s
p p p
Things to notice:Combine three orbitals, get three new orbitalsCombine an s and two p orbitals to get sp2 orbitalssp2 is the “name” of the orbital, just like s and p wereEnergy of sp2 orbital is between that of the s and the p orbital and higher than an sp orbital
sp2 sp2
p
sp2
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three sp2 orbitals and one left over p orbital
Atomic Orbital: sp2 another looks orbital
three p orbitalspx, py, pz
sp orbitals expanded outtrig. planar
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Atomic orbital hybridization sp3
s
p p p
Things to notice:Combine four orbitals, get four new orbitalsCombine an s and three p orbitals to get sp3 orbitalssp3 is the “name” of the orbital, just like s and p wereEnergy of sp3 orbital is between that of the s and the p orbital and higher than an sp or sp2 orbital
sp3 sp3 sp3 sp3
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four sp3 orbitals
Atomic Orbital: sp3 another looks orbital
three p orbitalspx, py, pz
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Atomic Orbital Hybridization: including d orbitals
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How to find hybridizationDraw Lewis Structure
Figure out the arrangement of ___________________
Determine steric number: number bonded atoms+electron pairs
gives you number of hybrid orbitals needed
If you need:then
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Double and Triple Bonds in Hybridization
C2H4
Each carbon has three atoms bondedhybridization= give three sp2 orbitals,
each one is used to bondwhich orbital is left over for the double bond?
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Double and Triple Bonds in Hybridization
C2H4
sp2
use left over p orbital for double bond
single bonds and first bond in double/triple bonds are called σ bonds, second and third are called π bonds. Each π bond has two lobes
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Double and Triple Bonds in Hybridization
C2H4
What orbital does H use to bond with?
What orbital does C use to bond with H?
What orbital does C use to bond with C in the σbond?
What orbital does C use to bond with C π bond?
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Double and Triple Bonds in Hybridization
What is hybridization on each carbon?
______________________
______________________
What orbitals are left over?______________________
Use p orbitals to make two π bonds.
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ExamplesWhat type of hybridization is present in the following and what does each atom use to bond with and where are lone pairs located (and for fun, what is the geometries):
BF3
CH4
NH3
H2O
PCl5CH3CHO
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Another look at CH3CHO
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Another look at CH3CHO
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Another look at CH3CHO
Video available online
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Things to RememberHowever many unhybridized orbitals you start with, is how many you end with.
Hybrid orbitals are a combination of the atomic orbitals and belong to the atom and are still “atomic” orbitals
These hybrid orbitals overlap with each other to form bonds.
Each atom in a molecule can have its own (different) hybridization type.
Sigma bonds are made from hybrid orbitals in hybridized atoms
Pi bonds are made from left over p orbitals in hybridized atoms
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Going Back to the Lewis Structures we did…..again…..
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Molecular Orbital Theory_________________________ treatment of bonding
Orbitals belong to ___________ instead of individual _____________________________ of atomic orbitals
Adds or subtract _________ orbitals to get ___________ orbitals
Treatment yields better agreement with experiment and better prediction
So why don’t we use it for everything?
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Bonding and Anti-bonding OrbitalsIn hybridization, when we combined a # of atomic orbitals we got back that same number of hybrid atomic orbitals
In MO theory we’ll combine a ____ atomic orbitals and get _____ molecular orbitals. Two _________ orbitals give two __________ orbitals
One will be the __________ of the two orbitals- this is the bonding orbital
electron here _________ to the bond
One will be the __________ of the two orbitals- this is the antibonding orbital
electrons here ___________ from the bond.
***Remember an orbital is described by the wavefunction Ψ
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Wave Interference
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MO Theory
Electrons act as waves and particles
Waves have interference ____________.
If they interfere ______________ they add.
If the interfere ______________ they subtract
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MO Theory: s orbitals
E
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MO Theory
H2 He2 Li2 Be2
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Bond Order
Electrons in bonding orbitals ___________________________
Electrons in anti-bonding orbitals ________________________
Each electron adds or subtracts ½ a bond. Think about how this relates to what you know about electrons in a bond in valence bond theory/lewis structures.
Formula. Bond Order=
Roughly corresponds to single, double, triple bonds. bond orders of decimal points____________________.
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MO Theory
H2 He2 Li2 Be2
BO
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MO Theory: p orbitals
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Steps to make MO diagram
Count valence electrons and fill into atomic orbital diagram along sides.
Decide on which order the energy diagram should follow.
Total valence electrons from each atom, fill into diagram from low energy to high
Pauli Exclusion, Hund’s and Aufbau principles still apply
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MO Theory
B2
BO
C2 N2 O2 F2 Ne2
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Molecular Orbital Diagrams of Ions
Add (__________) or subtract (________) electrons as needed.
When adding or subtracting electrons from atom orbitals obey Pauli, Aufbau and Hund’s rules are followed.
Add to Molecular orbitals in similar fashion
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Molecular Orbital Diagrams of Ions
N2N2- N2
+N22-
BO
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Case of Molecular Oxygen
O2
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Practice Use MO theory to explain or decide the following
Why don’t noble gasses form diatomic molecules?
Which species has the longer bond length N2 or O2?
For carbon monoxide, the porbitals combine such that the order of energy is π, σ,π*, σ*. Draw the MO diagram. Whatsthe bond order?
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Thing to double check when drawing MO diagrams
Did you label all the atomic orbitals?
Did you label all the molecular orbitals?
Did you remember to add the * to anti-bonding orbitals?
Did you add or subtract electrons as appropriate if you have an ion?
Did you use the proper order of orbitals for the MOs?
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Heteronuclear DiatomicsInstead of the atomic orbitals adding equally, one atom’s wave function adds disproportionately.
2 gives orbital
___________ in non polar covalent bond, everything is shared equally.
__________Orbital is more B-like than A-like
Happens in polar covalent bond
e.g. CO or NO
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Polyatomic Atoms
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A word on what you need to be able to do for the exam
Homonuclear diatomics: everything. I should be able to give you a blank sheet of paper and ask you to draw all of the MO diagrams and you need to be able to do it complete with labels.
Heteronuclear diatomics: I draw energy levels, you make labels and fill in electrons.
Triatomic molecules: I’ll draw and label, you simply will need to be able to fill in the electrons. (I’ll use the “box style notation when drawing so you know how many electrons can be held in orbital”)
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Revisit: The grass is green because…chlorophyll absorbs red light andreflects green light
eyes see reflected green light
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Chapter 5.1-5.6The rest will be covered in 1B
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What is an intermolecular force?
inter=
molecular=
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Types of Forces
Intramolecular forces are bonds ______________, i.e. covalent bonds
Intermolecular forces are much weaker.
Dipole-Dipole Forces
Hydrogen Bonds (a type of very strong dipole-dipole force)
Ion-Dipole Forces
Dispersion Forces
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Dipole-Dipole ForcesAttractive forces between
Positive side of one molecule attracts negative side of another molecule.
Negative side of one molecule attracts positive side of another molecule.
____________ of intermolecular forces (except hydrogen bonds which is a special type of dipole dipole force)
_________________as dipole of molecule increases
Present in molecules with a __________
Example: HCl
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Hydrogen BondsVery strong dipole-dipole interaction, ______________ of intermolecular forces
Occurs when H is _________________ to a highly electronegative atom (N, O, F)
Still occurs between two different molecules
The more hydrogen bond donors and acceptors the higher the difference in properties
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Video Time
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The more hydrogen bond donors and acceptors the higher the difference in properties
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Ion Dipole Forces
An _______ interacts with ___________________
___ ion draws the ____side of the ion toward it
___ ion draws the ___ side of the ion toward it
OR
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Dispersion ForcesWeakest of intermolecular forces
All molecules possess these
Responsible for non-ideal behavior in gasses
aka- induced dipole-induced dipole: aka London forces
Increases with increasing electronsThis means usually increases with increasing molar mass
One molecule induces a dipole in another, it induces a dipole in the next and so on……
Example: F2
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Gecko Feet
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“Van der Waals Forces”
Umbrella term for Dipole-dipole, dipole-induced dipole and dispersion forces.
Sometimes used as an umbrella term for all intermolecular forces
Shouldn’t be used when I ask for “Which molecular forces are present in this molecule?”
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Summary of Intermolecular Force Strengths and Molecules Having Interactions
Hydrogen BondingMolecules with H bonded to N, O or F (remember H-bond is still INTER-molecular)
Dipole-DipoleMolecules with a dipole
Dispersion forcesAll molecules- increases with increasing # of electrons
Strong
Weaker
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Going back to the Lewis Structures we did: What intermolecular forces are present:
N2
CH2OBH3NH3
XeF4
N2O
SF6
BH3
H2SO4
POCl3ClF4
-
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Effects on Boiling and Melting Points
Reminder: What happens when something Boils?
Reminder: What happens when something Melts?
If you increase the strength of the forces between molecules what is going to happen?
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Melting
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Boiling Melting
If you increase the strength of the forces between molecules what is going to happen? Boiling and melting point increase
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Boiling Point
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Boiling Point/Melting Point ExamplesRank the following by increasing Boiling and Melting point name the forces each have:
Ar, He, Ne, Xe
PH3, N2, CH4, H2O
Explain the followingBr2 has a lower melting point than NaBr
C2H5OH has a higher boiling point than butane, C4H10
H2O has a higher boiling point than H2Te
Acetic acid has a lower boiling point than benzoic acid
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Does this mean all H bonded molecules have higher bp/mp than non H bonded?
CH3F, HO-OH, NH3, H3C-O-CH3
Propanol, Dodecane, Propane
Lets look at some examples: Put in order of increasing boiling point.
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Hydrogen Bonding: Importance
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Hydrogen Bonding: Importance
http://idav.ucdavis.edu/~okreylos/ResDev/ProtoShop/ScreenShots.html
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Hydrogen Bonding Importance