Chapter 3- Molecular Shape and...

Chapter 3- Molecular Shape and Structure

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.

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

______________________________________________

Some DefinitionsSteric Number: _________________________

Coordination Number: ______________________

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.

Pull out your worksheet, candy and toothpicks!

Thanks to http://www.chemmybear.com/shapes.html for all animations on following slides!

Steric Number 2

Steric Number 3

Examples:SO3, BF3, CO3

2-Examples:SO2, CCl2

Steric Number 4

Examples:CH4, SiH4, PO4

Examples:NH3, PI3

Examples:H2O, OF2

Steric Number 5

Examples:PCl5, SbF5

Examples:TeCl4, SF4

Examples:ClF3, SeO3

2-Examples:I3

-, XeF2

Steric Number 6

Examples:SF6, Mo(CO)6

Examples:IF5, BrF5

Examples:XeF4, ClF4

-

Going back to the Lewis Structures we did:

N2

CH2OBH3NH3

XeF4

N2O

SF6

BH3

H2SO4

POCl3ClF4

-

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

Another Example: XeF2Cl2Two Possible Arrangements

Cl

F

CF4+ CF4

-

Going back to the Lewis Structures we did:

N2

CH2OBH3NH3

XeF4

N2O

SF6

BH3

H2SO4

POCl3ClF4

-

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

Greenhouse Gases?

N2 O2

CO

NO2

N2O

C

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 ______________

H2 Bonding

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

Sigma Bonds

Videos are available on the webpage

p Bonds

p Bonds

p

bond

s bonds bond

s

bond

p bonds

C2H2

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

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

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

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

three sp2 orbitals and one left over p orbital

Atomic Orbital: sp2 another looks orbital


sp orbitals expanded outtrig. planar

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

four sp3 orbitals

Atomic Orbital: sp3 another looks orbital


Atomic Orbital Hybridization: including d orbitals

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

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?


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


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?


What is hybridization on each carbon?

______________________

______________________

What orbitals are left over?______________________

Use p orbitals to make two π bonds.

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

Another look at CH3CHO

Another look at CH3CHO

Video available online

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

Going Back to the Lewis Structures we did…..again…..

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?

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

Wave Interference

MO Theory

Electrons act as waves and particles

Waves have interference ____________.

If they interfere ______________ they add.

If the interfere ______________ they subtract

MO Theory: s orbitals

E

MO Theory

H2 He2 Li2 Be2

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____________________.

MO Theory

H2 He2 Li2 Be2

BO

MO Theory: p orbitals

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

MO Theory

B2

BO

C2 N2 O2 F2 Ne2

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

Molecular Orbital Diagrams of Ions

N2N2- N2

+N22-

BO

Case of Molecular Oxygen

O2

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?

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?

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

Polyatomic Atoms

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”)

Revisit: The grass is green because…chlorophyll absorbs red light andreflects green light

eyes see reflected green light

Chapter 5.1-5.6The rest will be covered in 1B

What is an intermolecular force?

inter=

molecular=

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

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

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

Video Time

The more hydrogen bond donors and acceptors the higher the difference in properties

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

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

Gecko Feet

“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?”

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

Going back to the Lewis Structures we did: What intermolecular forces are present:

N2

CH2OBH3NH3

XeF4

N2O

SF6

BH3

H2SO4

POCl3ClF4

-

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?

Melting

Boiling Melting

If you increase the strength of the forces between molecules what is going to happen? Boiling and melting point increase

Boiling Point

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

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.

Hydrogen Bonding: Importance

Hydrogen Bonding: Importance

http://idav.ucdavis.edu/~okreylos/ResDev/ProtoShop/ScreenShots.html

Hydrogen Bonding Importance

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