Unit 3: Chemical Bonding and Molecular Structure Cartoon courtesy of NearingZero.net.

Unit 3: Unit 3: Chemical Chemical BondingBonding

and and MolecularMolecular Structure Structure

Cartoon courtesy of NearingZero.net

The The OctetOctet Rule RuleChemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.

Diatomic Fluorine

Hydrogen Chloride by the Octet Hydrogen Chloride by the Octet RuleRule

Formation of Water by the Octet Formation of Water by the Octet RuleRule

Comments About the Octet RuleComments About the Octet Rule

 2nd row elements C, N, O, F observe the octet rule.

 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.

 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.

 When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

Multiple Covalent Bonds:Multiple Covalent Bonds:Double bondsDouble bonds

Two pairs of shared electrons

Multiple Covalent Bonds:Multiple Covalent Bonds:Triple bondsTriple bonds

Three pairs of shared electrons

Bond Dissociation Energy

 It is the energy required to break a bond.  It gives us information about the strength of a bonding interaction.

ResonancResonancee  Occurs when more than one valid

Lewis structure can be written for a particular molecule.

  These are resonance structures. The actual structure is an average of the resonance structures.

Resonance in OzoneResonance in Ozone

Neither structure is correct.

HybridizationHybridization

The Blending of OrbitalsThe Blending of Orbitals

We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds.

Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it.

Lets look at amolecule of methane, CH4.

What is the expected orbital notation of carbon in its ground state?

(Hint: How many unpaired electrons does this carbon atom have available for bonding?)

Can you see a problem with this?

Carbon ground state configuration

You should conclude You should conclude that carbon only has that carbon only has TWOTWO electrons electrons available for bonding. available for bonding. That is not not enough!That is not not enough!

How does carbon overcome this problem so that How does carbon overcome this problem so that it may form four bonds?it may form four bonds?

Carbon’s Bonding Carbon’s Bonding ProblemProblem

The first thought The first thought that chemists had that chemists had was that carbon was that carbon promotes one of its promotes one of its 2s2s electrons… electrons…

…to the empty 2p orbital.

Carbon’s Empty Carbon’s Empty OrbitalOrbital

However, they quickly recognized a problem with such an arrangement…

Three of the carbon-hydrogen bonds would involve an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.

A Problem Arises

This would mean that three of the bonds in a methane molecule would be identical, because they would involve electron pairs of equal energy.

But what about the fourth bond…?

Unequal bond energy

The fourth bond is between a 2s electron from the carbon and the lone 1s hydrogen electron.

Such a bond would have slightly less energy than the other bonds in a methane molecule.

Unequal bond energy #2

This bond would be slightly different in character than the other three bonds in methane.

This difference would be measurable to a chemist by determining the bond length and bond energy.

But is this what they observe?

Unequal bond energy #3

The simple answer is, “No”.

Chemists have proposed an explanation – they call it Hybridization.

Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy.

Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane.

Enter Hybridization

In the case of methane, they call the hybridization sp3, meaning that an s orbital is combined with three p orbitals to create four equal hybrid orbitals.

These new orbitals have slightly MORE energy than the 2s orbital…

… and slightly LESS energy than the 2p orbitals.

sp3 Hybrid Orbitals

Here is another way to look at the sp3 hybridization and energy profile…

sp3 Hybrid Orbitals

While sp3 is the hybridization observed in methane, there are other types of hybridization that atoms undergo.

These include sp hybridization, in which one s orbital combines with a single p orbital.

Notice that this produces two hybrid orbitals, while leaving two normal p orbitals

sp Hybrid Orbitals

While sp3 is the hybridization observed in methane, there are other types of hybridization that atoms undergo.

These include sp hybridization, in which one s orbital combines with a single p orbital.

Notice that this produces two hybrid orbitals, while leaving two normal p orbitals

sp Hybrid Orbitals

Another hybrid is the sp2, which combines two orbitals from a p sublevel with one orbital from an s sublevel.

Notice that one p orbital remains unchanged.

sp2 Hybrid Orbitals

VSEPR Model

 The structure around a given atom is determined principally by minimizing electron pair repulsions.

(Valence Shell Electron Pair Repulsion)

Predicting a VSEPR StructurePredicting a VSEPR Structure

  Draw Lewis structure. Draw Lewis structure.

   Put pairs as far apart as possible. Put pairs as far apart as possible.

   Determine positions of atoms from Determine positions of atoms from the way electron pairs are shared. the way electron pairs are shared.

   Determine the name of molecular Determine the name of molecular structure from positions of the structure from positions of the atoms.atoms.

VSEPR and the water molecule

VSEPR and the ammonia molecule

Polar-Covalent bonds

Nonpolar-Covalent bonds

Covalent BondsCovalent Bonds

Electrons are unequally shared Electronegativity difference between .3 and 1.7

Electrons are equally shared Electronegativity difference of 0 to 0.3

PolarityPolarity

A molecule, such as HF, that has a A molecule, such as HF, that has a center of positive charge and a center of positive charge and a center of negative charge is said to center of negative charge is said to be polar, or to have a dipole moment.be polar, or to have a dipole moment.

+

FH

Relative magnitudes of Relative magnitudes of forcesforces

The types of bonding forces vary in The types of bonding forces vary in their strength as measured by their strength as measured by average bond energy.average bond energy.

Covalent bonds (400 kcal)

Hydrogen bonding (12-16 kcal )

(Van der Waals Forces) Dipole-dipole interactions (2-0.5 kcal)

London forces (less than 1 kcal)

Strongest

Weakest

Hydrogen Hydrogen BondingBonding

Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests.

Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen

Hydrogen Hydrogen Bonding Bonding in Waterin Water

Hydrogen Bonding between Ammonia and Water

Dipole-Dipole-Dipole Dipole

AttractionsAttractions

Attraction between oppositely charged regions of neighboring molecules.

The water dipole

The ammonia dipole

London Dispersion London Dispersion ForcesForces

The temporary The temporary separations of charge separations of charge that lead to the London that lead to the London force attractions are force attractions are what attract one what attract one nonpolarnonpolar molecule to its molecule to its neighbors.neighbors.

Fritz London Fritz London 1900-19541900-1954

London forces increase London forces increase with the size of the with the size of the molecules.molecules.

London Forces in HydrocarbonsLondon Forces in Hydrocarbons