READY

Molecular Shapes + VSEPR Model

• Objectives:
• SW apply the VSEPR Model of molecules to determine and describe the shapes of molecules.

​

Molecular Shapes

• Lewis structures are 2 Dimensional
• Accurately shows shape for 2-Dimensional molecules (linear, bent, planar)
• DOESN’T accurately show shape for 3-Dimensional molecules

Example: Methane (CH₄)

• Lewis Structure Molecular Shape

Terminology

• Bond Angles: the angle made by the lines joining the nuclei of the atoms in a molecule

​

Terminology

• Electron Domains: Regions in space where electrons are likely to be found
• Bonding electrons: Electrons that are shared between atoms. The region between the atoms is an electron domain
• Nonbonding pair of electrons: Electrons that are located principally on just one of the atoms. This also represents an electron domain
• In this example Nitrogen has
• 3 bonding domains
• 1 nonbonding domain
• 4 electron domains total

VSEPR: Valence Shell Electron Pair Repulsion Model

• Electrons repel each other, so the geometry around an atom is going to be determined by minimizing these repulsions between electron domains

A

A

AXE notation

• Shapes around a central atom are determined by how many electron domains there are around it.
• AXE notation will be a helpful way to keep track of electron domains and the resulting shapes
• A: The central atom
• X: number of atoms bonded to the central atom
• E: the number of unshared pairs of electrons
• AX₃E₁ : 3 atoms bonded to a central atom that also has 1 pair of unshared electrons on it
• To start, lets just look at examples with ONLY bonding domains

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

Organize Your Notes

• https://docs.google.com/document/d/1n68GMRalD13sRNiKZp8GeYSzb27Q4ZTrOSrxmRobl98/edit?usp=sharing

AX₁E₀

• Lewis Molecular Bond Angles

Electronic Geometry

A - X

H - H

180^o

Linear

AX₂E₀

• Lewis Molecular Bond Angles

Electronic Geometry

X - A - X

O = C = O

180^o

Linear

AX₃E₀

• Lewis Molecular Bond Angles

Electronic Geometry

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

120^o

Trigonal Planar

AX₄E₀

• Lewis Molecular Bond Angles

Electronic Geometry

109.5^o

Tetrahedral

AX₅E₀

• Lewis Molecular Bond Angles

Electronic Geometry

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

120^o

90^o

Trigonal Bipyramidal

PCl₅

AX₆E₀

• Lewis Molecular Bond Angles

Electronic Geometry

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

90^o

Octahedral

SF₆

Electronic Geometries vs Molecular Geometries

• Examples so far have been when electronic and molecular geometries have been the same
• What shape do the electronic domains make
• Molecular geometries can be different
• What shape do just the bonded ATOMS make
• ‘ignore’ the unshared pairs of electrons
• No unshared electrons means electronic geometry is the same as the molecular
• The total number of domains still dictate the overall shape around the atom, but the molecular shape just describes the shape the atoms make.

Electronic Geometry

Tetrahedral

​

Molecular Geometry

Bent

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

Electronic vs. Molecular Geometries

• Electronic Geometry Molecular Geometry

Tetrahedral Bent

AX₁E₁: 2 Domains; 1 bonding 1 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

N₂

180^o

Linear

Molecular Geometry

Linear

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₂E₁: 3 Domains: 2 bonding, 1 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

SO₂

120^o

Trigonal Planar

Molecular Geometry

Bent

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₁E₂: 3 Domains: 1 bonding, 2 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

O₂

180^o

Linear

Molecular Geometry

Trigonal Planar

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₃E₁: 4 Domains: 3 bonding, 1 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

NH₃

109.5^o

Tetrahedral

Molecular Geometry

Trigonal Pyramidal

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₂E₂: 4 Domains: 2 bonding, 2 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

H₂O

109.5^o

Tetrahedral

Molecular Geometry

Bent

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₁E₃: 4 Domains: 1 bonding, 3 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

H-Cl

180^o

Linear

Molecular Geometry

Tetrahedral

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

NOTE: Electrons are Space Hogs

• They want to occupy more space than bonded atoms
• Should they go where they have 90^o or 120^o of room?
• They will take the 120^o spot
• They also REPEL the bonded atoms and distort the bond angles (when they aren’t symmetrically around the central atom)

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₄E₁: 5 Domains: 4 bonding, 1 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

SF₄

120^o

90^o

Trigonal Bipyramidal

Molecular Geometry

Seesaw

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₃E₂: 5 Domains: 3 bonding, 2 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

ClF₃

90^o

Molecular Geometry

Trigonal Bipyramidal

T-Shaped

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₂E₃: 5 Domains: 2 bonding, 3 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

XeF₂

180^o

Molecular Geometry

Trigonal Bipyramidal

Linear

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₅E₁: 6 Domains: 5 bonding, 1 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

BrF₅

90^o

Octahedral

Molecular Geometry

Square Pyramidal

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₄E₂: 6 Domains: 4 bonding, 2 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

XeF₄

90^o

Octahedral

Molecular Geometry

Square Planar

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₄E₂: 6 Domains: 3 bonding, 3 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

90^o

Octahedral

Molecular Geometry

T- Shaped

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

AX₄E₂: 6 Domains: 2 bonding, 4 nonbonding

• Lewis Molecular Bond Angles

Electronic Geometry

180^o

Octahedral

Molecular Geometry

Linear

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

Try Some Practice

• Draw the Lewis structures for the following and determine their electronic AND molecular geometries!
• BrF₃ Electronic Molecular Bond Angles

​

​

​

• NO₂^-1 Electronic Molecular Bond Angles

​

Shapes of Larger

• Often times we will be asked the geometry around a specific atom in a more complex molecule
• Take the same approach and isolate the central atom
• CH₃COOH

Summarize

Can you

• Apply the VSEPR Model of molecules to determine and describe the shapes of molecules

READY

Molecular Shapes and polarity

Objectives:

• SW determine the polarity of molecules based on their shapes determined by the VSEPR model.

​

Electronegativity and Bond Polarity

• Electronegativity: an atoms attraction to electrons in a bond
• The higher the electronegativity; the greater the attraction
• ‘How much they ‘hog’ the electrons being shared’
• Bond Polarity
• When electrons are unevenly shared between two atoms, they create a polar bond
• Since electrons ‘hang around’ one of the atoms more than the other, that atom becomes partially negative

2.20

2.20

2.20

3.16

Electronegativity and Dipole Moments

• Bond Dipole: a measure of the separation between the positive and negative charges in a bond
• These are vector quantities; they have a magnitude AND a direction
• Point arrow towards more electronegative atom
• Other end gets a line through it (looks like a + sign)

​

​

2.20

2.20

2.20

3.16

Electronegativity and Dipole Moments

• Dipole Moment: a measure of the separation between the positive and negative charges in a molecule
• It is the SUM of the bond dipoles in a molecule
• If they cancel each other out (Dipole moment = 0) then the molecule is nonpolar
• If they don’t cancel out, you have a polar molecule

2.20

3.16

Example: CO₂

2.55

3.44

3.44

Example: CO₂ Non Polar

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

Example: H₂O

2.20

2.20

3.44

Example: H₂O

PhET Interactive Simulations, University of Colorado Boulder, https://phet.colorado.edu.

Try these ones! Is it a polar or nonpolar molecule?

• CCl₄ BF₃

​

​

​

• CH₃Cl BrF₅

​

A 2-D trick

• If the Lewis structure has more than 1 line of symmetry, its nonpolar
• CCl₄ BF₃

​

​

​

• CH₃Cl BrF₅

​

Summarize

Can you

• Determine the polarity of molecules based on their shapes determined by the VSEPR model.

​

• Electronegativity and bond polarity review
• Dipole moment and charge separation
• Charge separation affects the properties of molecules (hint to attractive forces and increased BP)
• Molecular polarity: depends on the bond polarity and geometry of the molecule
• Dipoles are vector quantities: depends on magnitude AND direction
• CO2 example of polar bonds but nonpolar molecule. No overall dipole moment because of their opposing directions
• H2O example of how they would add together
• Are there polar bonds?
• Are they asymmetrically placed?
• POLAR MOLECULE!

​

READY

Covalent Bonding and Orbital Overlap 

Objectives:

• Describe bonding between atoms in terms of orbital overlap

​

Orbital Overlap

• VSEPR Model predicts the shape well BUT…
• It doesn’t explain WHY atoms bond
• Valence Bond Theory: Combination of Lewis’ electron pair bonding AND atomic orbitals
• When a valence atomic orbital of one atom ‘merges’ and overlaps with another
• Two electrons of opposite spin share that space where they overlap
• Electrons are shared

​

H-H

• 1s¹ 1s¹

H-Cl

• H 1s¹ Cl: [Ne]3s² 3p⁵

Bond Distance and Energy

Energy (kJ/mol)

Bond Distance

0

Nuclei attraction to other atoms electrons > repulsive forces

Attractive and repulsive forces balanced.

This is the bond length

Nuclei get close and repel each other

Bond Length

• The distance where the attractive forces (Between Nuclei and electrons) are balanced by the repulsive forces (electron-electron and nucleus-nucleus)

Summarize

Can you

• Describe bonding between atoms in terms of orbital overlap.

​

​

• VSEPR models HOW they bond, but not WHY they bond
• Valence Bond Theory
• Electrons are ‘shared’ when orbitals overlap
• Two electrons of opposite spin share that space where they overlap
• Visual
• Nuclei attracted to other atoms electrons
• Forms a bond
• H₂ and HCl orbital diagrams and orbital overlaps. Visual with box diagrams
• Potential energy and distance chart and balance of attraction and repulsion

​

READY

Hybrid orbitals

Objectives:

• Describe what happens during hybridization and determine the hybridization present in various molecules

Valence-Bond Theory Shortcomings

• The geometries of atomic orbitals and the shapes that molecules make when bonding isn’t consistent.
• Nitrogen has 3 unpaired p electrons, with 90^o of separation between them
• Observed bond angles for NH₃ ARE NOT 90^o
• There are atoms that make more bonds than they have unpaired electrons
• Sulfur makes 6 bonds in SF₆ but only has 2 unpaired electrons

Hybrid Orbitals

• In order to account for Bonding AND Geometry, we need a new model
• Hybrid Orbital Theory: Atomic orbitals of an atom MIX and form new orbitals called hybrid orbitals.
• Hybridization: The process of mixing and, as a result, changing atomic orbitals as they approach each other to form bonds
• TOTAL NUMBER OF ATOMIC ORBITALS DOESN’T CHANGE
• 2 atomic orbitals mixed together 🡪 2 Hybrid orbitals
• 3 atomic orbitals mixed together 🡪 3 Hybrid orbitals

​

Hybrid Orbitals

• We Describe the hybridization of the orbitals by what went into creating them:
• A single ‘s’ orbital and a single ‘p’ orbital will hybridize to make TWO ‘sp’ hybrid orbitals
• A single ‘s’ orbital and two ‘p’ orbitals will hybridize to make THREE ‘sp²’ hybrid orbitals
• sp³: one s and three p orbitals
• sp³d
• sp³d²
• Paired Electrons can be ‘unpaired’ by being promoted to another orbital and end up unpaired in hybrid orbitals

s

p

sp

sp

Notes Organizer

sp Hybridization

• s + p 🡪 sp x2
• A single ‘s’ orbital and a single ‘p’ orbital hybridize to get two ‘sp orbitals’
• An ‘s’ electron gets promoted to a ‘p’ orbital and then they hybridize
• Shape of the orbital is similar but different to the s and the p orbitals
• 2 lobes
• One significantly longer than the other
• Better for overlapping with other orbitals
• Take on a linear geometry

​

2s

2p

2 sp

2p

Only large lobes:

1s

1s

Unhybridized

sp² Hybridization

• s + p + p 🡪 sp² x3
• Three sp² hybrid orbitals
• Trigonal Planar geometry (120^o between each)

2s

2p

sp²

2p

sp³ Hybridization

• s + p + p + p
• 4 sp³ hybrid orbitals
• 4 domains takes on a Tetrahedral shape

2s

2p

sp³

NOTE: Unshared pairs of electrons ALSO need a hybrid orbital

• Example: Oxygen in H₂O
• Expectation based on unhybridized orbitals
• p orbitals are at 90^o from each other
• If it was simple orbital overlap bond angles would be the same
• Actual bond angles reflect tetrahedral electronic geometry

2s

2p

NOTE: Unshared pairs of electrons ALSO need a hybrid orbital

2s

2p

O:

1s

1s

sp³

sp³d Hybridization

• s + p + p + p + d =
• 5 sp³d hybrid orbitals
• Forms a Trigonal Bipyramidal geometry

s

p

d

d

sp³d

sp³d² Hybridization

• s + p + p + p + d + d =
• 6 sp³d² hybrid orbitals
• Forms an Octahedral geometry

​

s

p

d

sp³d²

d

How to Determine What The Hybridization Is?

• 1) Draw the Lewis Structure
• 2) Determine the electron domain geometry using the VSEPR model
• Electron domains = Bonded atoms + unshared pairs of electrons
• Need 1 hybrid orbital for EACH domain
• 3) Determine which hybridization gives you that many hybrid orbitals

Practice: Predict the hybridization and electron geometry for each of the following

• BeF₂ BF₃ N₂

​

​

​

​

• CH₄ SF₄ SF₆

​

Summarize

Can You

• SW describe what happens during hybridization and determine the hybridization present in various molecules

READY

Multiple Bonds and Hybridization

Objectives:

• Describe multiple bonds (double and triple) in terms of sigma and pi bonds and the properties of such bonds

Single Bonds

• Electrons are being shared along the internuclear axis
• If you were to connect a line between the two nuclei, the shared electrons would be found along that line, in between the two nuclei.
• We call these kinds of bond sigma bonds (σ)

sp

sp

Single Bonds

• BH₃
• sp² hybrid orbitals
• Internuclear axis
• Sigma Bonds (σ)
• (a single bond is almost ALWAYS a sigma bond)

(σ)

(σ)

(σ)

Sigma Bonds can Spin

• Spinning in sigma bonds is allowed because as an atom spins there is still orbital overlap between them
• Molecules can ‘rearrange’ their 3D structure

Multiple Bonds (double and triple)

• Involve a different kind of bond
• Pi bonds (π): covalent bond that involves overlap of orbitals above and below the internuclear axis.
• The shared electrons wont be found between the two nuclei
• Results from the side-to-side overlap of unhybridized ‘p’ orbitals
• Can have 2 pi bonds total
• Overlap ‘above and below’
• Overlap ‘on the left and right’

​

π

π

Double Bond: C₂H₄

• Carbon:
• 3 electronic domains
• 3 sp² hybridized orbitals
• 1 unhybridized p orbital
• Bonding
• 1 sigma bond from overlapping sp² hybrid orbitals
• 1 pi bond from overlapping p orbitals
• PLANAR Molecule

2s

2p

sp²

2p

σ

π

Rigidity of pi Bonds

• pi bonds CANT SPIN
• pi bonds don’t occur along the internuclear axis
• p orbitals need to line up ‘above and below’ the axis
• If the atoms were to spin, the p orbitals wouldn’t overlap and the pi bond would need to break
• pi bonds are rigid and don’t allow for the rotation of the atoms involved

σ

π

Triple Bonds: C₂H₂

• Carbon:
• 2 electronic domains
• 2 sp hybridized orbitals
• 2 unhybridized p orbital
• Bonding
• 1 Sigma bond
• 2 pi bonds
• Left and right overlap
• Above and below overlap
• LINEAR shape

​

π

π

π

π

In General

• Double bonds:
• 1 sigma bond
• 1 pi bond
• Since you need an unhybridized p orbital, expect sp or sp² hybridization
• Triple bonds
• 1 sigma bond
• 2 pi bonds
• Since you need 2 unhybridized p orbitals, expect sp hybridization

​

Summarize

Can You:

• Describe multiple bonds (double and triple) in terms of sigma and pi bonds and the properties of such bonds

​

READY

Delocalized Electrons

Objectives

• Describe what delocalized electrons are, their effect on the structure of the molecule and determine when they occur.

​

Localized Electrons

• Localized electrons:
• electrons that are found in a specific region
• In this example, the bonding electrons ARE localized
• You will find the sigma bonding electrons along the internuclear axis
• You will find the pi bonding electrons “above and below” the internuclear axis

σ

π

Delocalized Electrons

• Sometimes the electrons ARENT localized:
• Delocalized electrons aren’t found shared just between 2 atoms, but MULTIPLE atoms and locations
• When does this occur?
• When a molecule has multiple resonance structures involving pi bonds

​

Ex: Benzene

• Carbon is sp² hybridized
• Trigonal planar with unhybridized p orbital

p_y

sp²

sp²

sp²

Ex: Benzene’s Resonance Structures

• Benzene has 2 equivalent resonance structures

Ex: Benzene’s Resonance Structures

• Benzene’s actual structure isn’t either of one of them
• If it was alternating single and double bonds with localized electrons, you would expect to see different bond distances
• Observed structure: ALL the bond distances are the same
• Delocalized Electrons can account for this difference

“Expected”

Observed

1.34 Å

1.54 Å

1.54 Å

1.54 Å

1.34 Å

1.34 Å

1.40 Å

1.40 Å

1.40 Å

1.40 Å

1.40 Å

1.40 Å

Ex: Benzene’s Resonance Structures

• Delocalized electrons: Electrons that aren’t localized between 2 atoms
• The electrons move between >2 atoms
• In benzene, they move between ALL 6 carbons in the regions above and below the molecule, where the pi orbitals would overlap
• Instead of being shared between 2 atoms, the ‘p’ electrons are shared among >2

Bond Lengths Reflect Delocalized Electrons

• The Carbons share their p electrons (in the pi bonds) with ALL of the carbons in Benzene
• Uniform bond lengths observed

Observed

1.40 Å

1.40 Å

1.40 Å

1.40 Å

1.40 Å

1.40 Å

Resulting Properties

• Delocalized electrons lead to:
• Increased stability of the molecule
• Instead of ‘messing’ up a bond for just 2 atoms, you’re ‘messing’ it up for more
• Colors of organic molecules

Are There Delocalized Electrons?

• Are there multiple resonance structures?
• Do they involve pi bonds?
• DELOCALIZED ELECTRONS!

Are There Delocalized Electrons?

• H₂CO O₃

​

​

​

• NO₃^-

​

​

​

​

Summarize

Can you

• Describe what delocalized electrons are, their effect on the structure of the molecule and determine when they occur.

​

​

READY

Molecular Orbital Theory

Objectives:

• Explain bonding in terms of molecular orbital theory.
• Determine bond order and draw molecular orbital diagrams for diatomic molecules

​

Molecular Orbital Theory

• Valence Bond Theory
• Good model for
• Bonding, Geometries
• NOT a good model for
• explaining excited sates of molecules, how molecules absorb light and are given color
• NEED A DIFFERENT MODEL for bonding to explain theses things
• Models: If they aren’t accurate, why do we use them?
• A Globe
• A good model for:
• What the earth looks like in 3d, Where different countries are in relation to each other
• NOT a good model for
• Driving directions, Trail maps

​

Molecular Orbital Theory

• Atomic Orbitals: how electrons exist (wave function) around a nucleus
• Molecular Orbital (MO): how electrons exist (wave function) on a molecule
• The electrons aren’t thought of as orbiting a single atom
• Orbiting the whole molecule

Atomic Orbitals

• On just 1 atom

BOTH

• Wave functions
• 2 electrons per orbital (opposite spins)
• Definite amount of energy

​

​

Molecular Orbitals

• Around the WHOLE molecule

Hydrogen Molecule (H₂)

• When 2 atomic orbitals overlap, 2 molecular orbitals are created
• 1 Bonding (lower energy):
• Orbitals merge together. Electron density in between nuclei
• Lower energy because electron is attracted to BOTH nuclei
• 1 Anti bonding (higher energy)
• Electrons are repelled from each other;
• very little electron density between the nuclei

​

‘Naming’ Molecular orbitals

• Describe where the electrons are and what the MO’s are made of
• Sigma: along the internuclear axis
• Bonding and antibonding
• Made from 1s atomic orbitals
• Here we have made
• σ_1s ^: “Sigma One s” (Bonding)
• σ^*_1s : “Sigma-star One S” (Antibonding)

​

​

​

​

​

σ_1s

σ^*_1s

Molecular Orbital Diagrams (H₂)

• Atomic Orbitals on left and right side
• Molecular orbitals in the middle
• Energy increases vertically
• Pauli’s Exclusion Principle still applies
• Place the electrons into the MO’s starting with the lowest energy ones

1s

1s

σ_1s

σ^*_1s

H

H

Practice: Draw the Molecular Orbital Diagram for He₂

• Atomic Orbitals on left and right side
• Molecular orbitals in the middle
• Energy increases vertically
• Pauli’s Exclusion Principle still applies
• Place the electrons into the MO’s starting with the lowest energy ones

​

1s

1s

σ_1s

σ^*_1s

He

He

Practice: Draw the Molecular Orbital Diagram for He₂⁺

• Atomic Orbitals on left and right side
• Molecular orbitals in the middle
• Energy increases vertically
• Pauli’s Exclusion Principle still applies
• Place the electrons into the MO’s starting with the lowest energy ones

​

1s

1s

σ_1s

σ^*_1s

He

He

Bond Order

• Bond order = ½ (#bonding electrons – antibonding electrons)
• Describes the stability of the covalent bond (# of shared pairs overall)
• Single bonds; Bond Order = 1
• Double bonds: B.O. = 2
• Triple bonds: B.O. = 3
• Bond order CAN be fractions

Bond Order for H₂

• B.O. = ½(bonding e^- - antibonding e^-)
• = ½ ( 2 – 0 )
• Bond Order = 1
• Single bond
• Consistent with what we know about H₂

​

Bond Order for He₂

• B.O. = ½(bonding e^- - antibonding e^-)
• = ½ ( 2 – 2 )
• Bond Order = 0
• No Bond Exist.

Note:

• Bonding Molecular Orbitals are at a lower energy than atomic orbitals
• Favored energetically
• Antibonding Molecular orbitals are at a HIGHER energy than atomic orbitals
• This higher energy OFFSETS the energy of the bonding MO
• When bonding = antibonding, total energy is slightly higher than individual atomic orbitals, so better off NOT bonding
• B.O. = 0 No Bond exists

​

Bond Order for He₂⁺

• B.O. = ½(bonding e^- - antibonding e^-)
• = ½ ( 2 – 1 )
• Bond Order = 1/2
• Fractional Bond Orders CAN exist!
• Can expect to see a He₂⁺ molecule, but NOT a He₂.

​

Summarize

Can you:

• Explain bonding in terms of molecular orbital theory.
• Determine bond order and draw molecular orbital diagrams for diatomic molecules

​

READY

Molecular Orbitals: 2nd row Diatomic molecules

Objectives:

• Accurately create molecular orbital diagrams for second row diatomic molecules

​

2^nd Row Diatomic Molecules

• What's different about the 2^nd row vs 1^st?
• 2s and 2p orbitals
• What’s the same?
• The # of MO’s = the number of atomic orbitals combined
• Atomic orbitals combine best with other atomic orbitals of similar energy (s+s, p+p, NOT s+p)
• More overlap = more effective bonding (and lower MO energy)
• Each MO can hold 2 electrons with opposite spins (Pauli’s Exclusion Principle)
• When MOs are being filled, one electron goes in each MO with the same energy before pairing of electrons (Hund’s Rule)

Li₂

• Li : 1s²2s¹
• 1s orbitals make molecular orbitals in the same way H₂ did
• 2s orbitals mix in the same way BUT
• can overlap MORE than the 1s orbitals
• Larger/further away from the nucleus
• Greater drop for the energy of σ_2s than the σ_1s.

​

​

1s²

1s²

Li (1s²2s¹)

Li (1s²2s¹)

Li₂ Bond Order

• BO = ½(Bonding electrons – antibonding electrons)
• = ½ (4 – 2)
• = 1

NOTE:

• The core electrons (1s²) fill the bonding and antibonding MOs, effectively canceling each other
• Core electrons don’t contribute to the bonding of molecules
• We can ‘ignore’ the core electrons and focus on just the valence shell
• BO = ½ (2 - 0) = 1

​

Be₂

• 1s²2s²
• Lets ignore the core electrons (1s²)
• Atomic orbitals
• Molecular orbitals
• Fill in electrons
• Bond order = ½ (2 – 2) = 0
• Be₂ does not exist!

​

2s²

2s²

Be

Be

Molecular Orbitals From 2p Atomic Orbitals

σ^*_2p

σ_2p

Head to head

Molecular Orbitals From 2p Atomic Orbitals

π^*_2p

π_2p

Side to side

(over and under)

Molecular Orbitals From 2p Atomic Orbitals

π^*_2p

π_2p

Side to side

(Left and right)

Molecular Orbitals From 2p Atomic Orbitals

σ^*_2p

σ_2p

π^*_2p

π_2p

π^*_2p

π_2p

Head to head

Side to side

Side to side

Molecular Orbitals From 2p Atomic Orbitals

• B₂, C₂, N₂
• π_2p < σ_2p

​

​

σ^*_2p

σ_2p

π^*_2p

π_2p

Molecular Orbitals From 2p Atomic Orbitals

• O₂, F₂, Ne₂
• σ_2p < π_2p

σ^*_2p

σ_2p

π^*_2p

π_2p

Why cant anything in chemistry be easy?

To Know:

• When to use which structure
• The reason they change relative energies is because of 2p and 2s orbitals interacting
• Large interactions for B₂, C₂, and N₂. (Larger s orbitals)
• Small for O₂, F₂, and Ne₂ (smaller s orbitals)

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σ^*_2p

σ_2p

π^*_2p

π_2p

σ^*_2s

σ_2s

σ^*_2p

σ_2p

π^*_2p

π_2p

σ^*_2s

σ_2s

B₂ C₂ N₂ O₂ F₂ Ne₂

B₂

• [He] 2s²2p¹
• Place the electrons
• Bond Order = ½ (4-2) = 1

2s²

2s²

2p¹

2p¹

C₂

• [He] 2s²2p²
• Place the electrons
• Bond Order = ½ (6-2) = 2

2s²

2s²

2p²

2p²

N₂

• [He] 2s²2p³
• Place the electrons
• Bond Order = ½ (8-2) = 3

2s²

2s²

2p³

2p³

O₂

• *MO structure changes*
• [He] 2s²2p⁴
• Place the electrons
• Bond Order = ½ (8-4) = 2

2s²

2s²

2p⁴

2p⁴

F₂

• [He] 2s²2p⁵
• Place the electrons
• Bond Order = ½ (8-6) = 1

2p⁵

2p⁵

Ne₂

• [He] 2s²2p⁶
• Place the electrons
• Bond Order = ½ (8-8) = 0

2p⁶

2p⁶

Heteronuclear Diatomic Molecules

• We’ve been looking at Homonuclear Diatomic Molecules
• Same nuclei, 2 atoms
• F-F
• Now lets take a peak at heteronuclear diatomic molecules
• Different nuclei, 2 atoms
• N-O
• The 2s orbitals of different elements have different energies BUT
• If they are similar enough you can treat them in the same way you did homonuclear diatomics
• The 2p atomic orbitals also have different energies

2p³

2p⁴

N

[He] 2s²2p³

O

[He] 2s²2p⁴

Heteronuclear Diatomic Molecules

• NO molecule
• Important biologically
• Memory
• Relaxing muscles
• Killing foreign cells
• Lewis structures suggest double bond
• Observations (shorter bond length) suggest higher bond order
• MO theory better explains the actual structure

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N = O

N = O

-1 +1

Heteronuclear Diatomic Molecules

• MO theory better explains the actual structure
• Lewis structure:
• Double bond
• MO Theory:
• BO = ½ (8 – 3 ) = 2.5
• Bond order 2.5 agrees more with the observed bond lengths

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2p³

2p⁴

N

[He] 2s²2p³

O

[He] 2s²2p⁴

Summarize

Can You

• Accurately create molecular orbital diagrams for second row diatomic molecules

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Diamagnetism and Paramagnetism

Objectives:

• Determine if a molecule is diamagnetic or paramagnetic and describe the properties of each.

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Paramagnetism

• Paramagnetism: Molecules with one or more unpaired electrons will be attracted to a magnetic field
• The more unpaired electrons there are, the more attracted to the magnetic field it would be.

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O₂

Paramagnetism

• A Note:
• O₂ ‘s Lewis structure doesn’t show the unpaired electrons
• But the Molecular Orbitals do!

O₂

O=O

Diamagnetism

• Diamagnetism: substances with NO unpaired electrons are weakly REPELLED from a magnetic field
• MUCH weaker than paramagnetism

N₂

Paramagnetic (P) or Diamagnetic (D)?

B₂ C₂ N₂ O₂ F₂ Ne₂

Summarize

Can you:

• Determine if a molecule is diamagnetic or paramagnetic and describe the properties of each.

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