Chapter 14: Phenomena Phenomena: Scientists knew that in order to - - PowerPoint PPT Presentation

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Chapter 14: Phenomena

Phenomena: Scientists knew that in order to form a bond, orbitals on two atoms must overlap. However, px, py, and pz orbitals are located 90˚ from each other and compounds like CH4 (which would form bonds using their p

• rbitals) do not have bond angles of 90˚. Therefore, scientists had to explain this discrepancy or go back and

reevaluate quantum mechanics. Scientists realized that because electrons have wave properties they should mix with each other forming new differently shaped orbitals. To determine what these new orbitals looked like scientists used computers to combine the orbitals on different atoms generating as many combinations as

• possible. The three picture set on the left shows three different combinations of two unmixed orbitals (green).

The picture set in the middle shows all possible combinations of mixing the two orbitals. The mixed orbitals are shown in purple. The black dots in the pictures represent the nuclei of the two atoms. What patterns do you notice when the orbitals mix? The picture set on the right shows all of the orbital mixing for NO and HF. What do you notice about the orbital diagrams for these compounds.

Initial Orbitals Mixed Orbitals

p p s s

nitrogen N

• xygen

O nitrogen monoxide NO

p s

Chapter 14 Covalent Bonding: Orbitals

• Local Electron Model

(Valence-Band Theory)

• Molecular Orbital Theory

2

Big Idea: Bonding can be described using two theories which take into account quantum

• mechanics. In the

Local Electron Model, bonds are formed from the overlap of atomic orbitals. In Molecular Orbital Theory, electrons are redistributed throughout the molecule and placed into new orbitals called molecular

• rbitals.

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Local Electron Model (Valence-Bond Theory)

VSPPR (Lewis Model): Did not take into

account quantum mechanic’s effects. Assumes bonds located directly between atoms, therefore, electrons did not have wavelike properties

Local Electron Model (Valence-Bond

Theory): Uses a quantum mechanical description of the distribution of electrons in bonds that provides a way of calculating the numerical values of bond angles and bond lengths

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Local Electron Model (Valence-Bond Theory)

 Overlap: The merging of

• rbitals belonging to different

atoms of a molecule.

 σ-bond: Two electrons in a

cylindrically symmetrical cloud between two atoms.

 Nodal Plane: A plane on

which electrons will not be found.

4

σ-bonds

Not

• te: σ-bonds contain no

no nodal planes along the internuclear axis. Not

• te: The greater the extent of orbital
• verlap, the stronger the bond.

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Local Electron Model (Valence-Bond Theory)

 A σ–bond is formed in HF when electrons in 1𝑡-

and 2pz- orbitals pair (where z is the direction along the internuclear axis). Notice that there is cylindrical symmetry and no nodal plane on the internuclear axis.

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Local Electron Model (Valence-Bond Theory)

 𝝆-bond: A bond

formed by the side-to- side overlap of two p-

• rbitals

 A σ-bond is formed by

the pairing of electron spins in the two 2pz-

• rbitals

 𝜌-bonds are formed

when electrons in two

• ther 2p-orbitals pair

and overlap side by side.

6

Not

• te: 𝜌-bonds contain a single nodal

plane along the internuclear axis

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Student Question

Local Electron Model (Valence-Bond Theory)

How many 𝜏 bond and 𝜌 bonds are there in CO2? Hint: Draw the Lewis structure.

a) 1 𝜏 bond and 1 𝜌 bonds b) 0 𝜏 bond and 2 𝜌 bonds c) 2 𝜏 bond and 2 𝜌 bonds d) None of the Above

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Local Electron Model (Valence-Bond Theory)

Promotion of an electron is possible if:

 There are empty p-orbitals  The energy gained by forming additional bonds

is greater than the energy needed to promote the electron to the p orbital

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Promotion Can Occur For Carbon Promotion Cannot Occur For Nitrogen

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Local Electron Model (Valence-Bond Theory)

These are the

bonding

• rbitals of C,

therefore, what angles should be between each H in CH4?

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Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Local Electron Model (Valence-Bond Theory)

 These hybrid orbitals can

be mathematically represented by linear combinations of the atomic orbitals (within

• ne atom).

 h1 = ½(s + px + py + pz)  h2 = ½(s - px - py + pz)  h3 = ½(s - px + py - pz)  h4 = ½(s + px - py - pz)

10

Not

• te: Since one s orbital and three p orbitals went in to forming the new hybrid
• rbitals, these hybrid orbital are referred to as sp3 hybridized orbitals.

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Local Electron Model (Valence-Bond Theory)

The new molecular orbitals have energies

that are at the same level.

The hydride orbitals show that CH4 should

be in a tetrahedral bonding configuration.

11

Not

• te: The number of atomic orbitals that go into the linear combinations are the

same number of hybrid orbitals that form.

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Local Electron Model (Valence-Bond Theory)

Hybrid orbitals can be formed from other

combinations of atomic orbitals.

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 h1 = s + 2py  h2 = s +

3 2px - 1 2py

 h3 = s -

3 2 px - 1 2py

 h1 = s + p  h2 = s - p

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Local Electron Model (Valence-Bond Theory)

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Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Local Electron Model (Valence-Bond Theory)

Assigning Hybridization

 Step 1: Draw Lewis structure.  Step 2: Count the number of bonds and lone pairs

• n the atom of interest.

 Step 3: Assign hybridization sup to 1 pup to 3 dup to 5

Describe Bonding using the local electron (LE) model

 Step 1: Draw Lewis structure (if possible obey the

• ctet rule).

 Step 2: Determine hybridization.  Step 3: Describe bonding.

14

Not

• te: All types of bonds (single, double, and triple) between two atoms count

as 1 bond.

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Local Electron Model (Valence-Bond Theory)

 The bonding model that we looked at before for

N2 was a little oversimplified. The sigma bonding should be looked at as taking place between two sp hybridized orbitals instead of between two pz orbitals. However, sp hybridized orbitals are very similar in shape to pz orbitals

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Local Electron Model (Valence-Bond Theory)

LE Description of Bonding

 Sulfur forms one 𝜏 bond to each oxygen atoms. The bonds are

formed from the overlap of a sp3 hybridized orbitals on both the sulfur and oxygen atoms. All the loan pair electrons on both sulfur and oxygen atoms are located in sp3 hybridized orbitals.

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SO3

2-

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Local Electron Model (Valence-Bond Theory)

LE Description of Bonding

 The carbon atom forms one 𝜏 bond to each of the single bonded

• xygen atoms. These bonds are formed from the overlap of sp2

hybridized orbitals on the carbon atom and sp3 hybridized orbitals on the single bonded oxygen atoms. A third 𝜏 bond is formed from the

• verlap of an sp2 hybridized orbital on the carbon atom and a sp2

hybridized orbital on the double bonded oxygen atom. The π bond between the double bonded oxygen atom and the carbon atom is formed from the overlap of the unhybridized p orbitals on both the carbon and oxygen atoms. The loan pair electrons on the double bonded carbon sit in sp2 hybridized orbitals and the loan pair electron

• n the single bonded oxygen atoms sit in sp3 hybridized orbitals.

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CO3

2-

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Student Question

Local Electron Model (Valence-Bond Theory)

Identify the hybrid orbitals used by the underlined atom in acetone, CH3COCH3. The O atom is double bonded to the central carbon atom.

a) sp3d b) sp2 c) None; pure pz-orbitals are used in

bonding.

d) sp3 e) sp

If you have extra time tell the person next to you the LE description of the molecule.

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Local Electron Model (Valence-Bond Theory)

What atoms can form double and triple bonds?

 Atoms in period 2 (especially C, N, O) readily form

double bonds with themselves and other period 2 atoms.

 However, atoms in period 3 and later have trouble

forming multiple bonds with other large atoms due to the fact that the atoms are so large and bond lengths so great that it is difficult for their p-orbitals to take part in effective side-by-side bonding.

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C O O N N

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Local Electron Model (Valence-Bond Theory)

Limits of Lewis Theory/VSEPR and LE Model

1.

Cannot draw some structures that are known to exist.

Ex: B2H6 (12 valence e-) Not enough electrons to make all of the bonds

2.

Does not explain resonance structures

3.

Paramagnetic/Diamagnetic Problems

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Local Electron Model (Valence-Bond Theory)

Paramagnetic: Having the tendency to

be pulled into a magnetic field; a paramagnetic substance is composed of atoms or molecules with unpaired spin.

Diamagnetic: A substance that tends to

be pushed out of a magnetic field; a diamagnetic substance is composed of atoms or molecules with no unpaired electrons.

21

Not

• te: Very weak response and is not observable in every day life.

Not

• te: Laymen would call these materials magnetic.

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Local Electron Model (Valence-Bond Theory)

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Molecular Orbital Theory

Molecular orbitals are formed by

superimposing atomic orbitals of all the atoms in the molecule.

The molecular orbitals that are formed

are a linear combination of atomic

• rbitals (LCAO).

23

Not

• te: This is similar to Local Electron Model (valence band theory), however, the

Local Electron Model only formed hybrid orbitals from σ bonds and lone pair electrons within one atom. Molecular Orbital Theory is going to use all of the atomic orbitals on all of the atoms. Not just σ orbitals and lone pair electrons. Not

• te: Superimposing just means adding together.

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

 + and – are the signs of the wavefunction  No electron density in bonding plane  Electron density in bonding plane

Molecular Orbital Theory

24

Not

• te: The number of atomic orbitals that go into making the molecular orbitals is the

number of molecular orbitals generated.

σ bonds H2

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Molecular Orbital Theory

Bonding Orbitals: A molecular orbital with

no nodes between neighboring atoms.

Antibonding Orbitals: A molecular orbital

with a node between all neighboring pairs of atoms.

25

Not

• te: These orbitals contribute to holding all the atoms together.

Not

• te: These orbitals contribute to pushing all the atoms apart.

Not

• te: These type of orbitals are denoted with a * next to them.

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Student Question Molecular Orbital Theory

Which molecular orbital is a σ?

26

Which molecular orbital is a σ*? Which molecular orbital is a π? Which molecular orbital is a π*? a) b) c) d)

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Molecular Orbital Theory

How do Electrons Fill Molecular Orbitals?

 Electrons are first accommodated in the

lowest-energy molecular orbitals, followed by

• rbitals of increasingly higher energy.

 According to the Pauli Exclusion Principle,

each molecular orbital can accommodate up to two electrons. If two electrons are present in one orbital they must be paired.

 If more than one molecular orbital of the same

energy is available the electrons enter them singly and adopt parallel spins (Hund’s Rule).

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Molecular Orbital Theory

MO diagram for homonuclear diatomic molecules Li2 through N2 MO diagram for homonuclear diatomic molecules O2 and F2

28

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Molecular Orbital Theory

Comparisons Between the VSEPR/LE/MO Theory

 VSEPR (N2): N2 has a triple bond  LE (N2): N2 has 1 σ bond that is formed from the

• verlap of sp hybridized orbitals on the nitrogen

atoms and 2 π bonds that are formed from the

• verlap of p orbitals on the nitrogen atoms

 MO (N2)

 Bond Order (b): The number of electron pairs (bonds)

that link a specific pair of atoms

 All three predict that N2 bonds with a triple bond

29

N≡N



b = ½ (N-N*)



N = # of e- in bonding orbitals



N* = # of e- in antibonding orbitals)

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Student Question

Molecular Orbital Theory

What is the bond order of O2?

a) 2.5 b) 2 c) 1.5 d) 1 e) None of the above

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Molecular Orbital Theory

Heteronuclear Diatomic Molecules

31

Not

• te: The energy level of atomic orbitals decreases as electronegativity increases.

Less Electronegative More Electronegative

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Molecular Orbital Theory

32

Not

• te: The molecular orbitals for CO fill in the same order as C2.

σ2s σ*2s σ2p σ*2p π*2p π2p

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Molecular Orbital Theory

 Similar to diatomic molecules, the number of

molecular orbitals in a polyatomic species equals the number of atomic orbitals that are available.

33 2a1 2t1 1t1 1a1 Energy Antibonding Bonding

Carbon 4 atomic orbitals (2s,2px,2py,2pz) Hydrogens 4 [1 atomic orbit each (1s)] 8 molecular orbitals total MO Diagram CH4

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Molecular Orbital Theory

 We already know that when N atomic

• rbitals merge together in a molecule,

they form N molecular orbitals. The same is true of a metal; but for a metal N is enormous (about 1023 for 10 g of copper). This results in the energy levels being so close together that they form a continuous band.

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 Highest Occupied Molecular Orbital

(HOMO): The highest-energy molecular orbital in the ground state

• f a molecule occupied by at least
• ne electron.

 Lowest Unoccupied Molecular Orbital

(LUMO): The lowest-energy molecular

• rbital that is unoccupied in the

ground state.

LUMO HOMO

2a1 2t1 1t1 1a1 Energy Antibonding Bonding

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Molecular Orbital Theory

35

Metal Insulator Semiconductor Conduction Band Valence Band Energy

 Valence Band: In the theory of solids, a band of

energy-levels fully occupied by electrons.

 Conduction Band: An incompletely occupied band

• f energy-levels in a solid.

Not

• te: In order to conduct electricity electrons must be promoted from the valence

band into the conduction band.

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Molecular Orbital Theory

Semiconductors can be doped to enhance electrical properties

 Doping: The addition of a known, small

amount of a second substance, to an

• therwise pure solid substance.

 n-type Semiconductor: Dopants are added to

the material that provide extra electrons.

 p-type Semiconductor: Dopants are added to

the material that provide extra holes (less electrons)

36

Exa xampl ple: : As added to Si Exa xample: ple: In added to Si

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Molecular Orbital Theory

MOSFET Devices

 When you apply a positive bias to the gate you end up

generating an e- channel at the oxide semiconductor interface which allows electrons to flow between the source and the drain.

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Oxide Gate

+ Bias (VD) Bias (VG = 0)

Semiconductor

Source Drain

h e- e- e- e- e- e- h h e- e- e- e- h h h h h h h

Oxide Gate

+ Bias (VD)

Semiconductor

Source Drain

h e- e- e- e- e- h h e- e- e- e- h h h h h h h

Bias (VG = +)

e- e- e- e- e-

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Take Away From Chapter 14

 Big Idea: Bonding can be described using two theories which

take into account quantum mechanics. In the Local Electron Model, bonds are formed from the overlap of atomic

• rbitals. In Molecular Orbital Theory, electrons are

redistributed throughout the molecules and placed into new

• rbitals called molecular orbitals.

 Local Electron Model (Valence Band-Theory)

 Know that bonding occurs from the overlap of atomic

• rbitals on neighboring atoms (12)

 σ Bonds: Cylindrical symmetry no nodal plane on internuclear

axis

 π Bonds: One nodal plane on internuclear axis

 Be able to determine the number and type (σ or π) of

bonds in a molecule (25, 26,27)

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Numbers correspond to end of chapter questions.

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Take Away From Chapter 14

 Local Electron Model (Valence Band-Theory)

(Continued)

 Know that electron promotion occurs to accommodate

more bonds when there are vacancies in the p-orbitals

 Know that in order to describe the experimentally seen

shape of molecules, a linear combination of atomic

• rbitals is needed (hybridization)

 Be able to determine the hybridization of atoms (21,22,28,79)

 Be able to write out the local electron (LE) description of

bonding (14)

 Be able to determine which atoms are in the same

bonding plane. (23,24)

 Know the limitations of the Local Electron Model/VSEPR

(Lewis Model)

 Not all molecules can be explained (B2H6)  Does not account for resonance  Incorrect magnetic properties (O2)

39

Numbers correspond to end of chapter questions.

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Take Away From Chapter 14

Molecular Orbital Theory

 Know that atomic orbitals, from all atoms in the molecule,

are mixed together to form molecular orbitals.

 The new molecular orbitals are delocalized over the entire

molecule (accounts for resonance structures). (51)

 Be able to visually identify bonding/antibonding orbitals

and σ and π molecular orbitals. (32,33)

 Memorize molecular orbital energy levels for homonuclear

diatomic atoms in periods one and two.

 Be able to write out molecular electron configuration

 Be able to determine the bonding order (40,41,44,53)

 𝐶𝑃 = #𝑐𝑝𝑜𝑒𝑗𝑜𝑕 𝑓− − #𝑏𝑜𝑢𝑗𝑐𝑝𝑜𝑒𝑗𝑜𝑕 𝑓− 2  If bonding order is positive - stable species.  If bonding order is 0 - not a stable species.

40

Numbers correspond to end of chapter questions.

Chapt pter er 14: Coval valent ent Bondin ding: g: Orbita tals ls

Take Away From Chapter 14

Molecular Orbital Theory (Continued)

 The greater the number of atoms in the molecule the

greater the number of molecular orbitals

 Be able to determine if a structure paramagnetic

(unpaired e-) or diamagnetic (paired e-)

41

Numbers correspond to end of chapter questions.