Deck 7: Theories of Chemical Bonding

Question 1

Use the orbital overlap model to explain the bonding in simple molecules.

Answer 1

Bonding orbitals are constructed by combining atomic orbitals from adjacent atoms. Only the valence orbitals are needed to describe bonding.

Question 2

Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

Answer 2

The number of hybrid orbitals generated by the hybridization process equals the number of valence atomic orbitals participating in hybridization. The steric number of an inner atom can be used to infer the hybrid orbitals it is using. Elements beyond the second row of the periodic table can form bonds to more than four ligands and can be associated with more than an octet of electrons.

Question 3

Describe the ? and $\pi$ bonding systems in multiple bonds.

Answer 3

A sigma (?) bond has high electron density distributed symmetrically along the bond axis. A pi ( $\pi$ ) bond has high electron density concentrated above and below the bond axis.

Question 4

Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Question 5

Describe the bonding in three-atom $\pi$ systems.

Question 6

Describe the bonding in extended $\pi$ systems.

Question 7

Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

Question 8

Which of the following molecules is NOT well described by the valence bond model?

A) SiH₄
B) H₂S
C) SbH₃
D) H₂Te
E) HF

Question 9

According to the valence bond model, which of the following is NOT a requirement of bonding?

A) overlap of atomic orbitals
B) singly occupied atomic orbitals
C) no two electrons in a molecule have identical descriptions
D) atoms in molecule contain unpaired valence electrons
E) orbitals must have same azimuthal quantum number to overlap

Question 10

Valence bond theory based unhybrized orbitals fails to adequately describe the structure of methane

A) as carbon has only two unpaired electrons.
B) because p orbitals on carbon cannot overlap with s type orbitals on H to form bonds.
C) because it fails to adequately describe the shape of the methane molecule.
D) because the Aufbau principle is violated.
E) because methane is a square planar molecule.

Question 11

The concept of hybridization of atomic orbitals was introduced

A) to predict shapes of molecules that cannot be described using traditional atomic orbitals.
B) to explain shapes of molecules that cannot be described using traditional atomic orbitals.
C) to account for the fact that carbon can form 4 bonds.
D) to explain $\pi$ bonding.
E) to explain delocalized $\pi$ bonding.

Question 12

What is the hybridization of the central atom in BeCl₂?

A) sp
B) sp²
C) sp³
D) sp⁴
E) sp³d

Question 13

What is the hybridization of the central atom in ICl₂^-1?

A) sp
B) sp²
C) sp³
D) sp⁴
E) sp³d

Question 14

What is the hybridization of the central atom in SO₂?

A) sp
B) sp²
C) sp³
D) sp⁴
E) sp³d

Question 15

Identify the steric number, shape, hybridization about the central atom and molecular polarity of sulphuryl chloride, SO₂Cl₂.

A) 4, square planar, sp³, non-polar
B) 4, tetrahedral, sp³, polar
C) 6, tetrahedral, sp³d², polar
D) 4, square based pyramid, sp³, polar
E) 6, octahedral, sp³d², non-polar

Question 16

Identify the steric number, shape, hybridization about the central atom in XeF₂Cl₂.

A) 4, tetrahedral, sp³
B) 4, square planar, sp³
C) 6, octahedral, sp³d²
D) 6, tetrahedral, sp³d²
E) 6, square planar, sp³d²

Question 17

What is the hybridization of Xe in the molecule XeF₄?

A) sp³
B) sp²
C) sp³d²
D) sp³d
E) sp⁴

Question 18

The best description of the CO bond in acetaldehyde, CH₃-(C=O)-H is

A) a sigma bond and two pi bonds.
B) two sigma bonds.
C) two pi bonds.
D) a sigma bond and a pi bond.
E) a delta bond and a pi bond.

Question 19

Carbon dioxide, CO₂, exists as discrete molecular compounds, whereas SiO₂ exists as a network covalent compound. Identify the hybridization and the number of sigma and pi bonds on each C and Si.

A) C(sp³_, 2,2) S(sp³,2,2)
B) C(sp_,2,2) S(sp³,0,4)
C) C(sp_, 2,2) S(sp³,4,0)
D) C(sp_, 2,2) S(sp,2,2)
E) C(sp³_,4,0) S(sp³,4,0)

Question 20

Given that the bond energy of C-O single bond is 360 kJ/mol, a C-O double bond is 750 to 800 kJ/mol and Si-O single bond is 450 kJ/mol, what is the bond energy of a Si-O double bond?

A) 900 kJ/mol
B) significantly greater than 900 kJ/mol
C) significantly less than 900 kJ/mol
D) not possible to predict
E) slightly greater than 900 kJ/mol

Question 21

Below is the line structure for the amino acid tryptophan. Determine the type of hybridization and number of each type for the starred atoms.

A) 1 sp, 2 sp², 2 sp³
B) 1 sp, 3 sp², 1 sp³
C) 2 sp², 3 sp³
D) 3 sp², 2 sp³
E) 1 sp², 4 sp³

Question 22

How many sigma and pi bonds are in acetic acid, CH₃CO₂H?

A) 4 $\sigma$ , 3 $\pi$
B) 5 $\sigma$ , 2 $\pi$
C) 6 $\sigma$ , 2 $\pi$
D) 7 $\sigma$ , 1 $\pi$
E) 7 $\sigma$ , 2 $\pi$

Question 23

Which of the following cations (O₂⁺, C₂⁺, N₂⁺; F₂⁺) would have a higher bond order than the neutral molecule?

A) all of them, removing electrons always increases bond order
B) none of them, removing electrons always decreases bond order
C) O₂⁺, C₂⁺
D) N₂⁺; F₂⁺
E) O₂⁺, F₂⁺

Question 24

What is the bond order in Mg₂?

A) 0
B) ½
C) 1
D) 1.5
E) 2

Question 25

What property of diatomic oxygen, O₂, is explained by MO theory but not by Lewis, valence bond on VSEPR?

A) bond strength
B) paramagnetism
C) molecular shape
D) bond length
E) ability to form multiple bonds

Question 26

Consider HF, a heteronuclear diatomic molecule. Which of the following atomic orbitals will combine to form a $\sigma$ molecular orbital?

A) H(1s) F(1s)
B) H(2p_x) F(2p_x)
C) H(1s) F(2s)
D) H(1s) F(2p_x)
E) H(1s) F(2p_x, 2p_y, 2p_z)

Question 27

Which of the following molecules/ions (O₃, NO₂^-, NCO^-, N₂O, NO₂) have the same bonding scheme as CO₂?

A) NO₂^-, NO₂
B) O₃, N₂O, NO₂
C) NCO^-, N₂O, NO₂
D) NCO^-, N₂O
E) O₃, NO₂^-, NCO^-, N₂O, NO₂

Question 28

Which of the following best shows one of the $\pi$ bonds in H₂C=C=CH₂?

A)
B)
C)
D)
E)

Question 29

Which of the following does NOT have delocalized $\pi$ bonding?

A) NO₂
B) SO₂
C) F₂O
D) NO₂⁺
E) O₃

Question 30

In the allene (H₂C=C=CH₂) molecule,

A) all the atoms lie in one plane.
B) the $\pi$ bonds are delocalized over all three C atoms.
C) the two H atoms on one C lie above and below the plane defined by the other CH₂ group.
D) the non-bonding orbitals are full.
E) delocalized bonds are formed from overlap of hybridized orbitals.

Question 31

What is needed to allow a $\pi$ orbital to be delocalized?

A) 4 or more atoms
B) more than 2 "p" orbitals with appropriate geometry
C) 6 $\pi$ electrons
D) an s and a p orbital in a molecular bond
E) carbon-carbon bonds

Question 32

P-O double bonds are stronger than P-S double bonds because

A) the O atom had greater electronegativity.
B) there is less polar character to the P-S bond.
C) the O atom has a smaller radius than the S atom.
D) the P atom is less electronegative than the O atom.
E) S has accessible d orbitals.

Question 33

Which of the following sulphur species has the greatest delocalization as judged by the number of resonance structures?

A) H₃SO₄⁺
B) H₂SO₄
C) HSO₄^-
D) SO₄^2-
E) SO₃

Question 34

Consider the perchlorate anion, ClO₄^-:

A) the steric number for this anion is 4, the best description includes resonance structures having two double bonds and two single bonds, electrons are delocalized throughout the ion.
B) the steric number for this anion is 4, the best description includes resonance structures having three double bonds and one single bonds, electrons are delocalized throughout the ion.
C) chlorine is not capable of forming multiple bonds.
D) the steric number for this anion is 4, the best description includes resonance structures having three double bonds and one single bonds, electrons are not delocalized since the structure is not conjugated.
E) the steric number for this anion is 7, the best description includes resonance structures having three double bonds and one single bonds, electrons are delocalized throughout the ion.

Question 35

Which of the following compounds would be expected to absorb the longest wavelength light?

A)
B)
C)
D)
E)

Question 36

Which of the following is true of a conjugate $\pi$ system?

A) The molecule is more reactive.
B) The molecule is more volatile.
C) The molecule is less reactive.
D) The molecule has the same reactivity.
E) The molecule reacts violently.

Question 37

Which of the following compounds will absorb the most visible light, i.e,. be the most deeply coloured?

A)

B)

C)

D)

E)

Question 38

How many delocalized $\pi$ electrons are in the following molecule?

A) 6
B) 8
C) 10
D) 12
E) 14

Question 39

Which of the following elements would be added to germanium to produce an n-type semiconductor?

A) gallium
B) silicon
C) aluminum
D) arsenic
E) tin

Question 40

In n-type seminconductor the

A) donor level of the dopant lies close in energy to the conduction band.
B) donor level of the dopant lies close in energy to the valence band.
C) acceptor level of the dopant lies close in energy to the conduction band.
D) the acceptor level of the dopant lies close in energy to the valence band.

Question 41

What will happen to a semiconductor made of GaP, if some of the P is replaced with As?

A) The band gap remains the same.
B) The band gap grows larger.
C) The band gap becomes smaller.
D) This cannot happen because As is not isoelectronic with P.
E) It will emit light of shorter wavelength.

Question 42

What would you dope a GaP semiconductor with to decrease the wavelength of light emitted?

A) In
B) As
C) N
D) Ge
E) The band gap cannot be changed.

Question 43

Draw a picture showing the bonding between a hydrogen atom and a chlorine atom telling what orbitals from each are making the bond.

Question 44

Describe the bonding between iodine atoms in molecular iodine, I₂. Make sure to include a drawing to symbolize the overlap of orbitals.

Question 45

NH₃ has bond angles of 107.3˚. Describe the bonding in NH₃ only using unhybridized orbitals. Discuss why this type of bonding model is faulty.

Question 46

H₂O has a bond angle of 104.5˚. Describe the bonding in H₂O only using unhybridized orbitals. Discuss why this type of bonding model is faulty.

Question 47

Methylene chloride, CH₂Cl₂, is a common industrial solvent. Sketch an orbital overlap picture of the bonds and describe the bonding present.

Question 48

Determine the hybridization of the central atoms of H₂CCHCH₃.

Question 49

Determine the hybridization of the central atoms in H₂NCHCH₂.

Question 50

Draw the Lewis structure of the sulphite ion, SO₃^-2. Determine what the hybridization of the central atom is.

Question 51

Draw two sketches describing the bonding in propene, C₃H₆. The first should show any sigma bonding framework and the second sketch should illustrate the $\pi$ bonds that are formed, if any. Write a brief description of the bonding, naming the orbitals involved.

Question 52

For the molecule ethene, C₂H₄, sketch a picture that shows the location of the $\pi$ bond as oriented with the plane of the hydrogen atoms.

Question 53

Explain why carbon dioxide includes both pi and sigma bonds whereas SiO₂ is formed from a network of sigma bonds.

Question 54

Draw a picture illustrating the atomic orbitals participating in the bonding in formaldehyde, H₂C=O. List the orbitals used in bond formation.

Question 55

The Lewis structure of vinyl chloride is shown below; Draw orbital pictures of (a) the sigma bonding framework and (b) the $\pi$ bond.

Question 56

What orbitals overlap to make the bond(s) between C and O in acetone, CH₃COCH₃?

Question 57

In which chemical species would you expect the strongest bond: C₂⁺; N₂⁺; O₂⁺? Write the electronic molecular orbital configuration and calculate the bond order for that species.

Question 58

Write the electron configuration and predict the bond order and number of unpaired electrons O₂^-.

Question 59

Write the electron configuration and predict the bond order and number of unpaired electrons B-C.

Question 60

Determine the bond order for F₂. How many electrons are in antibonding orbitals?

Question 61

Determine the bond order for O₂⁺ and tell how many electrons are in antibonding orbitals.

Question 62

Draw orbital pictures of the highest energy molecular orbital occupied in fluorine F₂ and the atomic orbitals that form it.

Question 63

Predict the bond order and number of unpaired electrons in NO⁺.

Question 64

Predict using MO theory if NO^- will be paramagnetic or diamagnetic.

Question 65

What kind of bonds are obtained by the overlap of the d_xy and the p orbitals shown below?

Question 66

In the text's problems, you were asked to find the Lewis structure of HNCO. Compare that Lewis structure with that of the isomer, NCOH and suggest why the HNCO isomer may be more stable.

Question 67

Draw the bonding and antibonding molecular orbitals for NO₂^-.

Question 68

Explain the difference between an isolated $\pi$ orbital and a delocalized $\pi$ orbital.

Question 69

What are non-bonding molecular orbitals of ozone, O₃, and where do they originate?

Question 70

Determine the bonding for the nitrite ion, NO₂^-1. Determine what hybridization the central atom has and how many delocalized $\pi$ orbitals there are.

Question 71

Draw a combination of a 3d orbital on phosphorus and a 2p orbital on O that allows $\pi$ bonding.

Question 72

Describe the bonding in the polyatomic anion, nitrate (NO₃^-1) by drawing the resonance structures and the bonding $\pi$ molecular orbital.

Question 73

Draw the band diagram that is appropriate for a ZnS semiconductor that has been doped with copper.

Question 74

Explain how band theory can explain the electrical conductivity observed in Mg metal given that atomic Mg has valence electrons fully occupying the 2p atomic orbitals.

Question 75

A new material is colourless, but is not an insulator. What can you deduce about the energy of the bandgap? Give an estimate of the bandgap in kJ.

Question 76

Draw a representation of a band gap for graphite and germanium and discuss any differences between them.

Question 77

Semiconductors are used for solar energy conversion in devices called photovoltaic cells. For high efficiency, the cell should be able to utilize as much light as possible. What band gap (in kJ/mole) would be needed to utilize light of 550 nm (about the middle of the visible spectrum) and shorter wavelengths?

Question 78

Describe what happens to an electron as it travels through an n $\rightarrow$ p-type junction of a semiconductor.