PS4

3.091 Intro. to Solid-‐State Chem. PS 4

3.091 Introduction to Solid-‐State Chemistry -‐ Fall 2011

Problem Set 4 No problem set quiz scheduled for PS4

Material will be covered on Test 1, October 1

1. Silicon (Si) reacts with chlorine (Cl), forming the compound SiCl4 (which is a homologue of CCl4, a liquid used as a spot cleaner). (a) Using Lewis notation, sketch this compound assuming that it has achieved octet

stability. (b) List the atomic orbitals which, on overlap, result in the formation of this

compound as well as the final molecular orbitals. (c) What do you expect to be the state of aggregation of this compound at room

temperature (a liquid, solid or gas), and what do you assume is the nature of the intermolecular forces acting (if any are active)?

(d) Knowing that the bond energies for Si-‐Si and Cl-‐Cl are respectively, 176 and 240 kJ/mol, determine the bond energy for the Si-‐Cl bond (in kJ/mol) according to Pauling.

2. Determine the minimum frequency (v) of radiation capable of breaking C-‐Cl bonds, given BECl-‐Cl = 243 kJ/mole and BEC-‐C = 245.4 kJ/mole. (a) Draw a 3-‐dimensional representation of the molecular geometry (not simply the

Lewis structure) of hydrazine, H2N-‐NH2. (b) Hydrazine and hydrogen chloride have nearly identical molecular weights, yet

hydrazine boils at 113.5°C and is a liquid at room temperature, while hydrogen chloride boils at -‐85°C and is a gas at room temperature. Explain with reference to electronic structure and bonding.

(c) Can visible light (400 <λ< 700 nm) break any bonds in hydrazine? Support your answer with calculations. Bond dissociation energies (kJ/mol): N-‐N = 160: H-‐H = 435.

3. (a) Draw a 3-‐dimensional representation of the molecular geometry around the central atom (not simply the Lewis structure) of AsClF4

2-‐. (b) Name the type of hybrid orbitals that the central atom forms. (c) Name the molecular geometry of the compound. (d) Estimate the Cl-‐As-‐F bond angle. Justify.


(e) Calculate the maximum wavelength, λ, of electromagnetic radiation capable of breaking the weakest bond in of AsClF4

2-‐. Bond Energies (kJ/mol): As-‐As 180; F-‐F 160; Cl-‐Cl 240.

4. (a) The ionization energies of aluminum. (gas-‐phase) are as follows: 1s 2s 2p 3s 3p 151 12.1 7.79 1.09 0.58 (MJ/mol) Sketch the photoelectron spectrum of Al (intensity vs energy). The diagram need not be drawn to scale; however, you must pay attention to relative magnitudes. (b) Calculate the average valence electron energy of aluminum.

5. (a) Construct an energy-‐level diagram using LCAO-‐MO theory (linear combination of atomic orbitals into molecular orbitals) to show that the nitric oxide ion, NO+, is stable. The filling sequence of the molecular 2p orbitals in NO+ is π2p, σ2p, π*2p, σ*2p. Indicate relative positions of the atomic and molecular energy levels as well as electron occupancy of the valence shell of each species. (b) The cyanide ion. CN-‐ is isoelectronic with NO+. Explain with reference to the relevant underlying physics why the C-‐N bond length in CN-‐ is greater than the N-‐O bond length in NO+. (c) Draw a 3-‐dimensional representation of athe molecular geometry around the central atom (not simply the Lewis structure) of arsine, AsH3. (d) Estimate the As-‐H bond angle in AsH3. (e) Is AsH3 polar or nonpolar? Justify.


6. (a) Draw a 3-‐dimensional representation of the molecular geometry around the central

atom (not simply the Lewis structure) of the bromoiodate anion,I2Br-‐. (b) Name the type of hybrid orbitals that the central atom forms. (c) Name the molecular geometry of the compound. (d) Is the molecule polar or nonpolar? Explain. (e) Calculate the bond energy of I-‐Br. Express your answer in kJ/mol. DATA: bond dissociation energies (kJ/mol): I-‐I = 150: Br-‐Br = 195.

7. Given the compound C2Cl4: (a) Draw a sketch for this compound with the electronic orbitals which on overlap form the intramolecular bonds (in this molecule) and label the orbitals involved. (b) List the total number of molecular orbitals formed in the compound and give the corresponding atomic orbitals involved in their formation.

8. Hydrogen and selenium react with each other (like hydrogen and oxygen). Do you expect the compound H2Se formed to be a solid, liquid or gas? (Rationalize your answer.) (a) Draw the Lewis structure of nitrosyl chloride (NOCl). (b) Is NOCl polar or nonpolar? Explain. (c) The value of the N-‐N bond energy is not easy to measure. Given the values of 190 kJ/mol for the N-‐Cl bond energy and 240 kJ/mol for the Cl-‐Cl bond energy, calculate the value of the N-‐N bond energy.


9. (a) Write the Lewis structure of BH4

-‐. (b) Is BH4

-‐ polar or nonpolar? Explain.

10. DATA: Compound boiling point compound boiling point compound boiling point (°C) (°C) (°C) CF4 -‐128 CCl4 76 CBr4 193 CH2F2 -‐52 CH2Cl2 40 CH2Br4 97 CH4 -‐162 CH4 -‐162 CH4 -‐162 (a) Explain why methane (CH4) boils at a lower temperature than carbon tetrafluoride (CF4). (b) Explain why difluororomethane (CH2F2) boils at a higher temperature than carbon tetrafluoride (CF4). (c) Explain why difluororomethane (CH2F2) boils at a higher temperature than carbon tetrafluoride (CF4) yet dichloromethane (CH2Cl2) boils at a lower temperature than carbon tetrachloride (CCl4), and dibromomethane (CH2Br2) boils at a lower temperature than carbon tetrabromide (CBr4).

11. Use molecular orbital theory to explain why the oxygen-‐oxygen bond is stronger in the O2 molecule than in the O2

2-‐ (peroxide) ion.

12. Use molecular orbital theory to predict whether the bond order in the superoxide ion, O2-‐,

should be higher or lower than the bond order in a neutral O2 molecule.

13. Use molecular orbital theory to predict whether the peroxide ion, O22-‐, should be

paramagnetic.

14. Write the electron configuration for the following diatomic molecules. Calculate the bond order in each molecule. (a) HF (b) CO (c) CN-‐ (d) ClO-‐ (e) NO+


15. Classify the following molecules as paramagnetic or diamagnetic. (a) HF (b) CO (c) CN-‐ (d) NO (e) NO+

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