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Assorted General Chemistry Problems

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8.20) a. The lattice energies of NaF and MgO are 910kJ/mol and 3795kJ/mol. Account for the difference in these two quantities. b. Account for the difference in lattice energies of MgCl2 (2326 kJ/mol) and SrCl2 (2127kJ/mol).

8.28) Calculate the lattice energy of CaCl2. Is this value greater than or less than the lattice energy of NaCl (788kJ/mol)? Explain.

8.50) For each of the following molecules or ions of sulfur and oxygen, write a single Lewis structure that obeys the octet rule, and calculate the oxidation numbers and formal charges on all the atoms: SO2, SO3, SO3 2-, SO4 2-.

8.52) Consider the nitryl cation, NO2+. a. Write one or more appropriate Lewis structures for the species. b. Ar3e resonance structures needed to describe the structure? c. With what familiar species is it isoelectronic?

8.62) Draw the Lewis structures for each of the following molecules or ions. Which do not obey the octet rule?
a. CO2 b. IO3- c. BH3 d. BF4- e. XeF2.

8.66) Using bond enthalpies, estimate ΔH for the following gas-phase reactions:
a. Br Br
Br--C-- H+Cl-- Cl  Br--C--Cl+H--Cl
Br Br

b. H H H H
H-- S-- C -- C-- S-- H+ 2H -- Br  Br -- C -- C -- Br+ 2H -- S-- H
H H H H

c. H H H
H -- N -- N -- H+Cl -- Cl  2H -- N -- Cl

8.98) One scale for electronegativity is based on the concept that the electronegativity of any atom is proportional to the ionization energy of the atom minus its electron affinity; electronegativity = k (IE-EA), where k is a proportionality constant. a. How does the definition explain why the electronegativity of F is greater than that of Cl even though Cl has the greater electron affinity? b. Why are both ionization energy and electron affinity relevant to the notion of electronegativity?

8.106) Average bond enthalpies are generally defined for gas phase molecules. Many substances are liquids in their standard state. Calculate average bond enthalpies in the liquid state for the following bonds, and compare these values to the gas-phase values.
a. Br--Br, from Br2 (l); b. C--Cl, from CCl4 (1); c. O--O, from H2O2 (1) (assume that the O--H bond enthalpy is the same as in the gas phase). d. What can you conclude about the process of breaking bonds in the liquid compared to the gas phase? Explain the difference in the ΔH values between the two phases.

9.22) Give the electron-domain and molecular geometries for the following molecules and ions: a. HCN, b. SO3 2-, c. SF4, d. PH6-, e. NH3Cl+, f. N3-

9.36) Predict whether each of the following molecules is polar or nonpolar: a. IF, b. CS2, c. SO3, d. PCl3, e. SF6, f. IF5

9.48) What set of hybrid orbitals is used by the central atom in a. SiCl4, b. HCN, c. SO3, d. ICl2-, e. BrF4-?

9.54) Ethyl acetate, C4H8O2, is a fragrant substance used both as a solvent and as an aroma enhancer. a. What is the hybridization at each of the carbon atoms of the molecule? b. What is the total number of valence electrons in ethyl acetate? c. How many of the valence electrons are used to make or Ơ bonds in the molecule? d. How many valence electrons are used to make π bonds? e. How many valence electrons remain in non-bonding pairs in the molecule?

9.66) Explain the following: a. The peroxide ion, O2 2-, has a longer bond than the superoxide ion, O2-. b. The magnetic properties of B2 are consistent with the π2p MOs being lower in energy than the Ơ2p MO.

9.68) a. What is meant by the term paramagnetism? b. How can one determine experimentally whether a substance is paramagnetic? c. Which of the following ions would you expect to be paramagnetic: O2+, N2 2-, Li2+, O2 2-? If the ion is paramagnetic, how many unpaired electrons does it possess?

9.90) The cyclopentadienide ion has a formula C5H5-. The ion consists of a regular pentagon of C atoms, each bonded to two C neighbors, with a hydrogen atom bonded to each C atom. All the atoms lie in the same plane. a. Draw a Lewis structure for the ion. According to your structure, do all five C atoms have the same hybridization? Explain. b. Chemists generally view this ion as having sp2 hybridization at each C atom. Is that view consistent with your answer to part a? c. Your Lewis structure should show one nonbonding pair reside? d. Are there resonance structures equilivant to the Lewis structure you drew in part A? If so, how many? e. The ion is often drawn as a pentagon enclosing a circle. Is this representation consistent with your answer to part d? Explain. f. Both benzene and the cyclopentadienide ion are often described as systems containing six π electrons. What do you think is meant by this description?

9.97) Sulfur tetrafluoride (SF4) reacts slowly with O2 to form sulfur tetrafluoride monoxide (OSF4) according to the following unbalanced reaction:
SF4 (g) + O2 (g)  OSF4 (g)
The O atom and the four F atoms in OSF4 are bonded to a central S atom. a. Balance the equation. b. Write a Lewis structure of OSF4 in which the formal charges of all atoms are zero. c. Use average bond enthalpies to estimate the enthalpy of the reaction. Is it endothermic or exothermic? d. Determine the electron-domain geometry of OSF4, and write two possible molecular geometries for the molecule based on this electron-domain geometry. e. Which of the molecular geometries in part d is more likely to be observed for the molecule? Explain.

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https://brainmass.com/chemistry/general-chemistry/assorted-general-chemistry-problems-161923

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Assorted General Chemistry Problems are solved with full working shown. Molecular geometry is strongly emphasized.

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Assorted General Chemistry Problems

1) List the following gases in order of increasing root-mean-square speed, all at the same temperature: H2, He, O2, Ne

2) Under what conditions of pressure and temperature is a real gas most expected to deviate from ideal gas behaviour?

3) What intermolecular interactions would be the strongest between molecules CH2Cl2?

4) List the following substances in order of increasing boiling point: H2O, H2S, H2Se, H2Te.

5) List the following compounds in order of increasing vapour pressure at 298K: pentane (C5H10), ethanol (C2H5OH), water (H2O)

6) Give one example of a molecular solid.

7) The Henry's Law constant for CO2 is 3.1X10^-2 molL^-1 atm^-1 at 298K. What is the concentration of dissolved CO2 in water exposed to an atmosphere of pure CO2 at 2.0atm? Report your answer to two significant figures.

8) What is the molarity of a solution prepared by dissolving 6.21g of glucose (C6H12O6) in 1.50L of water. Report your answer to the correct number of significant figures.

9) List the following solutions in order of increasing freezing points: 0.1 m NaCl(aq), 0.1 m MgCl2(aq), 0.1 m glucose.

10) Define internal energy and enthalpy. What is the difference between the two?

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