Download Inorganic Chemistry A Self-study exercises Chapters 1,2 and 3 1

Self-study exercises |1
Inorganic Chemistry A
Self-study exercises
Chapters 1,2 and 3
1. Calculate the value of Ar for naturally occurring chlorine if the distribution of
isotopes is 75.77% 1735 Cl and 24.23% 37
Cl . Accurate masses for 35 Cl and 37 Cl
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are 34.97 and 36.97.
2. If Ar for Cl is 35.45, what is the ratio of
atoms containing naturally occurring Cl?
35
Cl : 37 Cl present in a sample of Cl
3. Calculate the value of Ar for naturally occurring Cu if the distribution of
isotopes is 69.2% 63 Cu and 30.8% 65 Cu ; accurate masses are 62.93 and 64.93.
[Ans. 63.5]
4. Why in question 3 is it adequate to write 63 Cu rather than 2963 Cu ?
5. Calculate Ar for naturally occurring Mg if the isotope distribution is 78.99%
24
Mg , 10.00% 25 Mg and 11.01% 26 Mg ; accurate masses are 23.99, 24.99 and
25.98.
6. ‘Arsenic is monotopic.’ What does this statement mean?
7. Given that the principal quantum number, n, is 2, write down the allowed
values of l and ml , and determine the number of atomic orbitals possible for n
= 3.
8. Using the rules that govern the values of the quantum numbers n and l, write
down the possible types of atomic orbital for n = 1, 2 and 3.
9. Write down the possible types of atomic orbital for n = 4.
10.Which atomic orbital has values of n = 4 and l = 2?
11.Give the three quantum numbers that describe a 2s atomic orbital.
12.Which quantum number distinguishes the 3s and 5s atomic orbitals?
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13.Write down two possible sets of quantum numbers that describe an electron in
a 2s atomic orbital. What is the physical significance of these unique sets?
14.Write down two possible sets of quantum numbers to describe an electron in a
3s atomic orbital.
15.If an electron has the quantum numbers n = 2, l = 1, ml = −1 and ms = 
1
2
which type of atomic orbital is it occupying?
16.How many radial nodes does each of the following orbitals possess: (a) 2s; (b)
4s; (c) 3p; (d) 5d; (e) 1s; (f ) 4p?
17.Comment on differences between plots of R(r) against r, and 4r2R(r)2 against
r for each of the following atomic orbitals of an H atom: (a) 1s; (b) 4s; (c) 3p.
18.Write down the six sets of quantum numbers that describe the electrons in a
degenerate set of 5p atomic orbitals. Which pairs of sets of quantum numbers
refer to spin paired electrons?
19.For a neutral atom, X, arrange the following atomic orbitals in an approximate
order of their relative energies (not all orbitals are listed): 2s, 3s, 6s, 4p, 3p, 3d,
6p, 1s.
20. Using the concepts of shielding and penetration, explain why a ground state
configuration of 1s22s1 for an Li atom is energetically preferred over 1s22p1.
21.The bond dissociation enthalpies for the nitrogen–nitrogen bond in N2 and
[N2]− are 945 and 765 kJ mol−1 respectively. Account for this difference in
terms of MO theory, and state whether [N2]− is expected to be diamagnetic or
paramagnetic.
22.Account for the fact that [N2]+ is paramagnetic.
23.Using MO theory, rationalize why the N−N bond distance in [N2]+ is longer
(112 pm) than in N2 (109 pm).
24.Rationalize why the bond orders in [N2]+ and [N2]− are both 2.5.
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25.Show that N in NF3 obeys the octet rule.
26.Show that Se in H2Se obeys the octet rule.
27.In which of the following molecules is the octet rule apparently violated by the
central atom: (a) H2S; (b) HCN; (c) SO2; (d) CO2; (e) SO3?
28.Within the series of fluorides IF, IF3 and IF5, show that the octet rule appears to
be obeyed in only one instance for iodine.
29.Predict the structures of (a) XeF2 and (b) [XeF5]−.
30.Predict the structures of, BCl3 , [IF5]2− , [NH4]+ , SF6 , XeF4 , AsF5 , [BBr4]−.
31.Which of the following species are hydrogen-like: (a) H+; (b) He+; (c) He−; (d)
Li+; (e) Li2+?
32.Draw Lewis structures to describe the bonding in the following molecules: (a)
F2; (b) BF3; (c) NH3; (d) H2Se; (e) H2O2; (f) BeCl2; (g) SiH4; (h) PF5.
33.Use the Lewis structure model to deduce the type of nitrogen–nitrogen bond
present in (a) N2H4, (b) N2F4, (c) N2F2 and (d) [N2H5]+.
34.Use VB theory to describe the bonding in the diatomic molecules Li2, B2 and
C2. (b) Experimental data show that Li2 and C2 are diamagnetic whereas B2 is
paramagnetic. Is the VB model consistent with these facts?
35.Use MO theory to determine the bond order in each of [He2]+ and [He2]2+. (b)
Does the MO picture of the bonding in these ions suggest that they are viable
species?
36.Construct an MO diagram for the formation of O2; show only the participation
of the valence orbitals of the oxygen atoms. (b) Use the diagram to rationalize
the following trend in O−O bond distances: O2, 121 pm; [O2]+, 112 pm; [O2]−,
134 pm; [O2]2−, 149 pm. (c) Which of these species are paramagnetic?
37.Pick out pairs of isoelectronic species from the following list; not all species
have a ‘partner’: HF; CO2; SO2; NH3; PF3; SF4; SiF4; SiCl4; [H3O]+; [NO2]+;
[OH]−; [AlCl4]−.
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−
38.Use the VSEPR model to predict the structures of (a) H2Se, (b) [BH4] , (c)
+
−
+
NF3, (d) SbF5, (e) [H3O] , (f) IF7, (g) [I3] , (h) [I3] , (i) SO3, (g) SOF4.
39.Account for each of the following observations.
a. The first ionization energy of K is lower than that of Li.
b. BI3 is trigonal planar while PI3 is trigonal pyramidal in shape.
c. The second ionization energy of He is higher than the first despite the fact
that both electrons are removed from the 1s atomic orbital.
d. S2 is paramagnetic.
e. In the salt formed from the reaction of Br2 and SbF5, the Br–Br distance in
the Br2+ ion is 215 pm, i.e. shorter than in Br2.
40.Use VSEPR theory to account for the structure of NH3, and suggest an
appropriate hybridization scheme for the N atom.
41.Use VSEPR theory to account for the tetrahedral structure of [NH4]+.
42.Rationalize why H2O is bent but XeF2 is linear.
43.Give a suitable hybridization scheme for the central atom in each of the
following: (a) [NH4]+; (b) H2S; (c) BBr3; (d) NF3; (e) [H3O]+.
44.Draw a set of resonance structures (focusing only on those that contribute
significantly) for the nitrate ion. (c) Use an appropriate hybridization scheme
to describe the bonding in [NO3]−.
45.Use an appropriate hybridization scheme to describe the bonding in [BO3]3−.
46.Suggest an appropriate hybridization scheme for each central atom: (a) SiF 4;
(b) [PdCl4]2−; (c) NF3; (d) F2O; (e) [CoH5]4−; (f ) [FeH6]4−; (g) CS2; (h) BF3.
47.The I−I bond distance in I2 (gas phase) is 267 pm, in the [I3]+ ion is 268 pm,
and in [I3]− is 290pm (for the [AsPh4]+ salt). (a) Draw Lewis structures for
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these species. Do these representations account for the variation in bond
distance? (b) Use MO theory to describe the bonding and deduce the I−I bond
order in each species. Are your results consistent with the structural data?
48.Show that the structure of the unit cell for caesium chloride is consistent with
the formula CsCl.
49.MgO adopts an NaCl lattice. How many Mg2+ and O2− ions are present per
unit cell?
50.The unit cell of AgCl (NaCl type lattice) can be drawn with Ag + ions at the
corners of the cell, or Cl− at the corners. Confirm that the number of Ag + and
Cl− ions per unit cell remains the same whichever arrangement is considered.
51.With reference to the NaCl, CsCl and TiO2 lattice types, explain what is meant
by (a) coordination number, (b) unit cell, and (c) determination of the formula
of an ionic salt from the unit cell.
52.Determine the number of formula units of (a) CaF2 in a unit cell of fluorite,
and (b) TiO2 in a unit cell of rutile.
53.Confirm that the unit cell for perovskite is consistent with the stoichiometry
CaTiO3.
54. ReO3 is a structure-prototype. Each Re(VI) centre is octahedrally sited with
respect to the O2− centres. The unit cell can be described in terms of a cubic
array of Re(VI) centres, with each O2− centre at the centre of each edge of the
unit cell. Draw a representation of the unit cell and use your diagram to
confirm the stoichiometry of the compound.
55.Give explanations for the following observation: Raising the temperature of a
sample of -Fe from 298K to 1200K (at 1 bar pressure) results in a change of
coordination number of each Fe atom from 8 to 12.
56.Outline the similarities and differences between cubic and hexagonal closepacked arrangements of spheres, paying particular attention to (a) coordination
numbers, (b) interstitial holes and (c) unit cells.
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57.State the coordination number of a sphere in each of the following
arrangements: (a) ccp; (b) hcp; (c) bcc; (d) fcc; (e) simple cubic.
58.Gold (Au) crystallizes in a cubic close-packed structure (the face-centered
cube) and has a density of 19.3 g/cm3. Calculate the atomic radius of gold.
59.When silver crystallizes, it forms face-centered cubic cells. The unit cell edge
length is 408.7 pm. Calculate the density of silver.
60.The edge length of the NaCl unit cell is 564 pm. What is the density of NaCl
in g/cm3?
61.Copper crystallizes in a face-centered cubic lattice (the Cu atoms are at the
lattice points only). If the density of the metal is 8.96 g/cm3, what is the unit
cell edge length in pm?
62.Metallic iron crystallizes in a cubic lattice. The unit cell edge length is 287
pm. The density of iron is 7.87 g/cm3. How many iron atoms are within a unit
cell?
63.Europium crystallizes in a body-centered cubic lattice (the Eu atoms occupy
only the lattice points). The density of Eu is 5.26 g/cm3. Calculate the unit cell
edge length in pm.
With best wishes
Dr. Said M. El-Kurdi