ANSWERS – Exam 1 - 1998
1. (40 points) (a) Classify the solid forms of the following elements and compounds into one of the
following four types of crystals: (a) ionic, (b) extended covalent network, (c) metallic or (d)
molecular/atomic, and match up each group of elements/compound with the best available description [(i) >(iv)] listed below:
Ca, BaO, C, O, SiC, Bi, MgBr2, SO3, BaF2, F
(i) rel. low melting/boiling; transparent/colorless; low electrical conductivity; soft, mechanically weak
(ii) rel. high melting; transparent/colorless; low electrical conductivity in solid state; ionically conducting
melt; rel. hard, brittle.
(iii) high electrical and thermal conductivity; opaque/reflective; ductile.
(iv) high melting; mechanically strong/brittle; low to moderate electrical conductivity; opaque
(i)(d) O, SO3, F; (ii)(a) BaO, MgBr2, BaF2; (iii)(c) Ca, Bi; (iv)(b) C, SiC
(b) Choose the most ionic compound in this list and explain briefly the basis for your choice.
BaF2; F and Ba have the largest difference in electronegativity.
(c) For the (2) elements in this list that form descrite molecules (E and E'), give (i) the formula [En] for the
most stable molecular form for these elements and (ii) describe the bonding both by drawing a Lewis
structure and by giving a qualitative molecular orbital (MO) energy level diagram (filled in with the
valence electrons from each atom). Indicate the bond order and the number of any unpaired electrons in
the ground state that would be predicted by both the Lewis structure and the molecular energy level
diagram. Show all work for full credit.
O2 and F2; can use the F2 MO diagram (see text) for both: indicates a b.o of 2 for O 2 and 1 for F2 and
2 unpaired electrons for O2 (none for F2); Lewis structure for O2 also indicates a b.o. of 2 but no
unpaired electrons.
2. (30 points) (a) Write out the reactions that correspond to (i) the heat of formation, and (ii) the lattice
enthalpy, for cesium superoxide, CsO2(s) (a cesium salt of the [O2]- ion). Be sure to balance the
equations and to indicate the physical form of the species involved by mean of subscripts, such as (s), (l),
(g), (aq), etc.
(see answers to HW assign. 5 for this question)
(b) Using a Born-Haber type cycle, write an equation that would allow you to calculate an value for the
enthalpy of formation for the [O2]- ion [HEA [O2] ], if you were given the neccessary enthalpy terms.
Make a list of all of these enthalpy terms that would be needed for the calculation of HEA [O2] and for
each term write a balanced equation for the chemical reaction involved and indicate whether that enthalpy
value is likely to be positive, negative or approximately 0.
(ditto)
3. (25 points) (a) Assuming that they all had the same (rock salt) structure, arrange the following salts in
order of increasing lattice enthalpy and explain briefly the basis for your order of lattice enthalpies.
NaBr, BeS, BeO, MgS, CsBr
CsBr < NaBr < MgS < BeS < BeO
U is prop. to the charge on the ions squared over the sum of their radii.
The two Be2+ salts have the highest U because of the +2 charge on the
two ions (and the small size of the Be2+ ion), next is the Mg2+ and then
finally the two M+X- salts have the lowest U.
(b) Suppose that one of these salts actually had the cesium chloride structure (C.N. = 8) and another, the
zinc blende structure (C.N. = 4). Choose the most likely salts from this list to have these two other
structures and briefly explain the reason for your choice (assume that the anion is bigger than the cation in
each case)
CsBr would be most likely to have the CsCl structure, because the two
ions would be closest to having the same ionic radii, whereas BeS would
be most likely to have the lowest C.N. (zinc blende) structure, because
the two ions are the most different in size in this case (based on rel.
position of the atoms in the periodic table and the effect of the charge of
the ion on its size.
4. (15 points) Choose (i) the most soluble (in water) and (ii) the most thermally stable salt from the
following list. Explain briefly your choice for (i) and (ii):
NaClO4, KClO4, LiClO4, CsClO4, RbClO4
(i)
(ii)
LiClO4 (because it has the greatest difference in ionic radii)
CsClO4 (because it has the smallest difference in ionic radii)
5. (40 points)
(a) Give the formulas for the hydrides of the following elements:
Cl, Mg, Si, Rb, S, F, P, Ga
HCl, MgH2, SiH4, RbH, H2S, HF, PH3, GaH3
(b) Divide these hydrides into two groups, depending on whether they are "ionic" or "covalent" hydrides.
Ionic: MgH2, RbH,
Covalent: HCl, SiH4, H2S, HF, PH3, GaH3
(c) Choose one ionic hydride from your list and write a balanced equation for its reaction with water.
MgH2 + 2H2O  Mg(OH)2 + 2 H2
(d) Draw the Lewis structures and determine the molecular geometry (by using VSEPR) of all of the above
molecular hydrides that have two or more hydrogens attached to the central atom.
H
H
Si
H
H
S
H
P
H
H
H
H
H
Ga
H
Si
H
S
H
H
tetrahedral
H
H
P
H
bent (or
V shaped)
H
H
H
H
trigonal
planar
pyramidal
(e) Choose the hydride that would exhibit the strongest "hydrogen bonding" in the condensed state and briefly
describe at least one of the effects of this "hydrogen bonding" on the properties of this hydride (qualitatively
compare this hydride to the other hydrides of the same Group). Briefly explain the basis for your choice of hydrides
HF; the b.p. of HF is much higher than that of HCl, despite the increased size of the Cl atom; the remainder
of the HX compounds have b.p.s which increase with increasing size of the X atom (also HF (l) has a high
dielectric constant and a lower frequency (and broader) IR stretching band than does HF in the gas phase;
etc.)
6. (20 points) (a) Give the formulas for the hydrides of the following elements and then rank these hydrides in order
of increasing relative (protonic) acidity:
Al, Se, Si, P, Br, Mg
MgH2 < AlH3 < SiH4 < PH3 < H2Se < HBr
(b) Rank the following oxoacids in order of increasing relative (protonic) acidity:
HClO3, H2SeO3, HBrO3, H3AsO4, HClO4, H2SeO4
(for these oxoacids assume that all of the H's are attached to an oxygen)
H3AsO4 < H2SeO3 < H2SeO4 < HBrO3 < HClO3 < HClO4
7. (30 points) (a) Predict the directions of the following equilibria and explain your answer briefly,
identifying clearly (by formula and charge) the particular Lewis acids and bases involved in these two
reactions:
(i)
[Cl3N:AuCl3] + [Cl3Sb:CuCl3]- <==> [Cl3Sb:AuCl3] + [Cl3N:CuCl3]-
(ii)
[Cl3As:CuCl3]- + [CuCl4]3- <==> [Cl3As:CuCl3]2- + [CuCl4]2-
(i)
to the right (Au softer than Cu; Sb softer than N); Lewis bases: Cl 3N: and Cl3Sb:
(ii)
[AuCl3] + [CuCl3]to the right [Cu(I) is softer than Cu(II), Cl- harder than Cl3As:]; Lewis bases: Cl3As: and Cl-;
Lewis acids: [CuCl3]- + [CuCl3]2-
; Lewis acids:
(b) Classify each of the following metal ions according to whether they would be most likely to be found as
(or in) (i) oxides, carbonates or silicate minerals (lithophiles) or as (ii) sulfides (chalcophiles).
Al3+, Ag+, Cs+, Zr4+, Hg2+, Be2+
(i)
lithophiles: Al3+, Cs+, Zr4+, Be2+; (ii) chalcophiles: Ag+, Hg2+
(c) Which two of the following metal ions are most likely to behave as poisons towards animals (or humans)?:
K+, Ca2+, Ag+, Al3+, Cd2+, Si4+
Ag+ and Cd2+
(d) Give the formula for the (one) product that will precipitate when the following compounds are added to
1 M HCl(aq) (in the absence of light):
AuF(s) + AuF3(s) + H2S(g) + CsF(s)
Au2S(s