SJCC
Chem 1A Homework
Handout
Prof. John C. Song
Compilation of Homework Assignments for General Chemistry 1A as
supplement lecture materials at SJCC .
2
Homework (Matter and Measurement):
1.
Characterize each of the following dartboards according to the ACCURACY and PRECISION of the
results for each of the below:
(a) Which one of the above is the most accurate? Which one the least accurate?
(b) Which of the above is/are the most precise? Which the least precise?
2.
From looking at the periodic table, list the names and symbols of four elements of metal,
nonmetal and metalloid respectively and indicate where they gnerally lie in the periodic table.
3.
What are the symbols for t he following elements?
4.
(a)
Cadmium (used iin rechargeable batteries)
(b)
Beryllium (used in the space shuttle)
(c )
Germanium (used in semiconductors)
(d)
Arsenic (ued in pesticides)
Give the names corresponding to the following symbols:
(a) Re
(b) Pu
(c) Rh
(d) Ga
5.
The density of sulfuric acid, H2SO4, is 15.28 lb/gal. What is the density of sulfuric acid in g/mL?
6.
Vignaigrette salad dressing consists mainly of oil and vinegar. The desity of olive oil is 0.918
g/cm3, the density of vinegar is 1.006 g/cm3, and the two do not mix. If a certain mixtture of
olive oil and vinegar has a total mass of 397.8 g and a total volume of 422.8 cm3, what is the
volume of oil and what is the volume of vinegar in the mixture?
3
7.
A solid substance has a building unit that has a cubical "Unit Cell" with a unit length of 4.25 Ǻ
and density of 12.37 g/cm3. (Given: 1 Ǻ = 1 x 10-10 m)
(a)
Express the above density in the unit of kg/m3.
(b)
What is the mass of the "Unit Cell" in nanograms (ng)"
4
Homework (Atomic Theory, Atoms, Molecules, and Ions):
1.
In addition to carbon monoxide (CO) and carbon dioxide (CO2), there is a third compound of
carbon and oxygen carbon suboxide. If a 2.500 g sample of carbon suboxide contains 1.32 g of C
and 1.18 g of O, show that the law of multiple proportions is followed.
2.
Benzene, ethane, and ethylene are just three of a large number of hydrocarbons - compounds
that contain only carbon and hydrogen. Show how the following data are consistent with the
law of multiple proportions.
Compound
Mass of Carbon
In 5.00 g sample
4.61 g
4.00 g
4.29 g
Benzen e
Ethane
Ethylene
Mass of Hydrogen
In 5.00 g sample
0.39 g
1.00 g
0.71 g
3.
There are two binary compounds of titanium and chlorine. One compound contains 31.04 %
titanium by mass, and the other compound contains 74.76% chlorine by mass. What are the
ratios of titanium and chlorine atoms in the two compounds?
4.
What is the difference between an element's atomic number and its atomic mass?
5.
What is an isotope?
6.
Complete the following table:
Nucleus
Number of
Protons
Number of
Electrons
Number of
Neutrons
Nucleid
Symbol ( X )
N-15
Co-60
Pb-207
7.
Naturally occurring silver consists of two isotopes: Ag-107 (51.84%) with an isotopic mass of
106.9051 amu and Ag-109 (48.16%) with an isotopic mass of 108.9048 amu. What is the atomic
mass of silver? Check your answer in a periodic table.
8.
A sample of naturally occurring silicon consists of Si-28 (27.9769 amu), Si-29 (28.9765 amu) , and
Si-30 (29.9738 amu). If the atomic mass of silicon is 28.0855 amu and the natural abundance of
Si-29 is 4.67%, what is the natural abundances of Si-28 and Si-30?
5
9.
Which of the following bonds are likely to be covalent and which ionic? Explain,
(a) B Br
(b) Na
(c) Br Cl
(d) O Br
10.
Which of the following mixtures are homogeneous and which are heterogeneous?
(a) 18 karat gold
(d) clean air
11.
(b) salt and sand
(e) salt solution
Classify the following compounds as ionic or covalent:
(a) Iron (II) cyanide
(b) Calcium oxide
(c) Barium chloride
(d) Carbon monoxide
(e) Silicon dioxide
(f) Carbon tetrachloride
12.
Write formulas for the following compounds:
(a) Calcium acetate
(b) Sodium dichromate
(c) Chromium (III) sulfate
(d) Mercury (II) perchlorate
13.
14.
Name the following compounds:
(a) CCl4
(b) ClO2
(d) N2O3
(e) SO3
Write formulas for the following compounds:
(a) Calcium hydrogen sulfate
(b) Ruthenium(III) nitrate
(c) Ammonium carbonate
(d) Hydroiodic acid
(c) N2O
(c) A carbonated drink
6
Homework (Chemical Equations and Chemical Stoichiometry):
1.
Balance the following equations:
(a) Mg + HNO3 → H2 + Mg(NO3)2
(b) CaC2 + H2O → Ca(OH)2 + C2H2
(c) S + O2 → SO3
(d) UO2 + HF → UF4 + H2O
2.
How many moles are in one gram of each of the following substances?
(a) Cr
(b) Au
3.
How many moles of ions are in 2.5 mol of NaCl?
4.
How many moles of ions are I n 27.5 g of MgCl2?
5.
How many moles of cations are in 1.45 mol of K2SO4?
6.
How many moles of anions are in 35.6 g of AlF3?
7.
An average cup of coffee contains about 125 mg of caffeine, C8H10N4O2.
(a) How many moles of caffeine are in a cup?
(b) How many molecules of caffeine are present?
8.
An alternate method of preparing pure iron from Fe2O3 is by reaction with carbon monoxide:
Fe2O3 + CO → Fe + CO2
(unbalanced)
(a) Balance the equation.
(b) How many grams of CO are needed to react with 3.02 g of Fe2O3?
(c) How many grams of CO are needed to react with 1.68 mol of Fe2O3?
9.
Cisplatin [Pt(NH3)2Cl2], a compound used in cancer treatment, is prepared by reaction of
ammonia with potassium tetrachloroplatinate:
K2PtCl4 + 2 NH3 → 2 KCl + Pt(NH3)2Cl2
How many grams of cisplatin are formed from 55.8 g of K2PtCl4 and 35.6 g of NH3 if the reaction
takes place in 95% yield based on the limiting reactant?
7
10.
Vitamin C (ascorbic acid) contains 40.92% C, 4.58% H, and 54.50% O by mass.
(a) What is the empirical formula of ascorbic acid?
(b) If the molecular mass of ascorbic acid is 176 amu, what is the molecular formula?
11.
Aspirin contains C, H, and O in the following mass percents: C, 60.00%; H, 4.48%; and O, 35.52%.
Molecular mass of aspirin is 180 amu. What is the molecular formula of aspirin?
12.
Menthol, a flavoring agent obtained from pepperpint oil, contaiins carbon, hydrogen, and
oxygen. On combustion analysis, 1.00 g of menthol yields 1.161 g of H2O and 2.818 g of CO2.
What is the empirical formula of menthol?
13.
Convert the following percent composition into molecular formulas:
(a) Diborane: H 21.86%, B 78.14%; Molecular mass = 27.7 amu
(b) Trioxan:
14.
C 40.00%, H 6.71%, O 53.28%; Molecular mass = 90.08 amu
Lead Iodide (PbI2) can be prepared according to the following Double Replacement chemical
reaction:
Pb(NO3)2(aq) + 2 KI(aq) → PbI2(s) + 2 KNO3(aq)
(Given: molar mass of PbI2 = 461 g/mol)
(a) How many mL of 0.10 M Pb(NO3)2 is required to completely react with 50.0 mL of 0.10 M KI,
assuming the reaction goes to completion?
(b) If you react 25.0 mL of 1.0 M Pb(NO3)2(aq) with 25.0 mL of 1.0 M KI (aq) solution, how
much PbI2(s) will precipitate in moles?
(c) Which reactant is acting as the limiting reactant in the above reaction (b)? and give
explanation for your answer.
15.
Tartaric acid (H2C4H4O6) is the acid present in wine and can be neutralized by NaOH as follows:
H2C4H4O6 (aq) + 2 NaOH(aq) → Na2C4H4O6 (aq) + 2 H2O(l)
If 50.00 mL of wine required 12.6 mL of 0.1025 M NaOH (aq) solution to the endpoint,
(a) What is the concentration of tartaric acid in wine in Molarity?
(b) Express the above concentration of tartaric acid in wine in % (mass/volume).
(c) What is the net amount of tartaric acid in milligrams present in one 4.0 fl.oz of wine ?
16.
A commercially available sodium hydroxide (NaOH) is 50.0% (mass/mass) and its density is 1.32
g/mL. What is the Molarity (M) of the NaOH (aq)?
8
Homework (Atomic Structure & Periodicity):
1.
Calculate the energies of the following waves (in kilojoules per mole), and tell which member of
each pair has the higher value.
(a) An FM radio wave at 99.5 MHz and an AM radio wave at 1150 kHz.
(b) An X-ray with λ = 3.44 x 10-9 m and a microwave with λ = 6.71 x 10-2 m.
2.
At what speed (in meters per second) must a 145 g baseball be travelling to have a de Broglie
wavelength of 0.500 nm?
3.
Lines in the Brackett series of the hydrogen spectrum are caused by emission of energy
accompanying the fall of an electron from outer shells to the fourth shell. The lines can be
calculated using the Balmer-Rydberg equation:
=R
where m = 4, R = 1.097 x 10-2 nm-1, and n is an integer greater than 4.
(a) Calculate the wavelengths (in nanometers).
(b) Calculate energies (in kilojoules per mole) of the first two lines in the Brackett series.
(c) In what region of the electromagnetic spectrum do they fall respectively?
4.
Give the allowable combinations of quantum numbers for each of the following electrons:
(a) A 4s electron
(b) A 3p electron
(c) A 5f electron
(d) A 5d electron
5.
Give the orbital designations of electrons with the following quantum numbers:
(a) n = 3, l = 0, ml = 0
(b) n = 2, l = 1, ml = -1
(c) n = 4, l = 3, ml = -2
(d) n = 4, l = 2, ml =0
9
6.
Which of the following combinations of quantum numbers can refer to an electron in a groundstate cobalt atom (Z=27)?
(a) n = 3, l = 0, ml = 2
(b) n = 4, l = 2, ml = -2
(c) n = 3, l = 1, ml = 0
7.
Give the expected ground-state electron configurations for atoms with the following atomic
numbers:
(a) Z = 55
8.
(b) Z = 40
(c) Z = 80
(d) Z = 62
Draw orbital-filling diagrams for the following atoms. Show each electrons as an up or down
arrow, and use the abbreviation of the preceding noble gas to represent inner-shell electrons.
(a) Rb
(b) W
(c) Ge
(d) Zr
9.
What is the expected ground-state electron configuration of the recently discovered element
with Z =116?
10.
Following questions are for the ground state of Gallium (Ga). (atomic no. 31).
(a)
Write the full electron configuration for the Ga . (1s2, 2s2, etc.)
(b)
List the set of four (4) quantum numbers (n, l, ml, ms), for all the valence electron(s) of
Ga.
11.
Why do atomic radii decrease from left to right across a period of the periodic table?
12.
The amount of energy that must be added to remove an electron from a neutral atom to give a
cation is called the atom's ionization energy. Which would you expect to have the larger
ionization energy, Na or Mg? Explain.
13.
Why do halogens in group 7A have much higher electron affinities than metals in group 1A and
noble gases in group 8A?
14.
(a)
Calculate the “Ionization Energy” of Hydrogen atom in kJ/mol. (Given: E = - RH/n2,
where RH = 2.18 x 10-18 J, n = principal quantum number,Avogadro No.= 6.02 x 1023).
(b)
Using the Rydberg equation given below, calculate the wavelength of light
corresponding to the emission line n = 6 to m = 2) for hydrogen atom in nm. (Given:,
=R
where m, n = principal quantum numbers, R = 1.097 x 10-2 nm-1)
10
15.
The amount of energy necessary to remove an electron from an atom is a quantity called the
ionization energy, Ei. This energy can be measured by a technique called photoelectron
spectroscopy, in which light of wavelength λ is directed at an atom, causing an electron to be
ejected. The kinetic energy of the ejected electron (Ek) is measured by determining its velocity, v
(Ek = mv2/2), and Ei is then calculated using the conservation of energy principle. That is, the
energy of the incident light equals Ei plus Ek. What is the ionization energy of selenium atoms (in
kilojoules per mole) if light with λ = 48.2 nm produces electrons with a velocity of 2.371 x 106
m/s? The mass, m, of an electron is 9.109 x 10-31 kg.
16.
Ionization energy (IE) of an atom can be measured by photoelectron spectroscopy described in
no. 15 above. What is the threshold wavelength of the light beam in nm to ionize the cesium
atom if the ionization energy of cesium is 375.7 kJ/mol, kinetic energy (KE) of electron ejected is
1646 kJ/mol. (Given: hν = IE + KE, h = 6.63 x 10-34 J.s., c = 3.00 x 108 m/s, Avogadro No. =
6.022 x 1023)
11
Homework (Chemical Bonds: Ionic & Covalent):
1.
Find the lattice energy of LiBr from the Table in Text, and calculate the energy change (in
kilojoules per mole) for the formation of solid LiBr from the elements, i.e. Li(s) + 1/2 Br2(l) →
LiBr(s) (Given: sublimation energy for Li = +159.4 kJ/mol, bond dissociation energy for Br2(g) =
+224 kJ/mol, and heat of vaporization of Br2(l) to Br2(g) = 30.9 kJ/mol.)
2.
Define Lattice Energy and give an example with its lattice energy in kJ/mol.
3.
Which of the two has the larger lattice energy and explain your choice.
(a) LiF(s) vs. LiI(s)
(b) NaCl(s) vs. AlCl3(s)
4.
Which of the following substances are largely ionic and which are covalent?
(a) HI
(b) BBr3
(c) NaOH
(d) CH3Li
(e) PdCl2
5.
Define ELecronegativity and Order the following elements according to decreasing
electronegativity: C, Ca, Cs, Cl, Cu
6.
Use the electronegativity table to predict which bond in each of the following pairs is more
polar and for the more bond you have chosen, show the direction of polarity of the bond using
the → and δ+/δ- notations.
(a) C __ H or C __ Cl
(b) Si __ Li or Si__ Cl
(c) N__ Cl or N__ Mg
7.
8.
Draw electron-dot structures for the following molecules or ions:
(a) CBr4
(b) NCl3
(c) C2H5Cl
(d) BF4-
(e) O22-
(f) NO+
Draw elecron-dot structures for the following molecules,, which contain atoms from the third
row or lower.
(a) SbCl3
(e) H3PO4
(b) KrF2
(f) SeOCl2
(c)ClO2
(d) PF5
12
9.
Draw as many resonance structures as you can for the following nitrogen-containing
compounds:
(a) N2O
(b)NO
(d) N2O3 (use O-N-NO2 atomic arrangement)
10.
(c)NO2
Which of the following pairs of structures represent resonance forms, and which do not?
(a)
(b)
(c)
C
H
N
..
O:
..
____
: O:
. .___ ___ . .
S
O:
:O
..
..
..
: O:
and
..
___ . .
C ___ N____ O :
..
..
:O :
..
..
___ ___
O:
:O S
..
..
H
and
H
C
:O:
and
C
C
H
H
C:
H
H
H
(d)
H
H
H
H
H
H
H
H
and
H
H
H
11.
H
Assign formal charges to the atoms in the following structures:
H
H
(a)
H
..
O
..
N
..
H
(b)
H
. . ___
N
C
..
H
..
:O:
(c)
..
: Cl
..
P
: Cl :
..
..
Cl :
..
H
13
12.
Calculate formal charges for the C and O atoms in the following two resonance structures.
Which structure do you think is the more important contributor to the resonance hybrid?
Explain.
..
: O:
: O:
C
H
..
C
H
C
H
H
13.
H
How many charge clouds are there around the central atom in molecules that have the
following geometry?
(a) Tetrahedral
(d) Linear
14.
H
C
<---->
(b) Octahedral
(e) Square pyramidal
(c) Bent
(f) Trigonal pyramidal
What bond angles do you expect for each of the following?
(a) The Cl-P-Cl angle in PCl6(b) The Cl-I-Cl angle in ICl2(c) The O-S-O angle in SO42(d) The O-B-O angle in BO33-
15.
Use the MO diagram in the Text to describe the bonding in O2+, O2, and O2-.
(a) Which of the three should be stable?
(b) What is the bond order of each?
(c) Which contain unpaired electrons?
16.
(a) Draw three possible resonance structures for methyl isocyanate, CH3-N-C-O, a deadly toxic
gas intermediate to produce carbamate pesticides.
(b) Assign formal charges to the atoms in each resonance structure.
14
17.
In the cyanate ion, OCN-, carbon is the central atom.
(a) Draw as many resonance structures as you can for OCN-, and assign formal charges to the
atoms in each.
(b) Which resonance structure makes the greatest contribution to the resonance hybrid?
Which makes the least contribution? Explain.
(c) Is OCN- linear or bent? Explain.
(d) Which hybrid orbitals are used by the C atom, and how many π bonds does the C atom
form?
18.
(a) Draw as many resonance structures as you can for N2O (Dinitrogen
monoxide, laughing gas) and assign formal charges to the atoms in each.
(b) Choose the most important contributor among the above resonance
hybrids and provide reason(s) why.
15
19.
For the following molecules and ions:
(a)
Draw a Lewis electron dot structure for each of the following
(show all bonding and nonbonding electrons);
(b)
Indicate electronic geometry about the central atom for each;
(e.g. linear, bent, etc. No drawings please.)
(c)
Indicate molecular geometry about the central atom for each:
(e.g. linear, bent, etc. No drawings please.)
Molecule
SO2
IF5
Lewis Str.
Electronic
Molecu.
(a)
Geom. (b)
Geom. (c)
16
Homework (Thermochemistry):
1.
A reaction is carried out in a cylinder fitted with a moving piston, as shown here. The starting
volume is V = 5.00 L, and the apparatus is held at constant temperature and pressure. Assuming
that ΔH = -35.0 kJ and ΔE = -34.8 kJ, redraw the piston to show its position after reaction. Does
volume increase, decrease, or remain the same?
1 atm
V = 5.00 L
2.
Used in welding metals, the reaction of acetylene with oxygen has ΔHº = -1255.5 kJ:
C2H2(g) + 5/2 O2(g) → H2O(g) + 2 CO2(g)
ΔHº = -1255.5 kJ
How much PV work is done (in kJ), and what is the value of ΔE (in kJ) for the reaction of 6.50 g of
acetylene at atmospheric pressure if the volume change is -2.80 L?
3.
How much heat (in kJ) is evolved or absorbed in the reaction of 2.50 g of Fe2O3 with enough
carbon monoxide to produce iron metal? Is the process exothermic or endothermic?
Fe2O3(s) + 3 CO(g) → 2 Fe(s) + 3 CO2(g)
ΔHº = -24.8 kJ
4.
When 0.187 g of benzene, C6H6, is burned in a bomb calorimeter, the surrounding water bath
rose in temperature by 7.48ºC. Assuming that the bath contains 250.0 g of water and that the
calorimeter itself absorbs a negligible amount of heat, calculate combustion energies ΔE for
benzene in both kilojoules per gram and kilojoules per mole.
5.
When a solution containing 8.00 g of NaOH in 50.0 g of water at 25.0ºC is added to a solution of
8.00 g of HCl in 250.0 g of water at 25.0ºC in a calorimeter, the temperature of the solution
increases to 33.5ºC. Assuming that the specific heat of the solution is 4.18 J/(g•ºC) and that the
calorimeter absorbs a negligible amount of heat, calculate ΔH (in kJ) for the reaction
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
When the experiment is repeated using a solution of 10.00 g of HCl in 248.0 g of water, the
same temperature is observed. Explain.
17
6.
The standard heat of reaction for the following reaction as shown is -2219 kJ.
C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(l)
ΔHf for CO2(g) = -394 kJ/mol, ΔHf for H2O(l) = -286 kJ/mol
Calculate the ΔHf (in kJ/mol) for propane (C3H8(g)). (Caution: watch sign!)
7.
Given the following thermochemical equations, determine the standard heat of reaction (H)
for 2Al(s) + Fe2O3(s)  Al2O3(S) + 2Fe(s) in kJ.
(Given: 2Fe (s) +
2Al(s) +
8.
O2(g)  Fe2O3(s)
Hf = -824 kJ/mol
O2(g)  Al2O3(s)
Hf = -1676 kJ/mol )
Use the bond dissociation energy table in the Text to calculate an approximate ΔHº (in
kilojoules) for the industrial synthesis of isopropyl alcohol (rubbing alcohol) by reaction of water
with propene, as shown below.
OH
H3C CH
9.
CH2
+
H3C CH CH3
Isopropyl alcohol
H2O
Tell whether the entropy changes, ΔS, for the following processes are likely to be positive or
negative.
(a) The conversion of water liquid water to water vapor at 100ºC.
(b) The freezing of liquid water to ice at 0ºC.
(c) The eroding of a mountain by a glacier.
10.
Tell whether reactions with the following values of ΔH and ΔS are spontaneous or
nonspontaneous and whether they are exothermic or endothermic:
(a) ΔH = -128 kJ; Δ S = 35 J/K at 500 K
(b) ΔH = +67 kJ; ΔS = -140 J/K at 250 K
(c) ΔH = +75 kJ; ΔS = 95 J/K at 800 K
11.
Imagine that you dissolved 10.0 g of a mixture of NaNO3 and KF in 100.0 g of water and find that
the temperature rises by 2.22ºC. Using the following data, calculate the mass of each compound
in the original mixture. Assume that the specific heat of the solution is 4.18 J/(g •ºC).
NaNO3(s) →
KF(s) →
NaNO3(aq)
KF(aq)
ΔH = +20.5 kJ/mol
ΔH = -17.7 kJ/mol
18
Homework (Gases):
1.
Carry out the following conversions:
(a) 352 torr to kPa
(b) 0.255 atm to mm Hg
(c) 0.0382 mm Hg to Pa
2.
The height of the mercury in the right arm open to atmospheric pressure (760 mmHg) is 100
mm and the height in the left arm is 120 mm.
gas
120 mm
100 mm
What is the pressure of the gas in the bulb?
(A)
20 mmHg
(C)
740 mmHg
(B)
640 mmHg
(D)
780 mmHg
3.
Assume that you have an open-end manometer filled with ethyl alcohol (density = 0.7893 g/mL
at 20ºC) rather than mercury (density = 13.546 g/mL at 20ºC). What is the pressure (in pascals) if
the level in the arm open to the atmosphere is 55.1 cm higher than the level in the arm
connected to the gas sample and the atmospheric pressure reading is 752.3 mm Hg?
4.
A compressed air tank carried by scuba divers has a volume of 8.0 L and a pressure of 140 atm at
20ºC. What is the volume of air in the tank (in liters) at 0ºC and 1 atm?
5.
One mole of any gas has a volume of 22.414 L at 0 C and 1 atm. What are the densities of the
following gasses (in grams per liter) at the same Temperature and Pressure?
(a) CH4
6.
(b) CO2
(c) O2
(d) UF6
An unidentified gas has a density of 5.380 g/L at 23°C and 760 mmHg pressure. What is the
molar mass of the unknown gas? (Given PV = nRT; R = 0.0821 L-atm/mol.k)
19
7.
Ammonium nitrate can decompose explosively when heated according to the equation
2 NH4NO3(s) →
2 N2(g) + 4 H2O(g) + O2(g)
How many liters of gas would be formed at 450ºC and 1.00 atm pressure by explosion of 450 g
of NH4NO3?
8.
Natural gas is a mixture of many substances, primarily CH4, C2H6, C3H8, and C4H10. Assuming
that the total pressure of the gases is 1.48 atm and that their mole ratio is 94: 4.0: 1.5 : 0.50,
calculate the partial pressure (in atmospheres) of each gas.
9.
A 30.0 L flask contains a binary mixture of 10.00 g of Helium (He) gas and 10.00 g of Argon
(Ar) gas at 25ºC. (PV = nRT; R= 0.0821 L-atm/mol.K)
(a) What are the partial pressures of He and Ar gas in atmosphere?
(b) What is the total pressure of a mixture in atmosphere?
(c) What is the mole fraction of Ar gas in the mixture?
10.
Magnesium metal reacts with aqueous HCl to yield H2 gas:
Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g)
The gas that forms is found to have a volume of 3.557 L at 25ºC and a pressure of 747 mm Hg.
Assuming the gas is saturated with water vapor at a partial pressure of 23.8 mm Hg, what is the
partial pressure (in torrs) of the H2? How many grams of magnesium metal were used in the
reaction?
11.
Which gas will diffuse through a membrane more rapidly, CO or N2? Assume that the samples
contain only the most abundant isotopes of each element, 12C, 16O, and 14N.
12.
When a gaseous compound X containing only C, H, and O is burned in O2, 1 volume of the
unknown gas reacts with 3 volumes of O2 to give 2 volumes of CO2 and 3 volumes of gaseous
H2O. Assume all volumes are measured at the same temperature and pressure.
(a) Calculate a formula for the unknown gas, and write a balanced equation for the combustion
reaction.
(b) Is the formula you calculated an empirical formula or a molecular formula? Explain.
(c) Draw two different possible electron-dot structures for the compound X.
(d) Combustion of 5.000 g of X releases 144.2 kJ of heat. Look up ΔHºf values of CO2(g) and
H2O(g) in Thermochemical Data table in Text, and calculate ΔHºf for compound X.
20
Homework (Liquids, Solids and Phase Changes):
1.
Methanol (CH3OH; bp = 65 C) boils nearly 230º higher than methane (CH4; bp = -164ºC), but 1decanol (C10H21OH; bp = 229ºC) boils only 55º higher than decane (C10H22; bp = 174ºC). Explain.
2.
Draw three-dimensional structures of PCl3 and PCl5 and then explain why one of the molecules
has a dipole moment and one does not.
3.
Which one of the following substances would you expect to have a nonzero dipole moment?
Explain, and show the direction of each.
4.
(a) NF3
(b) CH3NH2
(c) XeF2
(d) PCl5
Dichloroethane, CH2Cl2. Is an organic solvent used for removing caffeine from coffee beans. The
following table gives the vapor pressure of dichloroethane at various temperatures. Fill in the
rest of the table and use the data to plot curves of Pvap versus T and ln Pvap versus 1/T.
Temp (K)
Pvap (mm Hg)
ln Pvap
1/T
263
80.1
?
?
273
133.6
?
?
283
213.3
?
?
293
329.6
?
?
303
495.4
?
?
313
724.4
?
?
5.
Ethanol has vapor pressure of 100 torr at 34.7ºC and the vapor pressure of 386 torr at 65.0ºC.
What is the heat of vaporization, ΔHvap, for ethanol in kilojoules per mole?
(Given: ln Pvap = (-ΔHvap/RT) + C, R=8.314 J/mol.K, C = constant)
6.
Copper metal crystallizes in a face-centered cubic structure. X-ray crystallography shows the
unit cell length a of 3.61 Å.
(a) Calculate the theoretical density of Copper metal in g/cm3. (1 Å = 1.0 x 10-10 m)
(b) What is the packing efficiency of the copper?
(packiing effciency = volume occupied by copper atoms per unit cell/volume of a
unit cell)
7.
Cesium chlroide (CsCl) crystallizes in a modified primitive cubic structure in which z+ = 1 and z- =
1. X-ray crystallography shows the unit cell length a of 4.11 Å.
(a) Calculate the theoretical density of CsCl in g/cm3. (1 Å = 1.0 x 10-10 m)
(b) Assuming the radius of Cs+ ion = radius of Cl- ion, what is the packing efficiency of the CsCl
ionic crystal structure.
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8.
Silver metal crystallizes in a face-centered cubic unit cell with an edge length of 408 pm. The
molar mass of silver is 107.9 g/mol, and its density is 10.50 g/cm3. Use these data to calculate
a value for Avogadro’s number.
9.
Refer to the phase diagram of CO2 in the Text.
(a) Under what state will CO2 exist at 6 atmosphere, and at -10º C?
(b) Under what pressure will CO2 liquify at -57º C?
(c) What is the maximum temperature under which CO2 gas can be liquefied?
10.
Oxygen has Tt = 54.3 K, Pt = 1.14 mm Hg, Tc=155 K, and Pc = 49.8 atm. The density of the liquid is
1.14 g/cm3, and the density of the solid is 1.33 g/cm3. Sketch a phase diagram for oxygen, and
label all points of interest.
11.
Refer to the phase diagram of oxygen in Number #10 above, and trace the following path
starting from a point at 0.0011 atm and -225C:
(a) First, increase P to 35 atm while keeping T constant.
(b) Next, increase T to -150C while keeping P constant.
(c) Then, decrease P to 1 atm while keeping T constant.
(d) Finally, decrease T to -215C while keeping P constant.
What is your starting phase, and what is your final phase?
22
Homework (Solutions and Properties):
1.
The dissolution of NH4ClO4(s) in water is exothermic, with ΔHsoln = +33.5 kJ/mol. If you prepare a
1.00 m solution of NH4ClO4 beginning with water at 25.0ºC, what is the final temperature of the
solution (in ºC)? Assume that the specific heats of both pure H2O and the solution are the same,
4.18 J/(K•g).
2.
How would you prepare each of the following solutions?
(a) 100 mL of a 155 ppm solution of urea, CH4N2O, in water.
(b) 100 mL of an aqueous solution whose K+ concentrataiton is 0.075 M.
3.
What is the mass percent concentration of the following solutions?
(a) Dissolve 0.655 mol of citric acid, C6H8O7, in 1.00 kg of water.
(b) Dissolve 0.135 mg of KBr in 5.00 mL of water
(c) Dissolve 5.50 g of aspirin, C9H8O4, in 145 g of dichloromethane, CH2Cl2.
4.
What is the molality of each solution prepared in Problem 3 above?
5.
A 0.944 M solution of glucose, C6H12O6, in water has a density of 1.0624 g/mL at 20C. What is
the concentration of this solution in the following units?
(a) Mole fraction
(b) Mass percent
(c) Molality
6.
Fish generally need an O2 concentration in water of at least 4 mg/L for survival. What partial
pressure of oxygen above the water (in atmospheres at 0ºC) is needed to obtain this
concentration? The solubility of O2 in water at 0ºC and 1 atm partial pressure is 2.21 x 10-3
mol/L.
7.
What is the vapor pressure (in mm Hg) of the following solutions, each of which contains a
nonvolatile solute? The vapor pressure of water at 45.0ºC is 71.93 mmHg.
(a) A solution of 10.0 g of urea, CH4N2O, in 150.0 g of water at 45.0ºC.
(b) A solution of 10.0 g of LiCl in 150.0 g of water at 45.0ºC, assuming complete dissolution.
8.
What is the freezing point in ºC of each of the solutions in Problem 7? (Kf for water = 1.86
C/molal)
23
9.
What is the normal boiling point in ºC of each of the solutions in Problem 7? (Kb for water = 0.51
C/molal)
10.
What osmotic pressure in atmosphere would you expect for 5.00 g of NaCl in 350.0 mL of
aqueous solution at 50ºC?
11.
Molecular weight of a carotene in carrot can be determined by measuring the osmotic
pressure. If 7.68 mg of carotene is dissolved in 10.0 mL of chloroform, the osmotic pressure was
26.57 mmHg at 25 C. What is the molecular weight of the caroten2?
12.
Calculate the freezing point of 500. g of water when 25.0 g of CaCl2 is dissolved in it. (assume
complete dissociation of the solute.) (Kf(H2O) = 1.86 C/m)
13.
A non-volatile compound of carbon, hydrogen and oxygen only was burned in excess oxygen
supply. 1.00 g of the compound produced 1.465 g of CO2 and 0.599 g of H2O upon complete
combustion. In another experiment, 1.00 g of the same unknown compound is dissolved in 50.0
g of water and its freezing point was measured at -0.248°C. (Given: ∆Tf = Kf x m; Kf for water is
1.86°C/m; m = molality)
(a) What is the molar mass of the unknown compound in grams per mole?
(b) Determine the molecular formula of the compound.