19 SUPPLEMENTAL PROBLEMS FOR CHEM 110 The problems are divided into chapters matching the book’s chapter numbers, and the chapters are arranged in order in which they are covered in lectures (see the syllabus). Ordered according to coverage: Chapter page 1 21 2 21 6 24 7 26 8 28 Organic 30 3 22 9 32 10 35 18 44 11 37 4 23 13 40 Reactions 45 5 23 15 42 Answer key 53 Ordered by chapter number: Chapter page 1 21 2 21 3 22 4 23 5 23 6 24 7 26 8 28 Organic 30 9 32 10 34 11 37 13 40 15 42 18 44 Reactions 45 Answer key 53
Transcript
19
SUPPLEMENTAL PROBLEMS
FOR CHEM 110
The problems are divided into chapters matching the book’s chapter numbers, and the chapters are arranged in order in which they are covered in lectures (see the syllabus).
Ordered according to coverage:
Chapter page 1 21 2 21 6 24 7 26 8 28
Organic 30 3 22 9 32
10 35 18 44 11 37
4 23 13 40
Reactions 45 5 23
15 42 Answer key 53
Ordered by chapter number:
Chapter page
1 21 2 21 3 22 4 23 5 23 6 24 7 26 8 28
Organic 30 9 32
10 34 11 37 13 40 15 42 18 44
Reactions 45 Answer key 53
20
21
Chapter 1 1. 13.2 nm is equal to
A. 1.32 × 10−6 cm
B. 1.32 × 10−6 mm
C. 1.32 × 10−8 cm
D. 1.32 × 10−8 mm
E. None of the above answers is correct 2. A 112.6 g mass of unknown material is submersed in
102 mL of water to yield a final volume of 126 mL. The apparent density of the unknown material is
A. 0.908 g/mL B. 1.10 g/mL C. 0.21 g/mL D. Not determinable from the information given E. None of the above answers is correct
3. The answer to the calculation
( )( )( ) 12949552.20925.8
6.413.180.9463.1025.1=
!+
to the correct number of significant figures is A. 20 B. 2.01 × 101 C. 2.0 × 101 D. 20.13 E. 2.012 × 101
The above calculation, when expressed to the correct number of significant figures and properly rounded, should be written A. 4.4 × 106 B. 4.40 × 106 C. 4.41 × 106 D. 4.4090 × 106 E. 4.4091 × 106
Chapter 2 1. Chlorine exists primarily as 37Cl and 35Cl and has an
average atomic weight of 35.453. The abundance of Cl with atomic weight of 35 is
A. 23% B. 91% C. 77% D. 45% E. None of the above is within 10% of the correct answer
2. How many protons, neutrons and electrons are
present in −2178O ?
A. p=8, n=9, e=10 B. p=8, n=9, e=8 C. p=8, n=9, e=6 D. p=9, n=8, e=11 E. p=9, n=8, e=9
3. X and Y are two species, each consisting of one
nucleus and a number of electrons. The two species are found to contain the same number of protons, the same number of neutrons, and different numbers of electrons. Which of the following statements about X and Y is correct?
A. X and Y are both neutral atoms. B. X and Y are isotopes of one another. C. The two have different atomic numbers. D. At least one of the species is an ion. E. There is no correct answer above since A⎯D are
all incorrect.
4. Which of the following chemical formulas is NOT the expected one for the compound named?
A. Ga2O3 – gallium oxide B. AlCl3 – aluminum chloride C. Li2O – lithium oxide D. MgBr – magnesium bromide E. SrI2 – strontium iodide
22
Chapter 2 continued 5. Which of the following formula–name combinations
B. 2 only C. 1 and 3 only D. 2 and 3 only E. 1 and 2 only
Chapter 3 1. The molecular weight of ethylene glycol, C2H6O2, is
A. 29 g/mole B. 34 g/mole C. 62 g/mole D. 94 g/mole E. 46 g/mole
2. One gram of alum, KAl(SO4)2 12H2O, has 1.3 × 1021
Al atoms. How many oxygen atoms are present in 1.0 g of alum?
A. 1.3 × 1022 B. 2.6 × 10 22 C. 1.6 × 1022 D. 1.0 × 1022 E. 2.1 × 1022
3. What is the weight percent of silver in silver nitrate,
AgNO3?
A. 63% B. 20% C. 43% D. 78% E. None of the above is correct to within 5%.
4. What is the empirical formula of a hydrocarbon
containing 84.2% C and 15.8% H by weight?
A. C8H18 B. C16H3 C. C3H8 D. C5H12 E. C4H9
5. Which of the following samples contains the largest total number of atoms?
A. 0.1 moles of P4O10 B. 0.2 moles of P4O6 C. 0.3 moles of N2O5 D. 0.4 moles of N2O4 E. 0.5 moles of BiF3
6. “X” is an unknown element which forms an acid,
HXO3. The mass of 0.0133 mol of this acid is 1.123g. Find the atomic mass of X and identify the element represented by X. The element X is
A. N B. Cl C. P D. Br E. I
7. The percent mass of hydrogen in the following
compound is
A. 5.2% B. 7.8% C. 8.7% D. 7.2% E. 8.4%
8. The carbon backbone for a molecule is given below.
What is the molecular weight of this compound?
A . 114 g/mol B . 115 g/mol C . 119 g/mol D . 120 g/mol E . 126 g/mol 9. A Chem 110 TA synthesizes a compound composed
of carbon, hydrogen, and nitrogen and submits 0.1156 g of it to combustion analysis. The TA recovers 0.3556 g of carbon dioxide and 0.0655 g of water. What is the empirical formula of the compound?
A. C11H11 B. C20H18N2 C. C5H5N D. C10H18 E. C10H9N
O
H2C
O
H2C CH
CH3
N CH3
H
H
H
H
23
Chapter 4 1. Which of the following aqueous solutions would you
expect to be the best conductor of electric current?
A. 1.0 M sugar (C6H12O6) B. 1.0 M CaCl2 C. 1.0 M ethanol (C2H5OH) D. 1.0 M acetic acid (HC2H3O2) E. 1.0 M NH4OH
A. 1 and 2 are weak electrolytes B. 1 and 3 are weak electrolytes C. 1, 2 and 3 are weak electrolytes D. 4 is the only weak electrolyte E. 1 is the only weak electrolyte
3. How much 0.154 M NaCl, “physiological saline,” can
be prepared by dilution of 100 mL of a 6.0 M NaCl solution?
A. 1.1 L B. 910 mL C. 90 mL D. 540 mL E. 3.9 L
4. How many milliliters of a 132.00 mL solution of 1.98
M AlCl3 must be used to make 162.00 mL of a solution that has a concentration of 0.630 M Cl–(aq)?
A. 1.61 mL B. 7.73 mL C. 17.2 mL D. 41.5 mL E. 51.5 mL
5. A 24.00 mL sample of a solution of Pb(ClO3)2 was
diluted with water to 52.00 mL. A 17.00 mL sample of the dilute solution was found to contain 0.220 M ClO3
–
(aq). What was the concentration of Pb(ClO3)2 in the original undiluted solution?
A. 3.60 × 10–2 M B. 7.19 × 10–2 M C. 0.238 M D. 0.156 M E. 0.477 M
Chapter 5 1. It is found that 6.00 g of potassium metal reacts with
excess water to release 29.8 kJ of heat. This means that, for the reaction
2 K(s) + 2 H2O(l) → 2 KOH(aq) + H2(g)
the enthalpy of reaction (per mole of H2 produced) is A. ΔH = +189 kJ B. ΔH = −4.97 kJ C. ΔH = +388 kJ D. ΔH = −194 kJ E. None of the above is within 1% of the correct
answer 2. Given the following reactions and their enthalpy
changes
2 C2H2(g) + 5 O2(g) → 4 CO2(g) + 2 H2O(l)
ΔHRXN = –2599.2 kJ
C(s) + O2(g) → CO2(g) ΔHRXN = –393.5 kJ 2 H2(g) + O2(g) → 2 H2O(l) ΔHRXN = –571.8 kJ Calculate ΔH°f for acetylene, C2H2(g). A. 620.2 kJ B. −620.2 kJ C. −453.4 kJ D. 226.7 kJ E. None of the above is within 10% of the correct
answer 3.
The value of ΔH° for the following reaction is –62.1 kJ. What is the value of ΔH°f (in kJ/mol) for HCl(g)?
SO2Cl2(g) + 2 H2O(l) → H2SO4(l) + 2 HCl(g)
A. −184 B. −372 C. −1079 D. 30 E. −92
Substance ΔH° f (kJ/mol) SO2(g) −297 SO3(g) −396
SO2Cl2(g) −364 H2SO4((l) −814 H2O((l) −286
24
Chapter 5 continued 4. For which one of the following equations is ΔH°rxn
equal to ΔH°f for the product? A. Xe(g) + 2F2(g) → XeF4(g) B. CH4(g) + 2Cl2(g) → CH2Cl2(ℓ) + 2 HCl(g) C. N2(g) + O3(g) → N2O3(g) D. 2 CO(g) + O2(g) → 2 CO2(g) E. None of the above have ΔH°rxn equal to ΔH°f
5. The standard heat of formation for H2O(g) is
–241.8 kJ/mol, and for NaOH(s) it is –425.6 kJ/mol. The enthalpy change for the reaction between Na(s) and H2O(g) to produce NaOH(s) and H2(g) is, per mole of H2(g) produced,
A. −367.6 kJ B. −183.8 kJ C. −667.4 kJ D. not determinable from the data given E. None of the above answers is correct
6. The standard heat of formation of ammonia, NH3(g),
is –46.2 kJ/mole. Which of the following is true about the reaction:
2 NH3(g) → 3 H2(g) + N2(g)
A. The reaction is exothermic with ΔH = −46.2 kJ. B. The reaction is endothermic with ΔH = −92.4 kJ. C. The reaction is exothermic with ΔH = 92.4 kJ. D. The reaction is endothermic with ΔH = 92.4 kJ. E. The reaction is endothermic with ΔH = 46.2 kJ.
7. 0.0100 mole of dry, solid KClO3 is added to 50.0 g of
water at 20.10°C in a coffee–cup calorimeter. The temperature is observed to drop to 18.10°C. ΔH of hydration for one mole of KClO3 is (within 2%) [Heat capacity of water = 4.18 J/°C g]
A. –0.836 kJ B. 0.836 kJ C. –42.8 kJ D. 42.8 kJ E. None of the above answers is correct.
Chapter 6 1. A radio station broadcasts on a frequency of 99.5
kilocycles/s. What is the wavelength of this radiation in km?
A. 7.25 × 103 km B. 2.99 × 107 km C. 2.99 × 104 km D. 3.02 km E. none of these
2. What is the energy possessed by one mole of
x–ray photons if the wavelength of the x–ray is 1.00 × 10−
9 m?
A. 1.29 × 10−27 J
B. 1.99 × 10−16 J
C. 1.21 × 10−10 J
D. 1.21 × 108 J E. 3.34 × 1010 J
3. Green light of wavelength 516 nm is absorbed by an
atomic gas. What is the energy difference between the two quantum states involved in the transition?
A. 5.81 × 1014 J B. 3.85 × 10−
19 J C. 1.28 × 10−
27 J D. 4.29 × 10−
36 J E. 1.43 × 10−
44 J 4. Which photon has an energy that is greater than the
energy of a blue photon?
A. microwave photon B. radio photon C. green photon D. infrared photon E. ultraviolet photon
5. A He-Ne laser (λ = 633nm) is used to heat up a
sample. How many photons are needed to transfer 12 J of heat to the sample?
A. 3.8 × 1019 B. 12,000 C. 6.3 × 1011 D. 3.2 × 10−
19 E. 1.6 × 106
25
Chapter 6 continued 6. If the Bohr model is used, what frequency of light
would be required for ionization of hydrogen?
A. 6.17 × 1014 Hz B. 1.31 × 103 Hz C. 3.29 × 1015 Hz D. 4.31 × 1010 Hz E. None of the above is within 5% of the correct
answer 7. Which of the following statements are true for the
Bohr model of the hydrogen atom?
1. The radius of the orbit increases as the principal quantum number increases.
2. The energy required to ionize the atom increases as the principal quantum number decreases.
3. Light emitted by the excited hydrogen atom corresponds to transitions from orbits of higher principal quantum number to lower principal quantum number.
A. 1 only B. 1and 2 only C. 2 and 3 only D. 1 and 3 only E. 1, 2, and 3
8. Which of the following electron transitions in a
hydrogen atom results in the greatest release of energy from the atom?
A. n = 3 to n = 4 B. n = 1 to n = 3 C. n = 6 to n = 4 D. n = 7 to n = 5 E. n = 2 to n = 5
9. For electron distributions, which of the following
statements are true?
1. d orbitals have a spherical shape. 2. p orbitals have a high electron density at the
nucleus. 3. s orbitals have no electron density at the
nucleus. A. 1 and 2 B. 2 only C. 2 and 3 D. 3 only E. None of the statements is true
10. Which series of quantum numbers describes the orbital in which the highest energy electron in potassium resides in the ground state?
n ℓ mℓ A. 4 3 0 B. 3 0 1 C. 4 1 0 D. 4 0 0 E. None of these
11. An atom of phosphorus 15P has how many electrons
with quantum number ℓ = 1?
A. 3 B. 9 C. 15 D. 5 E. None of the above
12. Which of the following could not be an orbital diagram
for an atom in its ground state?
1s 2s 2p 3s
A. (⇅) (⇅) (⇅) (↑) (↑)
B. (⇅) (⇅) (⇅) (⇅) (⇅) (↑)
C. (⇅) (⇅) (⇅) (⇅) ( )
D. (⇅) (⇅) (⇅) (⇅) (⇅) (⇅)
E. (⇅) (⇅) (↑) (↑) (↑) 13. Which of the following ground state electron
configurations can be ruled out by the Pauli Exclusion Principle?
1. 1s3 2s2 2p5
2. 1s2 2s2 2p7 3s2
3. 1s2 2s2 2p6 3s2 3p6 4s2 3d12
4. 1s2 2s2 2p6 3s2 3p6
A. 1 only B. 1 and 3 only C. 1, 2, and 3 only D. 1, 2, 3, and 4 only E. 1 and 4 only
26
Chapter 6 continued 14. In the ground state electronic configuration of Cr, how
many total shells, subshells and orbitals contain at least one electron, and how many unpaired electrons are present?
Shells Subshells Orbitals Unpaired Electrons
A. 4 3 8 4
B. 4 7 15 6
C. 4 7 15 4
D. 4 7 14 4
E. 4 7 14 6 15. Which orbital diagram(s) represents an excited state
electron configuration of C?
i.
ii.
iii.
A. i only B. iii only C. ii and iii D. i and iii E. i, ii and iii
16. When light of wavelength 420 nm is focused on a
metal surface, electrons are ejected with a speed of 7.00 × 105 m/sec. The binding energy of a mole of electrons to the metal is
A. 8.22 × 104 J/mole B. 1.50 × 105 J/mole C. 2.72 × 102 J/mole D. 5.63 × 101 J/mole E. None of the above is correct to within 5%
Chapter 7 1. Elements D, E, and G have atomic numbers Z, Z+1,
and Z+2 respectively. E is a noble gas (not helium). Which statement is true?
A. Elements D and G should form a compound which
is a gas at room temperature and which has formula GD.
B. Elements D and G should form a compound which is a solid at room temperature with formula GD.
C. Elements D and G should not react with each other.
D. Elements D and G should react to form a covalent compound with formula DG7.
E. None of the above is a true statement. 2. Atom X is in Group IIA and atom Y is in Group VIIA.
A compound formed between these two elements would have the formula
A. X2Y B. XY2 C. X2Y7 D. X7Y2 E. None of the above
3. Which of the following has the configuration of a
noble gas?
1. S2+ 2. Br− 3. Si4+
A. 1 only B. 2 only C. 1 and 2 D. 2 and 3 E. 1 and 3
4. The species which may not have the electron
configuration 1s2 2s2 2p6 3s2 3p6 4s1 4d1 is A. Sc2+ B. Sc+ C. Ca D. K1
− E. Ti2+
27
Chapter 7 continued 5. Which of the following is a true statement?
A. The configuration 1s22s22p4 is the electronic configuration of O2
−. B. The ground state for the configuration 1s22s22p3
has one unpaired electron spin. C. Sulfur is the only element in the third period with 2
unpaired electrons in the 3p subshell. D. The configuration [Ar]3d2 is the configuration for
Ti2+. E. None of the above is a true statement.
6. The electron affinity of the chlorine atom is the energy
change which occurs in the reaction
A. Cl(g) → Cl+(g) + e− B. Cl−(g) → Cl(g) + e− C. Cl(g) + e− → Cl+(g) D. Cl(g) + e− → Cl−(g) E. Cl(g) → Cl−(g) + e−
7. The ionization energy for Li2+ is expected to be
greater than that for H because:
A. Lithium is a metal whereas hydrogen is not. B. The nuclear charge in lithium is greater than the
nuclear charge in hydrogen. C. The size of the Li2+ ion is greater than the size of
the H atom. D. The principal quantum number of the electron is
greater in Li2+ than in H. E. The nucleus of the hydrogen atom has no
neutrons. 8. Pick the pair that has the lowest (easiest to ionize)
ionization energy and the most endothermic (most difficult to attach an electron) electron affinity.
Lowest I.E. Most endothermic E.A. A. Cs O B. O Na C. Na F D. Na O E. Cs Mg
9. For the series of Na+, O2−, F–, which statement is
FALSE?
A. The series is isoelectronic. B. Na+ has the greatest ionization energy (for forming
Na2+).
C. A valence electron in O2– experiences the least
effective nuclear charge. D. All three species have a filled quantum level
n = 2. E. F– has the smallest radius.
10. How many unpaired electrons are there in the Fe3+
ion when it is in its ground state?
A. 1 B. 2 C. 3 D. 4 E. 5
11. A correct orbital diagram for the ground state of the
O+ ion is
A. 1s ⇅ 2s ⇅ 2p ⇅ ↑ ↑ B. 1s ⇅ 2s ⇅ 2p ⇅ ⇅ ↑ C. 1s ⇅ 2s ⇅ 2p ⇅ ↑ _ D. 1s ⇅ 2s ⇅ 2p ↑ ↑ ↑ E. 1s ⇅ 2s ⇅ 3s ⇅ 4s ⇅
12. Which of the following gives the correct relationships
among the first ionization energies of the elements?
1. Na > Al > P 2. Rb > K > Li 3. Se > S > O 4. He > Ne > Ar
A. 1 only B. 2 only C. 3 only D. 4 only E. Both 3 and 4 are correct.
28
Chapter 8 1. Which of the following compounds would have the
largest lattice energy?
A. CaO B. CsI C. BaS D. NaF E. NaCl
2. Which of the following ionic crystals would you expect
to have the lowest melting point?
A. KF B. KBr C. CaO D. PbS E. ScN
3. Which of the following compounds would you expect
to possess a multiple bond?
A. SbH3 B. AlCl3 C. CBr4 D. C2H4 E. SiF4
4. A reasonable Lewis structure for the molecule CH2O
is
A. H C O H
B.
HC O
H
C. H C O H
D.
H
C O
H
E. H C O
H
5. What is a reasonable Lewis structure for cyanogen, C2N2?
A. N C NC
B. N N CC
C. C C NN
D.
C N
N C
E. C C NN
6. Nitrogen–nitrogen bond lengths of 1.10, 1.25 and
1.45 Å have been measured for different molecules.
Match the molecules N2F2, N2, N2F4 with appropriate N−N bond length. 1.10 Å 1.25 Å 1.45 Å A. N2 N2F4 N2F2 B N2 N2F2 N2F4 C. N2F4 N2 N2F2 D. N2F2 N2 N2F4 E. N2F4 N2F2 N2
7. The electronegativities of four elements (L, M, Q, and
R) are as follows:
L: 1.1 M: 2.1 Q: 2.4 R: 3.5
In which of the following diatomic molecules would the least polar bond be expected? A. LM B. MQ C. QR D. LR E. MR
8. Which of the following species violates the octet rule?
A. GeF4 B. TeF4 C. BH4− D. SO4
2−
E. SiH4
29
Chapter 8 continued 9. Read the following statements:
1. BF3 and PF5 are examples of violations of the octet rule.
2. BF4− and BF3−NH3 are examples of violations
of the octet rule. 3. Expanded valance shells occur most often
when the central atom is bonded to a small, electronegative element.
4. The central atoms most capable of having
expanded valence shells come from rows 3, 4, and 5 of the periodic table.
A. Only statements 1, 3,and 4 are correct. B. Only statements 1 and 4 are correct. C. Only statements 2, 3, and 4 are correct. D. Only statements 3 and 4 are correct. E. Only statements 2 and 4 are correct.
10. From Lewis structures for SO2 and SO3
2−, we can
predict that
A. SO2 has two different S−O bond lengths, one of which is the same as in SO3
2−.
B. All S−O bonds in SO32
− should be weaker than S−O bonds in SO2.
C. All S−O bonds in SO32
− are the same because there are two Lewis resonance structures for this ion.
D. S−O bonds in SO32
− should be shorter than those
in SO2. E. None of the above predictions follow from the
Lewis structures. 11. Which of the following is a TRUE statement?
A. SeO2 has two equivalent resonance structures. B. XeF4 obeys the octet rule. C. The structure NOBr is more stable than ONBr. D. Because of fluorine’s high electronegativity, the
F−F bond is very polar. E. None of the above statements is true.
12. Given the following bond energies C−C (348 kJ/mole), C=C (614 kJ/mole), H−H (436 kJ/mole) and the heat of reaction for the following process, estimate the C−H bond energy.
C2H4 + H2 → C2H6 ΔH = −137 kJ
A. 420 kJ B. 283 kJ C. 525 kJ D. 350 kJ E. 645 kJ
13. Using the bond energy data tabulated below, estimate
the enthalpy of formation (per mole) of NH3(g).
Bond Average bond energy (kJ/mol) N≡N(N2) 941 H−H 436 N−H 391
The correct answer (in kJ/mol) is closest to A. −422 B. −49 C. 49 D. 204 E. 422
14. Following are some average bond energies (kJ/mol):
C–H 413 O–H 463 C=C 614
C–C 348 O–O 146 C=O 799
C–O 358 O2 495
Estimate the heat released in the complete combustion of ethylene, H2C=CH2. A. 683 kJ/mol B. 802 kJ/mol C. 1297 kJ/mol D. 1403 kJ/mol E. 1792 kJ/mol
30
Organic 1. Which of the following are not acceptable structures
for C4H10? 1) 2) 3) 4) 5)
A. 1 and 2 B. 2 and 3 C. 3 and 4 D. 3 and 5 E. 1, 3 and 5
2. Which of the following is(are) not correct structure(s) of octane C8H18?
1) 2) 3)
A. 1 only B. 3 only C. 1 and 2 D. 2 and 3 E. all of these are possible structures for octane
3. Which of the following Lewis structures are incorrect?
I II III A. I only B. II only C. III only D. I and II E. I, II and III
4. Of the structures shown below, which is a structural
isomer of n-pentane?
a)
b)
c)
d)
e)
CH3 CH
CH3
CH3 CH3 CH2 CH2 CH3
C C C
H
H
H H
H H
H
CH
H
H
CH3 CH
CH3
CH2 CH3
C
CH3
HH
CH3
C H
H
H C C H
H H
H C
H H
C H H
C H H
C H H
C H H
H
H C
H
H
C
H
C
H
C
H
C
H H H
C
H
H
C
H
H
C
H
H
H
C C
C C
C H
H
H
H H
H H H
H
C H H C
H H C H
H
H
CHH
H
C HN
H
C
C C
CC
H
H
HH
H
CH2
CH2
CH2CH2CH2
CH2
CH3 CH2 CCH3
CH3
CH3
CH3 CH2 CH CH3CH3
CH3 CH CH CH2 CH3
CH3 CHCH3
CH2 CH2 CH3
O C C H
H N: H
: :
31
N
O
H CH3
Organic continued
5. Which of the compounds shown below can have a geometrical isomer?
1. 2. 3.
a) 1 only b) 2 only c) 3 only d) 1 and 2 e) 2 and 3
6. Which of the following is/are geometric isomers of the
structure below?
A. 1 B. 2 C. 3 D. 1 and 2 E. 1 and 3
7. Which one of these molecules is an alcohol? A. B. C. D.
E. none of these are alcohols
8. Identify the functional groups present in the following
structure.
1. ester 2. ether 3. amine 4. ketone 5. amide
A. 1 and 5 B. 2 and 5 C. 2 and 3 D. 4 and 5 E. 3 and 4
9. Threonine is a naturally occuring amino acid found in
many proteins. What functional groups are present in threonine?
a) alcohol, amine, ketone
b) alcohol, amine, ester
c) carboxylic acid, alcohol, amide
d) alcohol, carboxylic acid, amine
e) amide, amine, alcohol
H C C C C OHH OHH O
H H NH2
CH3 CH2C C
H
HCH3
CH3 CH2C C
CH3
HCH2CH3
CH3C C
H
CH3H
Br
H3C
C2H5
Cl
Br
H3C
Cl
C2H5
H3C
Br
Cl
C2H5
Br
Cl
1 2 3
CH3CCH2CH3O
CH2CH2CH3OH
CH3HCCH2CH3
O
32
Chapter 9
1. Which of the following are correct statements about
the valence–shell electron–pair repulsion (VSEPR) model of bonding?
1. Electron pairs orient themselves to give the
smallest angles possible. 2. Only bonding electron pairs are important in
the VSEPR model.
3. Electron–pair geometry in all cases
describes the spatial geometry of the atoms in the molecule.
A. 1 only B. 2 only C. 3 only D. 1, 2, and 3 E. None of the above
2. Which of the following is non–planar?
A. SO2 B. SO3 C. SO4
2−
D. NO3− E. BF3
3. The number of bonding pairs of electrons, non–
bonding pairs of electrons and molecular shape of the H3O+ ion are
Bonding
Pairs Non–bonding
Pairs
Molecular Shape
A. 4 0 tetrahedral
B. 3 1 tetrahedral
C. 2 2 bent
D. 3 1 trigonal pyramidal
E. 3 0 trigonal planar
4. Among the following gaseous molecules
BeF2 BF3 CF4 NF3 OF2 What is the correct trend in FMF bond angle?
LARGEST SMALLEST
A. BeF2 > BF3 > CF4 > NF3 > OF2
B. BeF2 = OF2 > CF4 > BF3 = NF3
C. BeF2 = OF2 > BF3 = NF3 > CF4
D. BeF2 > BF3 > OF2 > NF3 > CF4
E. CF4 > BF3 = NF3 > BeF2 = OF2 5. Which of the following molecules has a trigonal planar
structure?
1. CO32
− 2. SOCl2 3. H3O+ 4. SO3
2−
A. 1 only B. 1, 2 and 3 C. 1 and 3 D. 1, 3 and 4 E. 3 and 4
6. Which of the following molecules and ions possesses
a tetrahedral molecular structure?
A. TeI4 B. SeBr4 C. XeCl4 D. NH3 E. AlF4−
7. The bond angles in H2O2 are approximately:
A. 90° B. 105° C. 109.5° D. 120° E. 180°
33
Chapter 9 continued 8. Identify which of the following bonds is(are) polar and
which element in each is the more electronegative.
HCl NO Si2 Now consider these five statements about your conclusions.
1. HCl is polar, with H the more electronegative. 2. HCl is polar, with the Cl the more
electronegative. 3. NO is polar, with the N the more
electronegative. 4. NO is polar, with O the more electronegative. 5. Si2 is polar.
The correct statement(s) above is(are) A. 2 only B. 2 and 3 only C. 2 and 4 only D. 2, 3 and 5 only E. 1 and 4 only
9. Consider the following molecules.
1. BF3 2. NH3 3. SOCl2 4. SiF4 Which of these should have a dipole moment? A. 2 only B. 2, 3 and 4 only C. 1, 2 and 3 ony D. 1 and 4 only E. 2 and 3 only
10. Which of the following molecules has the largest
dipole moment?
A. XeF2 B. XeF4 C. PF5 D. SF4 E. SF6
11. Which of the following benzene–like molecules do you expect to have the largest dipole moment?
A.
F
F
F
FCl
Cl
B.
F
H
F
HH
F
C.
Cl
Cl
H
HH
H
D.
H
H
H
FH
F
E.
H
F
H
FH
H
12. Which of the following statements is true? A. There are two π bonds in B. The H–N–H angle in NH3 is slightly larger than the H–C–H angle in CH4. C. In CO2, the carbon is sp2 hybridized. D. The molecule below possesses a dipole moment. E. None of the above statements is true.
C C H H
H H
C C H C l
C l H
34
BNB
NBN
H
HH
HH
H
Chapter 9 continued 13. Which of the following is a correct set of electron–pair
geometry, hybridization, and bond angle(s)?
Geometry Hybridization Bond Angle(s)
A. linear sp2 180°
B. trigonal planar sp2 120°
C. octahedral d2sp3 60°
D. tetrahedral dsp3 109.5°
E. trigonal bipyramidal d2sp3 120°, 90°
14. Of the following statements about delocalized bonds,
which are true?
1. They occur in molecules with alternating single and double bonds.
2. They can be made from sp2 hybrid σ bonds. 3. They are most often found in bonds involving
heavy atoms. 4. They can be made from π bonds. 5. They rarely, if ever, involve bonds with
fluorine.
A. 1, 4, and 5 only B. 4 and 5 only C. 1, 3, and 4 only D. 1, 2, and 3 only E. All of the above
15. Which of the following statements is incorrect?
A. Hybridization accounts for the experimental observation that all F−C−F bond angles in CF4 are the same.
B. All C−C bond distances in benzene, C6H6, are the same.
C. The energy required to break a carbon–carbon
triple bond is greater than that needed to break a
carbon–carbon double bond.
D. The bonding in benzene, C6H6, includes 12 sigma (σ) bonds and 3 localized pi (π) bonds.
E. The π electrons in NO3− are delocalized.
16. The following molecule should be used
H
C1
H
C2
H
CC3
H
The hybridizations of C1, C2 and C3 respectively are
C1 C2 C3 A. sp3 sp3 sp2 B. sp3 sp sp2 C. sp3 sp3 sp3 D. dsp3 sp sp2 E. None of the above.
17. Borazine B3N3H6 has been called inorganic benzene.
What is the hybridization of the boron atoms?
A. sp
B. sp2 C. sp3 D. sp4 E. dsp3
35
Chapter 10 1. The volume of a large, irregularly shaped, closed tank
is determined as follows. The tank is first evacuated, and then it is connected to a 50.0 L cylinder of compressed nitrogen gas. The gas pressure in the cylinder, originally at 21.5 atm, falls to 1.55 atm without a change in temperature. What is the volume of the tank?
A. 100 L B. 644 L C. 25 L D. 744 L E. There is not sufficient information to determine the volume of the tank.
2. Which curve represents the relationship between the
volume of an ideal gas and its temperature in K at constant pressure?
(E) None of the above. 3. Which statement about ideal gases is incorrect?
A. Equal numbers of moles of two ideal gases of different molecular weight will fill the same volume at the same temperature and pressure.
B. If the temperature of a given amount of an ideal gas is held constant, its pressure is directly proportional to the volume it occupies.
C. The density of an ideal gas is independent of the volume it occupies at a given temperature and pressure.
D. The mass of an ideal gas is proportional to its pressure at a fixed volume and temperature.
E. The pressure of an ideal gas is proportional to its absolute temperature if its density is held constant.
4. Solid substances A and B react to form a solid product C and a gaseous product D. Solids A, B, and C are weighed, and the volume, temperature and pressure of gas D are measured. From this and the ideal gas law, the number of moles of D are calculated. The data are:
mass of A reacted: 5.00g mass of B reacted: 20.00g mass of C produced: 15.00g moles of D produced: 0.208 moles
The molecular weight of D is A. 48.1 B. 2.08 C. 72.1 D. 208 E. Impossible to determine from the data given.
5. A gaseous compound having the empirical formula
CH2 has a density of 1.88 g/L at STP. The molecular formula for this compound is (assuming ideal behavior)
A. CH2 B. C2H4 C. C3H6 D. C4H8 E. Not determinable from the data given.
6. A gas which has a density of 2.77 g/L at STP is found
to contain the following weight percents: 51.6% oxygen, 9.7% hydrogen, and 38.7% carbon. [Atomic weights are: C = 12.00 amu; H = 1.00 amu; O = 16.00 amu] The molecular formula for this gas is
A. CH3O B. C3H10O C. C3H9O3 D. C2H6O2 E. C2H4O
36
Chapter 10 continued 7. Consider the following gases: H2S, H2, and SO2.
Which of the following statements is accurate at any given temperature?
A. The average kinetic energy of H2S is the same as
that for H2. B. The rms speed of SO2 will be the same as that of
H2S. C. The rms speed of SO2 will be greater than that of
H2S. D. The average kinetic energy of SO2 will be greater
than that of either H2S or H2. E. All molecules of H2S will have the same kinetic
energy. 8. An ideal gas at 200.0 K and 1.50 atm pressure is
heated to 1000.0 K and kept at 1.50 atm pressure. Which of the statements 1–4 below are true?
1. The root mean square speed of the
molecules increases by a factor of five. 2. The average molecular kinetic energy
increases by a factor of twenty–five.
3. The collision frequency with a unit area of container wall increases.
4. The average molecular kinetic energy increases by a factor of five.
A. 1 and 2 B. 2 and 3 C. 3 and 4 D. 3 only E. 4 only
9. Ideal gases A and B are kept in a 5.0–liter container
at 20°C. The pressure of the gas mixture is 2.0 atm and it is known that 0.10 mole of gas A is present in the mixture. How many moles of B are present in the mixture? (MW of A = 25 amu; MW of B = 100 amu)
A. Insufficient information given to calculate moles of
B present. B. 0.32 moles C. 0.25 moles D. 0.42 moles E. 0.48 moles
10. 3.0 L of He gas at 5.6 atm pressure and 25°C and 4.5 L of Ne gas at 3.6 atm and 25°C are combined at constant temperature into a 9.0 L flask. What is the total pressure (in atm) in the 9.0 L flask?
A. 2.6 B. 9.2 C. 1.0 D. 3.7 E. 4.6
11. A student collects the following data in the process of
collecting nitrogen gas by water displacement:
Volume of gas collected: 73.2 mL Temperature of room air and of water: 22.5°C Barometric pressure: 752 mm Hg Vapor pressure of water at 22.5°C (interpolated):
20.45 mm Hg Assuming ideal gas behavior, which of the following is the number of moles of N2 the student collected? A. 0.00290 moles N2 B. 0.00299 moles N2 C. 2.21 moles N2 D. 2.27 moles N2 E. None of these is within 1% of the correct answer.
12. Which of the following statements is false?
A. Dalton’s law of partial pressures indicates that in gaseous mixtures at low pressure each kind of molecule behaves independently of the others.
B. At high pressures different gases give different values for the ratio PV/nRT.
C. The mean free path of a molecule depends on its size.
D. Collisions of molecules with the container walls give rise to the gas pressure.
E. At a given temperature, for a given gas, every molecule has the same speed.
13. If the temperature of an ideal gas is held constant
while its volume doubles, which one of the following happens?
A. The average velocity of the molecule increases. B. The frequency of collisions of gas molecules with
a given area of wall decreases. C. The mean free path of the molecules remains
unchanged. D. The density of the gas remains unchanged. E. None of the above is a true statement.
37
Chapter 10 continued 14. The properties of a real gas are most likely to deviate
from those properties predicted for an ideal gas when
A. the pressure is low. B. the temperature is high. C. the pressure is high and the temperature is low. D. the pressure is low and the temperature is high. E. the density of the gas is low.
15. The fact that the ideal gas law only approximately
describes the behavior of a gas might be partly explained by one or more of the following:
1. R is not really a constant. 2. Gas molecules really do not have zero
volume. 3. The kinetic energy of gas molecules is not
really directly proportional to the absolute temperature.
4. Gas molecules really do interact with each other.
Which are correct? A. 2 only B. 3 only C. 4 only D. 1 and 2 E. 2 and 4
Chapter 11 1. The boiling point of ammonia, NH3, is significantly
higher than that of analogous compounds from the same group, e.g., PH3, AsH3, and SbH3. The principal reason for the abnormally high boiling point is
A. London dispersion forces B. ion⎯dipole forces C. dipole⎯dipole interactions D. hydrogen bonding E. ion⎯induced dipole forces
2. Which of the following statements about intermolecular forces is true?
A. The large London dispersion forces between
helium atoms make helium difficult to liquefy. B. Intermolecular forces in a liquid are weaker than
intramolecular forces. C. Ion⎯ion forces in an ionic solution are
independent of the distances between ions. D. Dipole⎯dipole forces dominate in interactions
between carbon dioxide molecules. E. Van der Waals forces account for the cohesive
energy of an ionic liquid. 3. Elementary phosphorus is a solid consisting of P4
molecules and melts at 44°C. The principal forces between the molecules in the solid must be
A. London dispersion. B. ionic. C. covalent. D. dipolar. E. metallic.
4. The boiling points of the substances MgO, Ne, and
H2O should increase in the order
A. H2O < MgO < Ne B. Ne < MgO < H2O C. MgO < Ne < H2O D. MgO < H2O < Ne E. Ne < H2O < MgO
5. The viscosity of a liquid is an important factor in
A. its rate of evaporation. B. capillary rise. C. bubble formation. D. rate of flow through a tube. E. its vapor pressure.
6. Which of the following properties is not increased
when the strength of intermolecular forces in a liquid is increased?
A. vapor pressure B. heat of vaporization C. viscosity D. surface tension E. normal boiling point
38
Chapter 11 continued 7. Consider the three compounds sketched below:
The viscosities of these three compounds should increase in the order A. 1 < 2 < 3 B. 2 < 3 < 1 C. 3 < 2 < 1 D. 3 < 1 < 2 E. 1 < 3 < 2
8. A pure substance has a triple point in its phase
diagram at 216 K and 5.1 atm. This means that:
A. at a pressure of one atmosphere the substance can coexist as a solid and vapor.
B. at room temperature and one atmosphere pressure the substance is a liquid.
C. above the triple point pressure and temperature the substance can exist only as a vapor.
D. at a constant temperature of 216 K the substance is a solid at any pressure.
E. at a constant pressure of 5.1 atm the substance is a liquid at any temperature.
9. Solid CO2 sinks in liquid CO2, and the triple point of
CO2 is at –56.4°C and 5.11 atm pressure. What state of CO2 exists at –56.4°C and 6.00 atm pressure?
A. Gas B. Liquid C. Solid D. Supercritical fluid E. Mixture of solid and liquid
.
10.
The vapor pressure curves for two substances are shown in the figure. Identify which of X or Y has the higher normal boiling point and the larger intermolecular forces. Higher Boiling
Point Larger Intermolecular Forces
A. Y Y B. X X C. Y X D. X Y
E. Insufficient data given to determine the answer.
H
H H
H
H
H H H H
H C C C C F F O
H
H H
H
H
H C C C O C C O C
H H
H
H
H H
H
H 1 2 3
39
Chapter 11 continued 11. Which of the following compounds has the lowest
vapor pressure at 0 °C?
OH
HO OH
C
CH3
H3C
CH3
CH3
C
OH
HOH2C
H
CH2OH
A
B
C
D
E
12. Which of the following would you expect to have
the highest boiling point? A.
B. C.
D. E.
CH3 CH
CH3
CH2 CH3
CH3 CH
CH3
C
HCH3H
H C C C C C C H
H H H H H H
H H H H H
C
C
H
HH
HH
H C C C C C C H
H H H H H H
H H H H C H
H H H
C C
C C
C C
H H
H
H H
H H
H H
H H
H H
H
40
H H H
i) H—C—C—C—H
H H H
H Cl ii) H—C—C—Br H H Cl H H iii) H—C—C—C—Br
H H H
Chapter 11 continued 13. Rank these in order of decreasing boiling point
A. iii > ii > i B. i > ii > iii C. ii> i > iii D. iii > i > ii E. i > iii > ii
Chapter 13 1. The mole fraction of CO2 in a certain solution with
H2O as the solvent is 3.6 × 10−4. The molality of CO2
in this solution is about
A. 0.00036 m B. 0.0065 m C. 0.020 m D. 2.0 × 10−
5 m E. 6.5 m
2. A solution contains 4.00 g NaOH, 5.61 g KOH, and
1.03 g RbOH in 90.0 g of water. The solution has a density of 1.08 g/mL. What is the molality of OH−?
A. 2.33 m B. 2.09 m C. 2.52 m D. 2.16 m E. None of the above is within 2% of the correct
answer. 3. The mole fraction of HCl in a 36% by weight aqueous
solution of HCl is
A. 0.11 B. 0.22 C. 0.36 D. 0.64 E. 0.99
4. A solution whose density is 0.935 g/mL contains
30.0% by weight H2CO, 10.0% C2H5OH and 60.0% H2O. The molarity of the H2CO is
A. 9.34 M B. 2.32 M C. 0.0107 M D. 4.51 M E. 0.0214 M
5. 27.0 L of HCl gas at STP is dissolved in water, giving
785 mL of solution. The molarity of the HCl solution is:
A. 9.46 M B. 1.53 M C. 0.946 M D. 15.3 M E. None of the above is within 1% of the correct
answer.
41
Chapter 13 continued 6. Which one of the following statements is true?
A. The heat of solution in H2O for KOH is the same as that for KOH H2O.
B. For a compound to spontaneously dissolve in H2O at 25°C, the solution process must be exothermic.
C. Ne is less soluble in H2O at 2 atm pressure than at 1 atm pressure.
D. C2H4 is very soluble in liquid NH3. E. None of the above.
7. Given the solutes NaCl and I2 and the solvents NH3(l) and C7H16 (heptane), which is true?
A. NaCl and I2 are both more soluble in NH3 than in
C7H16. B. NaCl and I2 are both more soluble in C7H16 than in
NH3. C. NaCl is more soluble in NH3; I2 is more soluble in
C7H16. D. NaCl is more soluble in C7H16; I2 is more soluble in
C7H16. E. In order to determine relative solubilities in these
cases one must know whether the solution reactions are exothermic or endothermic.
8. The Henry’s law constant for He gas in H2O at 30°C is 3.7 × 10−
4 M/atm; that for N2 at 30°C is 6.0 × 10−
4 M/atm. If a gaseous He-N2 mixture that has a He mole fraction of 0.30 is placed over the water
A. the concentration of dissolved He will be greater than that of dissolved N2.
B. the concentration ratio of dissolved He to dissolved N2 would be 0.62.
C. the concentration ratio of dissolved He to dissolved N2 will be 0.26.
D. the concentration ratio of dissolved He to dissolved N2 will be 0.18.
E. the concentration ratio of dissolved He to dissolved N2 will be 0.49.
9. A saturated solution of silver chloride, AgCl is weakly conducting for electricity because
A. AgCl is essentially 100% ionized in solution, but is
not very soluble. B. AgCl is quite soluble, but only dissociates into ions
slightly, in solution. C. AgCl does not produce ions itself, but rather
induces hydrolysis to give AgOH and HCl, in water.
D. AgCl is only slightly ionized in solution, and also is only slightly soluble.
E. None of the above is a true statement.
10. Which one of the following would you expect to me MOST soluble in water?
A. CH3OH B. CH3CH2CH2OH C. CH3CH2 CH2 CH2CH2OH D. CH3CH2 CH2 CH2CH2CH3
E.
11. Which one of the following would you expect to be
LEAST soluble in water
A. CH3OH B. CH3CH2CH2OH C. CH3CH2 CH2 CH2CH2OH D. CH3CH2 CH2 CH2CH2CH3
E.
12. Rank these in order of decreasing solubility in water.
i) ii) iii) CH3CH2OH
A. i > ii> iii B. iii > i = ii C. ii > iii > i D. iii > ii > i E. i = ii > iii
13. Arrange the following aqueous solutions in order of
increasing boiling points:
1. 0.1 M Fe2(SO4)3 2. 0.2 M BaCl2 3. 0.3 M glucose (blood sugar) 4. 0.2 M LiCl 5. 0.2 M AlCl3
A. 3, 4, 2, 1, 5 B. 2, 4, 5, 1, 3 C. 5, 3, 1, 4, 2 D. 3, 4, 1, 2, 5 E. 3, 5, 2, 1, 4
CH2OH
CH2OH
CH3
CH3HO
42
Chapter 13 continued 14. Consider a solvent whose molecular weight is 100.0
amu and whose vapor pressure at 25°C is 25.00 mm Hg. 100.0 g of a nonvolatile solute whose molecular weight is 50.00 amu is dissolved in 1.000 kg of that solvent. The solute is not an electrolyte. What is the vapor pressure of the solution at 25°C?
A. 0.01 mm Hg B. 20.8 mm Hg C. 22.7 mm Hg D. 24.9 mm Hg E. 25.0 mm Hg
15. The compound ethylene glycol, C2H6O2, is widely
used as an automotive antifreeze. What is the minimum weight, in grams, of this compound that must be added to 8 kg of water to produce a solution that will protect an automobile cooling system from freezing at −20°F (−28.9°C)? Ethylene glycol is a non–electrolyte.
A. 32.0 B. 121 C. 7720 D. 965 E. 5342
16. 7.50 g of a compound having the empirical formula
CH2O is dissolved in water to form 100.0 mL of solution. This solution does not conduct electric current and is found to have an osmotic pressure of 12.2 atm. at 25°C. The molecular formula for the compound is
A. CH2O B. C3H6O6 C. C5H10O5 D. C6H12O6 E. None of the above is the correct answer
17. The addition of nonvolatile solute to a solvent
A. increases the boiling point, decreases the freezing point and increases the vapor pressure of the solvent.
B. decreases the boiling point, decreases the freezing point and decreases the vapor pressure of the solvent.
C. decreases the boiling point, increases the freezing point and increases the vapor pressure of the solvent.
D. increases the boiling point, decreases the freezing point and decreases the vapor pressure of the solvent.
E. increases the boiling point, decreases the freezing point but leaves the vapor pressure of the solvent unaffected.
18. Arrange the following aqueous solutions in order of increasing vapor pressure:
1. 0.5 M NaCl 2. 0.4 M Na2SO4 3. 0.1 M KCl 4. 0.1 M Al2(SO4)3 5. 0.1 M C6H12O6 (sugar)
A. 1, 2, 3, 4, 5 B. 5, 4, 3, 2, 1 C. 2, 1, 4, 3, 5 D. 4, 2, 3, 1, 5 E. 5, 3, 4, 1, 2
Chapter 15 1. Kc for the reaction HCl(g) + NH3(g) → NH4Cl(s) is
A. [ ][ ][ ]3
4NHHClClNH
B. [ ][ ][ ]ClNH
NHHCl
4
3
C. [ ][ ]3NHHCl1
D. [HCl] [NH3]
E. [NH4Cl] 2. Consider the following three reactions:
For which of these would Kp = Kc? A. 1 only B. 1 and 2 only C. 1 and 3 only D. 3 only E. 2 and 3 only
3. At 50°C, Kc = 2.2 × 103 for the reaction
3 Fe(s) + 4 H2O(g) ⇄ Fe3O4(s) + 4 H2(g)
What is the value of Kp at 200°C for this reaction? A. 8.8 × 103 B. 2.2 × 103 C. 5.5 × 102 D. 3.5 × 104 E. This question cannot be answered with the
information provided
43
Chapter 15 continued 4. For the following reaction, Kp = 1.96 at 700 K.
NOCl(g) ⇄ NO(g) + ½ Cl2(g) At this same temperature, Kp for the reaction
Cl2(g) + 2 NO(g) ⇄ 2 NOCl (g) is A. 1.96 B. 3.85 C. 0.260 D. 0.509 E. None of the above is within 5% of the correct
answer. 5. Kc for the reaction F2(g) + Cl2(g) ⇄ 2 FCl(g) equals
125 at a particular temperature. Suppose a system involving this reaction is already at equilibrium and the concentrations of F2 and Cl2 are found to be [F2] = 0.115 M and [Cl2] = 0.221 M. What is the concentration of FCl in the system?
A. [FCl] = 2.54 × 10−
2 M B. [FCl] = 3.17 M C. [FCl] = 1.78 M D. [FCl] = 1.43 × 10−
2 M E. None of the above is within 5% of the correct
answer 6. At a certain temperature, 0.300 moles of NO, 0.200
moles of Cl2, and 0.500 moles of ClNO were placed in a 25.0 L vessel and allowed to reach equilibrium:
2 NO(g) + Cl2(g) ⇄ 2 ClNO(g).
At equilibrium, 0.600 moles of ClNO were present. The number of moles of Cl2 present at equilibrium is A. 0.100 B. 0.150 C. 0.200 D. 0.250 E. 0.300
7. Consider the reaction 2 NH3(g) ⇄ N2(g) + 3 H2(g). Suppose 6 moles of pure NH3 are placed in a 1.0–liter flask and allowed to reach equilibrium. If X represents the concentration in moles per liter of N2 present in the system once equilibrium is reached, which one of the following will represent the concentration of NH3 at equilibrium in moles per liter?
A. 6 – (X/2) B. 6 – X C. 6 – 2X D. 3 – 2X E. None of the above is correct
8. For the reaction N2(g) + O2(g) ⇄ 2 NO(g) Kc = 4.0 at a
particular temperature. Suppose we begin an experiment by mixing 1.0 mol of N2 and 1.0 mol of O2 in a 1.0–liter container. What will be the concentration of NO once equilibrium is reached at the given temperature?
A. [NO] = 0.25 M B. [NO] = 0.50 M C. [NO] = 1.0 M D. [NO] = 0.67 M E. None of the above is within 5% of the correct
answer 9. At 750°C, Kp = 0.770 for the reaction
H2(g) + CO2(g) ⇄ H2O(g) + CO(g)
If 0.200 atm of H2 and 0.200 atm of CO2 are admitted into a rigid container and allowed to reach equilibrium, what should the equilibrium partial pressure of CO be? A. 0.0935 atm B. 0.109 atm C. 1.43 atm D. 0.100 atm E. 0.770 atm
44
Chapter 15 continued 10. At high temperatures one mole of hydrogen gas
reacts with one mole of bromine gas to form hydrogen bromide. At a given temperature the equilibrium constant is 57.6. If at the same temperature, a mixture of 4.67 × 10−
3 M bromine gas, 2.14 × 10−3
hydrogen gas, and 2.40 × 10−2 M hydrogen bromide
gas is made, then
A. the system is at equilibrium. B. the system is far from equilibrium and will shift to
form more hydrogen gas. C. the system is far from equilibrium and will shift to
form more hydrogen bromide gas. D. nothing can be deduced since we do not know
whether the reaction is endothermic or exothermic. E. nothing can be deduced since we do not know
whether the equilibrium constant is Kc or Kp. 11. Which one of the following equilibriums is least
affected by a change in the volume of the system?
A. 2 C(s) + O2(g) ⇄ 2 CO(g) B. 2 NO2(g) ⇄ N2O4(g) C. H2(g) + S(ℓ) ⇄ H2S(g) D. H2O(ℓ) ⇄ H2O(g) E. 2 NO(g) + Cl2(g) ⇄ 2 NOCl(g)
12. Nickel (II) oxide can be reduced to nickel metal by
treatment with carbon monoxide as indicated in the reaction
CO(g) + NiO(s) ⇄ CO2(g) + Ni(s) Kp=20 at 500°C
If the reaction chamber contains some solid Ni and NiO, 400 mm Hg of CO2 and 20 mm Hg of CO, all at equilibrium, which one of the following changes will lead to the reduction of more nickel oxide at 500°C? A. Doubling the amount of NiO(s) present. B. Adding CO2 to raise its pressure to 700 mm Hg. C. Adding CO to raise its pressure to 40 mm Hg. D. Removal of half of the NiO(s) present. E. Doubling the volume of the reaction chamber at
500°C.
13. For the reaction FeO(s) + CO(g) ⇄ Fe(s) + CO2(g),
A. the usual expression for the equilibrium constant is Kc =
[CO] [FeO]][CO [Fe] 2
B. addition of CO2(g) will increase Kc. C. increasing the volume of an equilibrium mixture at
constant temperature will cause the number of
moles of CO2 to increase as the mixture re–equilibrates.
D. adding more FeO to an equilibrium mixture will cause the number of moles of CO2 to increase as
the system re–equilibrates.
E. None of the above statements is true Chapter 18 1. Select the substance that is thought to be partially
responsible for depleting the concentration of ozone in the upper atmosphere.
A. CFCl3 B. CO2 C. O2 D. N2 E. NO2
2. The C−Cl bond dissociation energy in CF3Cl is 339
kJ/mol. What is the maximum wavelength of photons that can rupture this bond?
A. 275 nm B. 45.0 nm C. 742 nm D. 353 nm E. 137 nm
3. Which species is primarily responsible for absorbing
ultraviolet radiation in the stratosphere?
A. N2 B. O3 C. CO D. CH2Cl2 E. CO2
45
REACTIONS 1. When the following equation
C9H15S + O2 → CO2 + H2O + SO2
is balanced with the smallest possible set of integer coefficients (no fractions), the coefficient preceding CO2 is
A. 9 B. 27 C. 36 D. 18 E. None of the above is correct
2. When the equation for the complete combustion of
ethanol, C2H5OH, to form CO2 and H2O is balanced with lowest–integer coefficients, the coefficient for O2 is
A. 7 B. 6 C. 5 D. 3 E. None of the above answers is correct
3. A 100.0 g sample of a metal M is burned in air to
produce 103.7 g of its oxide with the formula M2O. What is the atomic mass of the metal?
A. 73 amu B. 150 amu C. 216 amu D. 290 amu E. 370 amu
4. 4.31 g of a hydrocarbon is completely combusted in
oxygen and found to yield 6.36 g of water. The empirical formula for the hydrocarbon is
A. C3H7 B. CH4 C. C7H4 D. C21H4 E. CH
5. 1.00 g of a compound is combusted in oxygen and found to give 3.14 g of CO2 and 1.29 g of H2O. From these data we can tell that
A. the compound contains C, H, and some other
element of unknown identity, so we can’t calculate the empirical formula.
B. the compound contains only C and H and has the empirical formula of C6H.
C. the compound contains C, H, and O and has the empirical formula of CH3O.
D. the compound contains only C and H and has the empirical formula of CH2.
E. None of the above is a true statement 6. Mixing solutions of K2SO4(aq) and BaCl2(aq)
produces an insoluble salt. Which of the following is the correct list of spectator ions for this reaction?
A. K+ , SO4
2– , Ba2+ , Cl– B. K+ , SO4
2– C. K+ , Cl – D. Ba2+ , Cl– E. Ba2+ , SO4
2– 7. When aqueous solutions of Fe2(SO4)3 (aq) and
Pb(NO3)2(aq) are mixed, an insoluble salt forms. Which of the following reactions is the correct net ionic equation that describes this reaction?
A. 2Fe+3(aq)+ 3SO4
2− (aq) + 3Pb2+(aq)+ 6NO3−(aq) →
2Fe(NO3)3(s) + 3Pb2+(aq) + 3SO42
−(aq)
B. Fe+3(aq)+ SO42
−(aq) + Pb2+(aq) + 3NO3−(aq) →
Fe(NO3)3(s) + PbSO4(s)
C. Fe+3(aq) + 3NO3−(aq) → Fe(NO3)3(s)
D. Pb2+(aq)+ 2NO3−(aq) → Pb(NO3)2(s)
E. Pb2+(aq)+ SO42
−(aq)→ PbSO4(s) 8. Which of these reactions will form a precipitate in
water?
A. Mg(OH )2 + 2 HCl → Mg(Cl)2 + 2H2O
B. Mg(OH)2 + (NH4)2SO4 → MgSO4 + 2 NH3 + 2 H2O
C. CaCO3 + 2 HCl → CaCl2 + CO2 + H2O
D. HCl + CH3COOK → CH3COOH + KCl
E. Ba(NO3)2 + K2SO4 → BaSO4 + 2 KNO3
46
REACTIONS continued 9. Which one of the following (unbalanced) equations
describes a spontaneous reaction?
A. Ag(s) + HCl(aq) → AgCl(s) + H2(g)
B. Hg(l) + Al3+(aq) → Hg2+(aq) + Al(s)
C. Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
D. Ni(s) + H2O(l) → Ni(OH)2(s) + H2(g)
E. None of the above describes a spontaneous reaction.
10. Given that the three standard reactions shown below
proceed as written, which of the subsequent 3 reactions do not occur.
Standard reactions
Mg(s) + ZnCl2(aq) → Zn(s) + MgCl2(aq)
Zn(s) + FeCl2(aq) → Fe(s) + ZnCl2(aq)
Fe(s) + CuCl2(aq) → Cu(s) + FeCl2(aq)
I. Mg(s) + FeCl2(aq) → Fe(s) + MgCl2(aq)
II. Zn(s) + CuCl2(aq) → Cu(s) + ZnCl2(aq)
III. Fe(s) + ZnCl2(aq) → Zn(s) + FeCl2(aq)
A. I only
B. II only
C. III only
D. I and II
E. I and III 11. Which of the following reactions are redox reactions?
1. Ni(OH)2(s) + H2SO4(aq) → NiSO4(aq) + 2 H2O()
2. 8 NH3(g) + 6 NO2(g) → 7 N2(g) + 12 H2O()
3. 3 PbO(s) + 2 NH3(aq) → N2(g) + 3 H2O() + 3 Pb(s) A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 2 and 3
12. The oxidation number of Mo (molybdenum) in Mo7O24
6− is:
A. +6 B. +5 C. +4 D. +3 E. +2
13. 100.0 mL of Ca(OH)2 solution is titrated with
5.00 × 10−2 M HBr. It requires 36.5 mL of the acid
solution for neutralization. The concentration of the Ca(OH)2 solution is
A. 9.12 × 10−
3 M B. 1.82 × 10−
4 M C. 9.12 × 10−
4 M D. 1.82 × 10−
3 M E. None of the above is within 1% of the correct
answer 14. Given the balanced equation for the oxidation of
ethanol, C2H5OH, by potassium dichromate:
3C2H5OH + 2K2Cr2O7 + 16HCl →
3C2H4O2 + 4CrCl3 + 4KCl + 11H2O
Calculate the volume of a 0.600 M K2Cr2O7 solution needed to generate 0.1665 moles of C2H4O2 from a solution containing excess ethanol and HCl. A. 278 mL B. 185 mL C. 111 mL D. 6.60 mL E. None of the above answers is correct
15. How many grams of solid NaOH will be needed to
neutralize 85.0 mL of a 0.275 M solution of H2SO4?
A. 0.94 g B. 19.1 g C. 1.87 g D. 0.47 g E. 6.41 g
16. When 47.3 mL of 0.107 M HCl is added to 54.7 mL of
0.213 M NaOH, the resulting OH− concentration is:
A. 0.028 M B. 0.065 M C. 0.114 M D. 0.120 M E. 0.213 M
47
REACTIONS continued 17. A method of removing CO2(g) from a spacecraft is to
allow it to react with solid NaOH:
2 NaOH(s) + CO2(g) → Na2CO3(s) + H2O(ℓ) How many liters of CO2(g) at 26.0°C and 755 mm Hg can be removed per kg NaOH? A. 24.7 B. 16.3 C. 12400 D. 618 E. 309
18. A gaseous sample of pure methane, CH4, is found to
occupy a volume of 20.0 L at 25°C and 1 atm pressure. What volume of O2, at the same temperature and pressure, should be just sufficient to permit complete combustion of the methane?
A. 10.0 L B. 20.0 L C. 30.0 L D. 40.0 L E. None of the above is correct.
19. The following reactions involve only gases. Which of
them, if carried out in a closed container at constant temperature, would proceed with a decrease in pressure?
A. 2 CO + O2 → 2 CO2 B. 2 NH3 → N2 + 3 H2 C. N2 + O2 → 2 NO D. 2 O3 → 3 O2 E. 2 HI → H2 + I2
20. Which of the following compounds would react most
rapidly with bromine (Br2) at room temperature?
a) pentane CH3 - CH2 - CH2 - CH2 - CH3 b) toluene c) ethanol CH3 - CH2 - OH
21. Choose the appropriate words to fill in the blanks. A
condensation reaction between a _______ and a _______ will form an _________.
A. amide, ester, alcohol B. amine, carboxylic acid, amide C. alcohol, amine, amide D. alcohol, alcohol, ester E. alcohol, carboxylic acid, amide
C H 3
48
Chem 110 Supplemental Problems
Answer Key Chapter 1 1. A 2. E 3. B 4. C 5. A Chapter 2 1. C 2. A 3. D 4. D 5. D Chapter 3 1. C 2. B 3. A 4. E 5. D 6. B 7. B 8. D 9. E Chapter 4 1. B 2. E 3. E 4. C 5. C Chapter 5 1. E 2. D 3. E 4. A 5. A 6. D 7. D
Chapter 6 1. D 2. D 3. B 4. E 5. A 6. C 7. E 8. C 9. E 10. D 11. B 12. C 13. C 14. B 15. A 16. B Chapter 7 1. B 2. B 3. D 4. A 5. D 6. D 7. B 8. E 9. E 10. E 11. D 12. D Chapter 8 1. A 2. B 3. D 4. B 5. E 6. B 7. B 8. B 9. A 10. B 11. A 12. A 13. B 14. C
Organic 1. D 2. D 3. D 4. A 5. C 6. B 7. B 8. E 9. D Chapter 9 1. E 2. C 3. D 4. A 5. A 6. E 7. B 8. C 9. E 10. D 11. D 12. E 13. B 14. A 15. D 16. E 17. B Chapter 10 1. B 2. B 3. B 4. A 5. C 6. D 7. A 8. E 9. B 10. D 11. A 12. E 13. B 14. C 15. E
Chapter 11 1. D 2. B 3. A 4. E 5. D 6. A 7. D 8. A 9. C 10. A 11. E 12. E 13. A Chapter 13 1. C 2. A 3. B 4. A 5. B 6. E 7. C 8. C 9. A 10. A 11. D 12. D 13. D 14. B 15. C 16. C 17. D 18. C Chapter 15 1. C 2. C 3. E 4. C 5. C 6. B 7. C 8. C 9. A 10. A 11. C 12. C 13. E
Chapter 18 1. A 2. D 3. B
Chemical Reactions 1. C 2. D 3. C 4. A 5. D 6. C 7. E 8. E 9. C
10. C 11. E 12. A 13. A 14. B 15. C 16. B 17. E 18. D 19. A 20. E 21. B
