Download Extra Material

CHEM 1301
POSSIBLE EXTRAS REVIEW
Chapter Ten Stuff
Differences between solids, liquids and gases.
Intermolecular Forces – important in solid state
Stronger they are, higher the melting and boiling points:
1.
IONIC
all ionic compounds have IONIC intermolecular forces
2.
METALLIC all metals have METALLIC intermolecular forces
Both of these are “infinite” networks – go on and on and on
COVALENT molecules have:
3.
HYDROGEN BONDS for H attached to F, O, N (or Cl)
4.
DIPOLE-DIPOLE if the molecule is polar [has a dipole – opposite ends]
5.
LONDON DISPERSON if the molecule is not polar. i. e., has no polar bonds or the polar
bonds cancel out.
Typical questions:
For each of the following solids, assign whether the intermolecular forces are
(A)
(D)
ionic
dipole-dipole
(B)
(E)
1.
5.
9.
NaCl
H2O
KBr
11.
Which of the following would be expected to have the highest boiling point?
(A)
NaCl (B)
12.
Which substance has the highest boiling point?
(A)
CH4
13.
Which group of substances is arranged in order from the highest to the lowest melting point?
(A)
HF>H2>NaF (B)
2.
6.
10.
Na
(B)
metallic
London dispersion
3.
7.
Cl2
NH3
CH4
(C)
He
Cl2
(C)
Hg
FBr
(D)
all the same
(C)
HF
NaF>H2>HF (C)
hydrogen bonding
4.
8.
(E)
can’t tell
(D)
Cl2
HF>NaF>H2 (D)
1
Fe
CO2
NaF>HF>H2
14.
Arrange KCl, NH3, and CH4 in order of increasing boiling point.
(A)
(C)
CH4<KCl<NH3
CH4<NH3<KCl
15.
Which has the highest melting point?
(A)
S8
16.
Which of the following molecules will not form hydrogen bonds?
(B) NH3<KCl<CH4
(D) NH3<CH4<KCl
(B)
H2O
(C)
Ar
(D)
BaF2
O
(A)
H3C
CH2
C
(B)
OH
HF
H
O
H3C
(C)
CH2
C
(D)
CH3
H3C
CH2
N
17.
In hydrogen iodide __________________ are the most important intermolecular forces.
(A)
(C)
dipole-dipole forces
hydrogen bonding
(B)
(D)
CH3
London dispersion forces
ionic
GASES
Pressure
Temperature
Volume
n
measured in atmospheres or mmHg [1 atm = 760 mm Hg]
measured in Kelvin [K = °C + 273]
measured in liters
moles
Pressure is proportional to temperature and number of moles and inversely proportional to volume
Put together the relationships using the correct units:
PV = nRT
[R = 0.082 l atm / mol K]
18.
When the temperature in Kelvin of a fixed quantity of an ideal gas is quadrupled and the
pressure is doubled, what is the net effect on the volume of the gas?
(A)
(C)
The volume remains constant.
The volume increases four–fold.
(B)
(D)
The volume increases two–fold.
The volume increases eight–fold.
2
19.
Air is sealed in a vessel at 273 °C and then cooled to 0 °C. If the vessel itself does not
contract, the pressure inside the vessel will become
(A)
(D)
twice its original value.
one-fourth of its original value.
20.
Both the pressure and the absolute temperature of a certain gas sample are doubled. In the
absence of dissociation, the volume of the gas is
(A)
(D)
quadrupled.
decreased by one–half.
21.
If a gas at 1 atm and 273 K occupies 3.36 L, how many moles does it contain?
(A)
0.15
22.
A gas sample occupies a volume of 16.4 L at 27 °C and 0.300 atm. How many moles of gas
are present?
(A)
0.200
23.
What is the volume of 2.00 mol of helium gas at 27 °C and 3.00 atm?
(A)
6.1 x10–2 L
24.
How many moles of gas are in a sample with a volume of 500 mL at 25 °C and 0.460 atm?
(A)
0.00941 mol (B)
(B)
(B)
(E)
0.81
(B)
none of these.
(C)
one-half of its original value.
decreased by one–fourth.
unchanged.
(C)
0.450
(B)
(B)
(E)
3.4
(C)
16.4 L
3.50
(C)
10.2 mol
1.48 L
(C)
0.160 mol
(C)
(D)
zero.
doubled.
6.7
(D)
10.0
(D)
44.8 L
(D)
13.4 mol
This can be combined into the Roadmap:
grams of A
g
grams of B
g
Particles
NB
Particles
NA
n = g / MM
N = 6.02 x 1023 n
23
N = 6.02 x 10 n
Moles of A
nA
Volume of Gas,
V at Temp, T, and
Pressure, P
nB = (b/a)nA
pV = nRT
M=n/L
Volume of
Solution, L
Moles of B
nB
Volume of Gas,
V at Temp, T, and
pV = nRT
Pressure, P
Volume of
Solution, L
3
25.
What is the number of molecules in 1.00 mL of an ideal gas at 1 atm and 273 K?
(A)
2.69 x 1022
26.
How many grams of CO2 would occupy an 8.8 L flask at 300 K and 1.1 atm?
(A)
14 g
27.
What volume will 5.10 g of sulfuryl fluoride, SO2F2, occupy at 1 atm and 273 K?
(A) 0.056 L
28.
(B)
(B)
(B) 1.12 L
2.69 x 1019
17 g
(C) 2.24 L
6.02 x 1020
(C)
(C)
22 g
(D) 11.2 L
(D)
(D)
2.24 x 1019
32 g
(E) 22.4 L
What volume does 16.00 g of oxygen gas (O2) occupy at 546 °C and 2.00 atm?
(A) 5.6 L
(B) 11.2 L
(C) 22.4 L
(D) 67.2 L
(E)
16.8 L
29.
When 18.0 g of water is heated to steam at 100 °C, the volume at 1 atm pressure is
approximately?:
(A)
13.2 mL
30.
Benzene, C6H6, can be burned in oxygen according to the equation:, 2C6H6(l) + 15O2(g) →
12CO2(g) + 6H2O(g). If 5.0 L of oxygen measured at 1 atm and 273 K were required to burn
a given amount of benzene, the 1 atm and 273 K volume of CO2 formed would be
(A)
17.9 L
31.
What mass of CaCO3 will produce 8.0 L of CO2 (measured at 1 atm and 273 K) in the
reaction, CaCO3(s) → CaO(s) + CO2(g) ?
(A)
4.46 g
32.
How many liters of CO2 gas at 1 atm and 273 K can be obtained by completely burning one
C3H8(g) + O2(g) → CO2(g) + H2O(g)
mole of C3H8?
(A)
11.2
33.
What volume of pure N2O at 1 atm and 273 K could be prepared by the controlled
decomposition of 8.00 g of ammonium nitrate - NH4NO3(s) → N2O(g) + 2H2O(l)
(A)
1.12 L
(B)
(B)
(B)
(B)
(B)
18.6 L
4 L
12.5 g
44.8
2.24 L
(C)
1800 mL
(C)
33.4
(C)
35.7 g
(C)
67.2
(C)
3.36 L
4
(D)
(D)
(D)
11 L
(E)
5.5 L
280 g
(D) 112
(D)
30.6 L
4.48 L
34.
How many liters of hydrogen gas with an excess of nitrogen at 1 atm and 273 K are required
to prepare 45.0 g of ammonia? 3 H2(g) + N2(g) → 2 NH3(g)
(A)
39.5
35.
How many liters of hydrogen at 1 atm and 273 K can be produced by the reaction of 9.00 g
of Al with excess dilute H2SO4: 2Al(s) + 3H2SO4(aq) → Al2(SO4)3(aq) + 3H2(g)
(A)
5.61 L
(B)
30.0
(B)
(C)
11.2 L
59.3
(C)
(D)
33.6 L
(D)
88.9
67.2 L
More Reactions
We started talking about REDOX equations. You need to be able to:
a)
b)
c)
assign oxidation numbers:
look at an equation and say whether it is a redox equation by change in charge or oxidation
number
if redox, which atom is reduced and which is oxidized
36.
Calculate the oxidation number of the chlorine in perchloric acid, HClO4, a strong acid.
(A)
-1
37.
Calculate the oxidation number of sulfur in sodium metabisulfite, Na2S2O5.
(A)
-2
38.
Sodium tripolyphosphate is used in detergents to make them effective in hard water.
Calculate the oxidation number of phosphorus in Na5P3O10.
(A)
+3
39.
Calculate the oxidation number of iodine in I2.
(A)
-1
40.
The oxidation numbers of P, S and Cl in H2PO2-, H2S and KClO4 are, respectively
(A)
(E)
-1, -1, +3
-1, -2, +3
41.
Identify which species is reduced in the following redox reaction.
(B)
(B)
(B)
+
Cu(s)
Cu(s)
(D)
+10
(C)
+1, -2, +7
(D)
+4
(C)
0
(B)
+5
(C)
+5
(B)
Hg2+(aq)
(C)
+2
(B)
Hg2+(aq)
(A)
+4
+1
(C)
+1, +2, +7
→ Cu2+(aq)
+
5
+5
(D)
+15
(D)
+7
(D)
-1, -2, +7
Hg(l)
Cu2+(aq)
(C)
+7
(D)
Hg(l)
42.
Sodium thiosulfate, Na2S2O3, is used as a “fixer” in black and white photography. Identify
which atom is oxidized in the reaction of thiosulfate with iodine.
2S2O32-(aq)
+
I2(aq)
(B)
→
I-
S4O62-(aq)
2I-(aq)
+
(A)
I
43.
Which one of the following is not a redox reaction?
(A)
(B)
(C)
(D)
(E)
2H2(g) + O2(g) → 2H2O(l)
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
H2O(l) + NH3(g) → NH4+(aq) + OH-(aq)
6FeSO4(aq) + K2Cr2O7(aq) + 7H2SO4(aq) →
Cr2(SO4)3(aq) + 3Fe2(SO4)3(aq) + K2SO4(aq)
Cl2(g) + 2KBr(aq) → Br2(l) + 2KCl(aq)
44.
Which one of the following is not a redox reaction?
(A)
(B)
(C)
(D)
(E)
2H2O2(aq) → 2H2O(l) + O2(g)
N2(g) + 3H2(g) → 2NH3(g)
BaCl2(aq) + K2CrO4(aq) → BaCrO4(aq) + 2KCl(aq)
2Al(s) + Fe2O3(s) → Al2O3(s) + 2Fe(s)
2H2O(g) → 2H2(g) + O2(g)
45.
Which one of the following is a redox reaction?
(A)
(B)
(C)
(D)
(E)
2Na(g) + Cl2(g) → 2NaCl(s)
Ba2+(aq) + SO42-(aq) → BaSO4(s)
K2Cr2O7(aq) + 2KOH(aq) → 2K2CrO4(aq) + H2O(l)
Na2CO3(s) + 2HCl(aq) → 2NaCl(aq) + CO2(g) + H2O(l)
H2O(l) → H+(aq) + OH-(aq)
46.
Which one of the following is not a redox reaction?
(A)
(B)
(C)
(D)
(E)
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
H2(g) + Cl2(g) → 2HCl(g)
2H2O2(aq) → 2H2O(l) + O2(g)
Fe2O3(s) + 3H2SO4(aq) → Fe2(SO4)3(aq) + 3H2O(l)
2KMnO4(aq) + 10FeSO4(aq) + 8H2SO4(aq) →
K2SO4(aq) + 2MnSO4(aq) + 5Fe2(SO4)3(aq) +
(C)
S
(D)
+
O
7H2O(l)
8H2O(l)
We discussed the relative “strength” of metals and getting electricity from redox reactions.
6
Acids and Bases
a)
b)
c)
d)
e)
Arrhenius definition
Bronsted-Lowry definition
strong and weak
neutralization
pH = -log [H+] or –log [H3O+]
47.
Which of the following is a weak acid?
(A)
H2SO4
48.
Which of the following is a strong acid?
(A)
H3PO4
49.
Which of the following is a strong base?
(A)
NH3
50.
Which of the following is a weak base?
(A)
NH3
51.
Which one of the following substances is a strong acid?
(A)
(D)
HNO3
CH3COOH
52.
What is the pH of a 0.75 M HNO3 solution?
(A)
0.12
53.
What is the pH of a 0.00200 M HClO4 solution?
(A)
0.995
54.
What is the pH of a 0.050 M HBr solution?
(A)
0.89
(B)
(B)
(B)
(B)
HNO3
HNO3
(B)
(B)
(C)
Ba(OH)2
Sr(OH)2
(B)
(E)
(B)
(C)
0.29
1.378
1.12
HF
(D)
HBr
HF
(D)
CH3COOH
(C)
Al(OH)3
(D)
B(OH)3
(C)
Ba(OH)2
(D)
NaOH
H2CO3
H3PO4
(C)
(C)
NH3
(D)
0.82
2.699
(D)
6.215
1.30
(D)
3.00
0.63
(C)
(C)
Answers:
1 – 30:
31-54:
AEBBC|CDEAE|ACDCD|CABEE|AABAB|BBEEB|
CCBDB|DCBBB|ACCCA|DCBBA|AACC
7