Download xmas review questions 01516 with hints

Name ___________________________________ AP Solstice Greetings Review 015/16 
Justify each on separate paper – Large and legible!
Questions 1-3
(A) F
(B) S
(C) Mg
(D) Ar
(E) Mn
1. Forms monatomic ions with 2– charge in solutions
This is probably and easy one. Goes back to the element type and tendency to form various ions.
2. Forms a compound having the formula KXO4
Can you identify the charge that K forms? That tells you the overall charge of the XO4 ion and will
lead you to the identity of X. You should be able to narrow this down to two choices.
3. Forms oxides that are common air pollutants and that yield acidic solution in water
This element forms one of the XO acids which is one of the 5 (or 7) common strong acids.
Questions refer to atoms for which the occupied atomic orbitals are shown below.
4. Represents an atom that is chemically unreactive
What makes atoms truly stable?
5. Represents an atom in an excited state
Recall that atoms fill their orbitals according to the aufbau rule?
6. Represents an atom that has four valence electrons
The valence shell includes both s and p orbitals.
7. Represents an atom of a transition metal
Where do added electrons go when moving across the middle of the periodic table?
Questions refer to the following descriptions of bonding in different types of solids.
(A)
(B)
(C)
(D)
(E)
Lattice of positive and negative ions held together by electrostatic forces
Closely packed lattice with delocalized electrons throughout
Strong single covalent bonds with weak intermolecular forces
Strong multiple covalent bonds (including Pi-bonds) with weak intermolecular forces
Macromolecules held together with strong polar bonds
8. Cesium chloride, CsCl(s)
What type of bonding forms between metals and non-metals?
9. Gold, Au(s)
What type of bonding?
10. Carbon dioxide, CO2(s)
Draw the Lewis e dot and see!
11. Methane, CH4(s)
Draw the Lewis e dot for this one also.
Questions refer to the following elements.
(A) Lithium
(B) Nickel
(C) Bromine
(D) Uranium
(E) Fluorine
12. Is a gas in its standard state at 298 K
Small diatomic molecules have weak IM forces and tend to be gases.
13. Reacts with water to form a strong base
Another term for base is the name for the group from where this element comes.
Questions
(A)
Heisenberg uncertainty principle
(B)
Pauli exclusion principle
(C)
Hund’s rule (principle of maximum multiplicity)
(D)
Shielding effect
(E)
Wave nature of matter
14. Can be used to predict that a gaseous carbon atom in its ground state is paramagnetic
Paramagnetism is caused by unpaired electrons.
15. Explains the experimental phenomenon of electron diffraction
What is diffraction? What can be diffracted?
16. Indicates that an atomic orbital can hold no more than two electrons
That’s because they have opposite spins.
17. Predicts that it is impossible to determine simultaneously the exact position and the exact velocity
of an electron
Because an electron has both wave and particle properties, the more we know about one quality the
less we know about the other.
1s2 2s2 2p6 3s2 3p3
18. Atoms of an element, X, have the electronic configuration shown above. The compound most likely
formed with magnesium, Mg, is
(A) MgX
(C) MgX2
(E) Mg3X2
(B) Mg2X
(D) MgX3
The number of valence electrons will indicate its group # and its likely oxidation state value.
19. The elements in which of the following have most nearly the same atomic radius?
(A) Be, B, C, N
(D) C, P, Se, I
(B) Ne, Ar, Kr, Xe
(E) Cr, Mn, Fe, Co
(C) Mg, Ca, Sr, Ba
Process of elimination. Recall what trends in radius occur and the reason.
20. What number of moles of O2 is needed to produce 14.2 grams of P4O10 from P? (Molecular weight
P4O10 = 284)
(A) 0.0500 mole
(D) 0.250 mole
(B) 0.0625 mole
(E) 0.500 mole
(C) 0.125 mole
Did you notice that 142 is half of 284? 14.2 grams is what fraction of a mole of P 4O10? How many
moles of O would you need?– and obviously half that amount of O2! You could also create a
balanced equation and do simple stoichiometry. That might be easier.
CH4(g) + 2 O2(g) CO2(g) + 2 H2O(l)
∆H˚ = –889.1 kJ
∆Hf˚ H2O(l) = –285.8 kJ/mol
∆Hf˚ CO2(g) = –393.3 kJ/mol
21. What is the standard heat of formation of methane, ∆Hf˚ CH4(g), as calculated from the data
above?
(A) –210.0 kJ/mole
(D) 75.8 kJ/mole
(B) –107.5 kJ/mole
(E) 210.0 kJ/mole
(C) –75.8 kJ/mole
Recall that H for a reaction equals the  H formation of products minus the  H formation of
reactants? Set it up and solve for the unknown.
22. Pi (π) bonding occurs in each of the following species EXCEPT
(A) CO2
(C) CN–
(E) CH4
(B) C2H4
(D) C6H6
Try to visualize the molecular structure. Eliminate the one that you know only sigma bonds.
3 Ag(s) + 4 HNO3
3 + NO(g) + 2 H2O
23. The reaction of silver metal and dilute nitric acid proceeds according to the equation above. If 0.10
mole of powdered silver is added to 10. milliliters of 6.0–molar nitric acid, the number of moles of
NO gas that can be formed is
(A) 0.015 mole (C) 0.030 mole (E) 0.090 mole
(B) 0.020 mole (D) 0.045 mole
Limiting reactant prob. How many moles is 10 mL (0.010 L) of 6 M HNO3? Find which one will
obviously be used up first and use the mole ratios to find the moles of NO produced.
24. Which, if any, of the following species is in the greatest concentration in a 0.100–molar solution of
H2SO4 in water?
(A) H2SO4 molecules
(C) HSO4– ions
(B) H3O+ ions
(D) SO42– ions
(E) All species are in equilibrium and therefore have the same concentrations.
Sulfuric acid (a strong acid) ionizes completely.
25. Which of the following represents the ground state electron configuration for the Mn 3+ ion?
(Atomic number Mn = 25)
(A) l s2 2s22p6 3s23p63d4
(B) 1s2 2s22p6 3s23p63d5 4s2
(C) 1s2 2s22p6 3s23p63d2 4s2
(D) 1s2 2s22p6 3s23p63d8 4s2
(E) 1s2 2s22p6 3s23p63d3 4s1
Which electrons are lost first during formation of ions in the d block (transition) metals?
26. When 70. milliliter of 3.0-molar Na2CO3 is added to 30. milliliters of 1.0-molar NaHCO3 the resulting concentration of Na+ is
(A) 2.0 M
(C) 4.0 M
(E) 7.0 M
(B) 2.4 M
(D) 4.5 M
A dilution problem? Each solution is diluted to 100 mL but be careful. Combine the new molarities to
find the final …be careful though; the 3 M solution is a “two for one” if you see what I mean.
27. The net ionic equation for the reaction that occurs during the titration of nitrous acid with sodium
hydroxide is
(A) HNO2 + Na+ + OH–  NaNO2 + H2O
(B) HNO2 + NaOH  Na+ + NO2– + H2O
(C) H+ + OH–H2O
(D) HNO2 + H2O  NO2– + H3O+
(E) HNO2 + OH–  NO2– + H2O
Which solutions are strong electrolytes? Dissociate and cancel spectators.
28. A student wishes to prepare 2.00 liters of 0.100–molar KIO3 (molecular weight 214). The proper
procedure is to weigh out
(A) 42.8 grams of KIO3 and add 2.00 kilograms of H2O
(B) 42.8 grams of KIO3 and add H2O until the final homogeneous solution has a volume of 2.00
liters
(C) 21.4 grams of KIO3 and add H2O until the final homogeneous solution has a volume of 2.00
liters
(D) 42.8 grams of KIO3 and add 2.00 liters of H2O
(E) 21.4 grams of KIO3 and add 2.00 liters of H2O
Find moles (and then mass) of KIO3 required then recall that the solute is part of the solution when
combined with the solvent.
29. A 20.0–milliliter sample of 0.200–molar K2CO3 solution is added to 30.0 milliliters of 0.400–molar
Ba(NO3)2 solution. Barium carbonate precipitates. The concentration of barium ion, Ba2+, in
solution after reaction is
(A) 0.150 M
(C) 0.200 M
(B) 0.160 M
(D) 0.240 M
(E) 0.267 M
A semi-complex math problem involving limiting factors. Ba ions combine with CO3 ions to form
BaCO3 precipitate. Need to find the number of moles of Ba and moles of CO3 in their original
solutions, and then see how much Ba is left over in excess. Oh… remember also, the final solution
volume is the sum of the two smaller volumes. Molarity = moles/liters.
30. A 27.0–gram sample of an unknown hydrocarbon was burned in excess oxygen to form 88.0 grams
of carbon dioxide and 27.0 grams of water. What is a possible molecular formula of the hydrocarbon?
(A) CH4
(C) C4H3
(E) C4H10
(B) C2H2
(D) C4H6
Did you see the significance of the 88 g CO2 and 27 g H2O? CO2 is 44 g/mole; while H20 is 18
g/mol. Convert mass of CO2 into moles for moles of C; convert mass of H2O to moles of H2O
and then to moles of H (notice 2 mol of H in each mole H2O?). That’s gets your empirical
formula. Most of the answers don’t fit that ratio anyway. Mass of original hydrocarbon sample
is just a distraction… I think ;-)
31. In which of the following compounds is the mass ratio of chromium to oxygen closest to 1.62 to
1.00 ?
(A) CrO3
(B) CrO2
(C) CrO
(D) Cr2O
(E) Cr2O3
I’m not sure what the best strategy is for this. I looked up the atomic mass of Cr and that of O
and looked first at the 1:1 ratio. I think a quick guess and check would be your best strategy.
Try doing it without the calculator first. Check it after if you like.
H2(g) + 1/2 O2(g)  H2O(l)
ΔH0 = –286 kJ
2 Na(s) + 1/2 O2(g)  Na2O(s)
ΔH0 = –414 kJ
Na(s) + 1/2 O2(g) + 1/2 H2(g) NaOH(s) ΔH0 = –425 kJ
32. Based on the information above, what is the standard enthalpy change for the following reaction?
Na2O(s) + H2O(l)  2 NaOH(s)
(A) –1,125 kJ
(B) –978 kJ
(C) –722 kJ (D) –150 kJ
(E) +275 kJ
Recall that H for a reaction equals the  H formation of products minus the  H formation of
reactants? Round all values to 1 sig fig in your “setup” and calculate. Choices are far enough apart
to get you close enough to the correct answer.
33. Given that a solution is 5 percent sucrose by mass, what additional information is necessary to
calculate the molarity of the solution?
I. The density of water
II. The density of the solution
III. The molar mass of sucrose
(A) I only
(B) II only
(C) III only
(D) I and III
(E) II and III
Think of this as a two-step dimensional analysis problem. You are converting percent (grams of
sucrose per 100 grams of solution) into molarity (moles of sucrose per liter of solution.)
34. The system shown above is at equilibrium at 28_C. At this temperature, the vapor pressure of
water is 28 millimeters of mercury. The partial pressure of O2(g) in the system is
(A) 28 mm Hg
(B)
56 mm Hg
(C) 133 mm Hg
(D)
161 mm Hg
(E) 189 mm Hg
A question involving partial pressures: Recall that the total pressure inside the spherical container is
the sum of two gases – oxygen and…._______. The difference in levels within the mercury filled tube
is the total pressure exerted.
Mass of an empty container
Mass of the container plus the
solid sample
Volume of the solid sample
3.0 grams
25.0 grams
11.0 cm3
35. The data above were gathered in order to determine the density of an unknown solid. The density
of the sample should be reported as
(A) 0.5 g/cm3 (B) 0.50 g/cm3
(C) 2.0 g/cm3 (D) 2.00 g/cm3
E) 2.27 g/cm3
Careful with this one. It’s a rounding problem. Recall that subtracting to find the mass of the solid
involves rounding using a different rule from the rule applied when you do the M/V division.
36. A sample of an ideal gas is cooled from 50.00C to 25.00C in a sealed container of constant volume.
Which of the following values for the gas will decrease?
I. The average molecular mass of the gas
II. The average distance between the molecules
III. The average speed of the molecules
(A) I only
(B)
II only (C)
III only
(D) I and III
(E)
II and III
A Kinetic theory problem. How does cooling a gas affect its molecules. Notice the container
doesn’t change?
37. Which of the following solids dissolves in water to form a colorless solution?
(A) CrCl3
(B)
FeCl3 (C)
CoCl2
(D) CuCl2
(E)
ZnCl2
One of these metals is not a true transition metal (although it is, like the others; a member of
the d-block) You recall that many transition metals form brightly colored ions in solution?
38. Which of the following has the lowest conductivity?
(A) 0.1 M CuSO4
(B)
0.1 M KOH
(C) 0.1 M BaCl2,
(D)
0.1 M HF
(E) 0.1 M HNO3
Conductivity is a function of dissolved ions (remember “electrolytes”?) One is a weak electrolyte
which fails to dissociate significantly into ions even when dissolved.
39. When dilute nitric acid was added to a solution of one of the following chemicals, a gas was
evolved. This gas turned a drop of limewater, Ca(OH)2, cloudy, due to the formation of a white
precipitate. The chemical was
(A) household ammonia, NH3
(B) baking soda, NaHCO3
(C) table salt, NaCl
(D) epsom salts, MgSO4.7H2O
(E) bleach, 5% NaOCl
This question will be easier following a further discussion of acids and bases in an upcoming
topic. For now try to connect it to a gas producing reaction in our reaction type lab when an acid
(vinegar) was added to one of these. The ion of interest is usually insoluble but not with a group
1 ion. Write the equations for both reactions to justify.
40. If 87 grams of K2SO4 (molar mass 174 grams) is dissolved in enough water to make 250 milliliters of
solution, what are the concentrations of the potassium and the sulfate ions?
[K+]
[SO42–]
(A) 0.020 M
0.020 M
(B) 1.0 M
2.0 M
(C) 2.0 M
1 .0 M
(D) 2.0 M
2.0 M
(E) 4.0 M
2.0 M
Notice that 87g is half of 174? The moles of K doubles after dissolving. Setup your Molarities.
Dividing by .25 L is like multiplying by 4). Again try without a calculator by doing metal
math.
41. All of the following statements concerning the characteristics of the halogens are true EXCEPT:
(A) The first ionization energies (potentials) decrease as the atomic numbers of the halogens
increase.
(B) Fluorine is the best oxidizing agent.
(C) Fluorine atoms have the smallest radii.
(D) Iodine liberates free bromine from a solution of bromide ion.
(E) Fluorine is the most electronegative of the halogens.
“True (except)” means you are looking for the false statement. I used process of elimination.
Four are clearly true. The false one makes sense in light of their relative “activity” (see
regents reference table J for a reminder)
42. What volume of 0.150–molar HCl is required to neutralize 25.0 milliliters of 0.120–molar Ba(OH)2?
(A) 20.0 mL
(B) 30.0 mL
(C) 40.0 mL
(D) 60.0 mL
(E) 80.0 mL
A titration problem. Create a balanced equation showing the reaction of HCl with Ba(OH)2 to
form water and barium chloride (BaCl2?) Then do a quick stoichiometry. Recall
MaVa=2MbVb. The molarity of the base doubles when dissolved. Ba(OH)2
43. It is suggested that SO2 (molar mass = 64.1 grams), which contributes to acid rain, could be
removed from a stream of waste gases by bubbling the gases through 0.25–molar KOH, thereby
producing K2SO3. What is the maximum mass of SO2 that could be removed by 1,000. liters of the
KOH solution?
(A) 4.0 kg (B) 8.0 kg (C) 16 kg (D) 20. kg (E) 40. Kg
You should recognize this as a stoichiometry problem since you are given info relating to KOH but
asked about SO2. You will need to create a balanced equation and in the process you will notice
that water must be another product. Find moles of KOH present in the solution and from that the
moles and mass of SO2 that can be reacted (removed).
44. When a 1.00–gram sample of limestone was dissolved in acid, 0.38 gram of CO2 was generated. If
the rock contained no carbonate other than CaCO3, what was the percent of CaCO3 by mass in the
limestone?
(A) 17% (B) 51% (C) 64% (D) 86% (E) 100%
A semi-stoichiometry problem. Since the CaCO3 becomes CO2 you should focus on the carbon.
Finding the moles of CO2 will tell you the moles of C, and thus the moles of CaCO3 it stated
as. From this the mass and percent of CaCO3 can be calculated.
I2(g) + 3 Cl2(g) 2 ICl3(g)
45. According to the data in the table below, what is the value of H for the reaction represented
above?
Bond Average Bond Energy
(kilojoules/mole)
I–I
149
Cl–Cl
239
I–Cl
208
(A) –860 kJ (B) –382 kJ (C) +180 kJ (D) +450 kJ (E) +1,248 kJ
The normal way to look at this based on bond energy: H =  H bonds broken -  H bonds formed:
Or you could switch the sign to negative for the bonds formed (since bond forming is exothermic) and
then add them together (Hess’ law). Hint: the overall reaction is exothermic.
46. At 250C, a sample of NH3 (molar mass 17 grams) effuses at the rate of 0.050 mole per minute.
Under the same conditions, which of the following gases effuses at approximately one–half that
rate?
(A) O2 (molar mass 32 grams)
(B) He (molar mass 4.0 grams)
(C) CO2 (molar mass 44 grams)
(D) Cl2 (molar mass 71 grams)
(E) CH4 (molar mass 16 grams
Recall temperature (kinetic energy) is equal between two samples at the same temp and
KE=1/2mV2. For the two gases then 1/2mV2 for one gas is equal to 1/2mV2 for the other. If
you let V=2 for the NH3 then V for the other gas = 1 (one half of the NH3’s )
so set the two equal and then solve for mass: ½(17)(2)2 = ½ m(1)2
47. Which of the following molecules has a dipole moment of zero?
(A) C6H6 (benzene) (B) NO (C) SO2 (D) NH3 (E) H2S
The molecule must be totally nonpolar. Process of elimination will lead you to an answer if nothing
else works.
48. When a solution of sodium chloride is vaporized in a flame, the color of the flame is
(A) blue
(B)
yellow (C)
green
(D) violet
(E)
white
Recall the color of the sodium flame? Hint: sodium is a common contaminant of other chemicals.
49. A hot-air balloon, shown above, rises. Which of the following is the best explanation for this
observation?
(A) The pressure on the walls of the balloon increases with increasing temperature.
(B) The difference in temperature between the air inside and outside the balloon produces
convection currents.
(C) The cooler air outside the balloon pushes in on the walls of the balloon.
(D) The rate of diffusion of cooler air is less than that of warmer air.
(E) The air density inside the balloon is less than that of the surrounding air.
This is a question which you may recognize based on past experience. A balloon is a constant
pressure system. Air pressure inside and outside must be the same. Warm air molecules
move faster…so you need less of them to create the same pressure.
50. The melting point of MgO is higher than that of NaF. Explanations for this observation include
which of the following?
I. Mg2+ is more positively charged than Na+.
II. O2– is more negatively charged than F–.
III. The O2– ion is smaller than the F– ion.
(A) II only (B) I and II only (C) I and III only (D) II and III only (E) I, II, and III
Lattice energy is a function of coulombs law: Eel = kq1q2/r2 a function of. Difference in charge vs.
distance between charges.
H2Se(g) + 4 O2F2(g)  SeF6(g) + 2 HF(g) + 4 O2(g)
51. Which of the following is true regarding the reaction represented above?
(A) The oxidation number of O does not change.
(B) The oxidation number of H changes from -1 to +1.
(C) The oxidation number of F changes from +1 to -1.
(D) The oxidation number of Se changes from -2 to +6.
(E) It is a disproportionation reaction for F.
Check out each choice in turn until you get to a true statement. Recall that compounds are
neutral and general rules include – atoms in the front of a formula are positive, the element
in back is usually negative. F is always -1 unless its elemental (as in F2) O is usually -2 but
there are exceptions.
52. Types of hybridization exhibited by the C atoms in propene, CH3CHCH2, include which of the
following?
I. sp II.
sp2
III.
sp3
(A) I only (B) III only (C) I and II only (D) II and III only (E) I, II, and III
Draw it out recalling that H forms only one bond while C forms 4. Recall that hybridization
follows from electron domains.
53. A 1.0 L sample of an aqueous solution contains 0.10 mol of NaCl and 0.10 mol of CaCl 2. What is the
minimum number of moles of AgNO3 that must be added to the solution in order to precipitate all
of the Cl– as AgCl(s) ? (Assume that AgCl is insoluble.)
(A) 0.10 mol (B) 0.20 mol (C) 0.30 mol (D) 0.40 mol (E) 0.60 mol
Seems easy…notice that NaCl has one Cl, while CaCl2 has 2 Cl’s ?
Ionization Energies for element X (kJ mol-1)
First
Second
Third
Fourth
Fifth
580
1,815
2,740
11,600
14,800
54. The ionization energies for element X are listed in the table above. On the basis of the data,
element X is most likely to be
(A) Na (B) Mg (C) Al (D) Si (E) P
Notice the jump between the third and 4th IE? This represents which electrons are being
subsequently removed.
55. Of the following molecules, which has the largest dipole moment?
(A) CO
(B) CO2
(C) O2 (D) HF (E) F2
Which molecule is most polar? Recall polarity is related to arrangement and polarity of bonds:
the difference in electronegativity between bonded atoms.
56. A rigid metal tank contains oxygen gas. Which of the following applies to the gas in the tank when
additional oxygen is added at constant temperature?
(A) The volume of the gas increases.
(B) The pressure of the gas decreases.
(C) The average speed of the gas molecules remains the same.
(D) The total number of gas molecules remains the same.
(E) The average distance between the gas molecules increases.
Uh…seems simple. Adding more molecules to a fixed volume. More molecules in the same space.
Visualize how the systems would be different. Eliminate all but the correct response?
57. When hafnium metal is heated in an atmosphere of chlorine gas, the product of the reaction is
found to contain 62.2 percent Hf by mass and 37.4 percent Cl by mass. What is the empirical
formula for this compound?
(A) HfCl (B) HfCl2 (C) HfCl3 (D) HfCl4 (E) Hf2Cl3
You might recognize that the atomic mass of Hf is 5 times the mass of Cl, but 62 is only 1 2/3
greater than 37.4. The number of Cl’s must increase to reduce the ratio. By how much?
You could of course do a quick mole conversion – recall ballpark round off to avoid calculator use
– and find the mole ratio of Hf to Cl just as quickly.
58. Which of the following techniques is most appropriate for the recovery of solid KNO 3 from an
aqueous solution of KNO3?
(A) Paper chromatography
(B) Filtration
(C) Titration
(D) Electrolysis
(B) Evaporation to dryness
You can’t need a hint, but do explain the process which would lead to creating a sample of solid
KNO3.
59. In the periodic table, as the atomic number increases from 11 to 17, what happens to the atomic
radius?
(A) It remains constant.
(B) It increases only.
(C) It increases, then decreases.
(D) It decreases only.
(E) It decreases, then increases.
This one runs counter to logic. Atomic size depends on the number of occupied energy shells, and
the charge of the nucleus pulling the electrons in.
60. Which of the following is a correct interpretation of the results of Rutherford’s experiments in
which gold atoms were bombarded with alpha particles?
(A) Atoms have equal numbers of positive and negative charges.
(B) Electrons in atoms are arranged in shells.
(C) Neutrons are at the center of an atom.
(D) Neutrons and protons in atoms have nearly equal mass.
(E) The positive charge of an atom is concentrated in a small region.
Describe the results of the experiment that lead to this conclusion.
W+XY+Z
61. Gases W and X react in a closed, rigid vessel to form gases Y and Z according to the equation above.
The initial pressure of W(g) is 1.20 atm and that of X(g) is 1.60 atm. No Y(g) or Z(g) is initially
present. The experiment is carried out at constant temperature. What is the partial pressure of
Z(g) when the partial pressure of W(g) has decreased to 1 .0 atm?
(A) 0.20 atm (B) 0.40 atm (C) l.0 atm (D) 1.2 atm (E) l.4 atm
Since pressure is proportional number of molecules (moles) the decrease in pressure of gas W
should produce a stoichiometrically equivalent increase the pressure from Z.
10 HI + 2 KMnO4 + 3 H2SO4  5 I2 + 2 MnSO4 + K2SO4 + 8 H2O
62. According to the balanced equation above, how many moles of HI would be necessary to produce
2.5 mol of I2, starting with 4.0 mol of KMnO4 and 3.0 mol of H2SO4?
(A) 20. (B) 10. (C) 8.0 (D) 5.0 (E) 2.5
Stoichiometry: a strange question. In order to produce 2.5 moles of I2 there must already be
sufficient quantities of KMnO4 and H2SO4. The question relates to the amount of HI you must
react to produce the desired quantity of I2.
63. A yellow precipitate forms when 0.5 M NaI(aq) is added to a 0.5 M solution of which of the
following ions?
(A) Pb2+(aq) (B) Zn2+(aq) (C) CrO42–(aq) (D) SO42–(aq) (E) OH–(aq)
Since group 1 ions are always soluble the precipitate must be one of I- ? Which ion is an
exception to the solubility of I-? I think you produced this in the lab.
64. A 40.0 mL sample of 0.25 M KOH is added to 60.0 mL of 0.15 M Ba(OH)2. What is the molar
concentration of OH–(aq) in the resulting solution? (Assume that the volumes are additive.)
(A) 0.10 M (B) 0.19 M (C) 0.28 M (D) 0.40 M (E) 0.55 M
If you mix two solution with different concentrations together where do you expect the final
concentration to be. You could also look at it as two dilution problems.
NH4NO3(s) --> N2O(g) + 2 H2O(g)
65. A 0.03 mol sample of NH4NO3(s) is placed in a 1 L evacuated flask, which is then sealed and heated.
The NH4NO3(s) decomposes completely according to the balanced equation above. The total
pressure in the flask measured at 400 K is closest to which of the following? (The value of the gas
constant, R, is 0.082 L atm mol-1 K-1.)
(A) 3 atm (B) 1 atm (C) 0.5 atm (D) 0.1 atm (E) 0.03 atm
Each mole of reactant breaks into 3 moles of gaseous product.
C2H4(g) + 3 O2(g)  2 CO2(g) + 2 H2O(g)
66. For the reaction of ethylene represented above, ΔH is -1,323 kJ. What is the value of ΔH if the
combustion produced liquid water H2O(l), rather than water vapor H2O(g) ? (ΔH for the phase
change H2O(g)  H2O(l) is –44 kJ mol-1.)
(A) -1,235 kJ (B) -1,279 kJ (C) -1,323 kJ (D) -1,367 kJ (E) -1,411 kJ
The condensation of water vapor to liquid water is also exothermic.
67. Equal numbers of moles of He(g), Ar(g), and Ne(g) are placed in a glass vessel at room
temperature. If the vessel has a pinhole-sized leak, which of the following will be true regarding
the relative values of the partial pressures of the gases remaining in the vessel after some of the
gas mixture has effused?
(A) PHe < PNe < PAr
(B) PHe < PAr < PNe
(C) PNe < PAr < PHe
(D) PAr < PHe < PNe
(E) PHe = PAr = PNe
Because KE = ½ mV2 , molecules effuse at a rate inversely proportional to the square roots of their
masses. (this is Graham’s law)
68. In which of the following processes are covalent bonds broken?
(A) I2(s) I2(g)
(B)
CO2(s)  CO2(g)
(C) NaCl(s)  NaCl(l)
(D)
C(diamond)  C(g)
(E) Fe(s)  Fe(l)
Really, AP? Between what types of elements do we expect to find covalent bonds?
69. What is the final concentration of barium ions, [Ba2+], in solution when 100. mL of 0.10 M BaCl2(aq)
is mixed with 100. mL of 0.050 M H2SO4(aq)?
(A) 0.00 M
(B) 0.012 M
(C) 0.025 M
(D) 0.075 M
(E) 0.10 M
Ba combines with SO4 to form BaSO4 precipitate. Which of the two ions will limit the
precipitate produced? The excess must be Ba2+ . How much is left over after all the SO4 is
precipitated? Be careful, the final volume is now larger so the final molarity is lower.
70. When 100 mL of 1.0 M Na3PO4 is mixed with 100 mL of 1.0 M AgNO3, a yellow precipitate forms
and [Ag+] becomes negligibly small. Which of the following is a correct listing of the ions remaining
in solution in order of increasing concentration?
(A) [PO43–] < [NO3–] < [Na+]
(B) [PO43–] < [Na+] < [NO3–]
(C) [NO3–] < [PO43–] < [Na+]
(D) [Na+] < [NO3–] < [PO43–]
(E) [Na+] < [PO43–] < [NO3–]
Do the math to justify, but can you do this one logically? Na + is a spectator producing 3 moles of
Na+ ions for every mole of Na3PO4, PO43- is reacted but is in excess so there is some left over.
How much was there initially and how much reacts may be a question. NO3 is also a
spectator.
71. After completing an experiment to determine gravimetrically the percentage of water in a hydrate,
a student reported a value of 38 percent. The correct value for the percentage of water in the
hydrate is 51 percent. Which of the following is the most likely explanation for this difference?
(A) Strong initial heating caused some of the hydrate sample to spatter out of the crucible.
(B) The dehydrated sample absorbed moisture after heating.
(C) The amount of the hydrate sample used was too small.
(D) The crucible was not heated to constant mass before use.
(E) Excess heating caused the dehydrated sample to decompose.
Think about how the mass of water lost from the original sample is calculated. Only one of these
would create a lower than expected value for this.
72. The volume of distilled water that should be added to 10.0 mL of 6.00 M HCl(aq) in order to
prepare a 0.500 M HCl(aq) solution is approximately
(A) 50.0 mL (B) 60.0 mL (C) 100. mL (D) 110. mL (E) 120. mL
Dilution MaVa=MbVb. But realize you’re already starting with 10ml of volume.
73. Which of the following gases deviates most from ideal behavior?
(A) SO2 (B) Ne (C) CH4 (D) N2 (E) H2
Molecules deviate from ideal behavior since real gas molecules have attractive forces and real gas
molecules have volume. When molecules are close together due to high pressure and are moving
slowly these deviations become more apparent.