AP Chemistry
Fall Semester Practice Multiple Choice
Name ________________________
Period _____
Multiple Choice: Briefly show/explain why the multiple choice answer is correct in the space provided (no calculator).
1. Copper has two naturally occurring isotopes, 63Cu and
Questions 6-8
65Cu. What is the abundance of 65Cu if the average atomic
(A) Heisenberg uncertainty principle
mass of copper is 63.5?
(B) Pauli exclusion principle
(A) 90%
(B) 70%
(C) 50%
(D) 25%
(C) Hund's rule (principle of maximum multiplicity)
(D) Shielding effect
6. Can be used to predict that a gaseous carbon atom in its
ground state is paramagnetic.
2. Which of the following particles is emitted by an atom of
39Ca when it decays to produce an atom of 39K?
(A) 10n
(B) 11H
(C) (D) +
7. Indicates that an atomic orbital can hold no more than two
electrons.
3.
After 195 days, a 10.0 g sample of pure 95Zr has decayed
to the extent that only 1.25 g of the original 95Zr remains.
The half-life of 95Zr is closest to
(A) 195 days
(B) 98 days
(C) 65 days
(D) 49 days
Questions 4-5 The diagram shows the energy levels (in eV) for
hydrogen gas.
8.
Predicts that it is impossible to determine simultaneously
the exact position and the exact velocity of an electron.
9.
Which set of quantum numbers (n, l, ml, ms) best describes
the valence electron of highest energy in a ground-state
gallium atom (Z = 31)?
(A) 4,0,0,½ (B) 4,0,1,½ (C) 4,1,1,½ (D) 4,1,2,½
Questions 10-12 refer to neutral atoms for which the atomic
orbitals are represented below.
(A) 1s()
(B) 1s() 2s() 2p()()( )
(C) 1s() 2s() 2p()()()
(D) [Ar] 4s() 3d()()()()()
10. Is in an excited state
11. Has exactly five valence electrons
12. Has the highest first ionization energy
4.
What is the energy, in eV, of a photon emitted by an
electron as it moves from the n = 6 to the n = 2 energy
level in a hydrogen atom.
(A) 0.38 eV (B) 3.02 eV (C) 3.40 eV (D) 13.60 eV
5.
A photon having energy of 9.4 eV strikes a hydrogen atom
in the ground state. Why is the photon not absorbed by
the hydrogen atom?
(A) The atom's orbital electron is moving too fast
(B) The photon striking the atom is moving too fast.
(C) The photon's energy is too small.
(D) The photon is being repelled by electrostatic force.
13. The first seven ionization energies of element X are shown
in the table below.
Ionization Energy (kJ•mol-1)
1st
2nd
3rd
4th
5th
6th
7th
787
1,580 3,200 4,400 16,000 20,000 24,000
On the basis of these data, element X is most likely a
member of which of the following groups of elements?
(A) Alkaline earth metals (B) Boron group
(C) Carbon group
(D) Nitrogen group
14. In which of the following are the chemical species correctly
ordered from smallest radius to largest radius?
(A) B < C < N
(B) At < Xe < Kr
(C) CI < S < S2(D) Na < Na+ < K
25. Which of the following molecules contains bonds that have
a bond order of 1.5?
(A) N2
(B) O3
(C) NH3
(D) CO2
26.
15. Of the following elements, which would be expected to
have chemical properties most similar to those of sulfur, S?
(A) Br
(B) CI
(C) P
(D) Se
16. Which pair of ions should have the highest lattice energy?
(A) Na+ and Br(B) Li+ and F(C) Cs+ and F(D) Li+ and O2-
17. Which molecule has the weakest bond?
(A) CO
(B) O2
(C) Cl2
(D) N2
18. Which pair of atoms should form the most polar bond?
(A) F and B
(B) C and O
(C) F and O
(D) N and F
19. Which species has a valid non-octet Lewis structure?
(A) GeCl4 (B) SiF4
(C) NH4+
(D) SeCl4
20. For which molecule are resonance structures necessary to
describe the bonding satisfactorily?
(A) H2S
(B) SO2
(C) CO2
(D) OF2
21. The Lewis structure for SeS2 with zero formal charge has a
total of
(A) 2 bonding pairs and 7 nonbonding pairs of electrons.
(B) 2 bonding pairs and 6 nonbonding pairs of electrons.
(C) 3 bonding pairs and 6 nonbonding pairs of electrons.
(D) 4 bonding pairs and 5 nonbonding pairs of electrons.
Questions 22-24 refer to the following molecules.
(A) CO
(B) CH4
(C) HF
(D) PH3
22. Contains two -bonds
23. Has the highest dipole moment (most polar)
24. Has a molecular geometry that is trigonal pyramidal
CCl4, CO2, PCl3, PCl5, SF6
Which does NOT describe any of the molecules above?
(A) Linear
(B) Octahedral
(C) Square planar
(D) Tetrahedral
27. According to the VSEPR model, the progressive decrease
in the bond angles in the series of molecules CH4, NH3,
and H2O is best accounted for by the
(A) increasing strength of the bonds
(B) decreasing size of the central atom
(C) increasing electronegativity of the central atom
(D) increasing number of unshared pairs of electrons
28. Which of the following is a formula for an acid?
(A) CH3–CO–CH3
(B) CH3–CH2–COOH
(C) CH3–CH2–CH2OH
(D) CH3–CH2–O–CH3
29. Which is NOT a structural isomer of 2-methylpentane?
(A) hexane
(B) 3-methylpentane
(C) 2,2-dimethylbutane (D) 4-methylpentane
30. Types of hybridization exhibited by carbon atoms in a
molecule of propyne, CH3CCH, include which of the
following?
I. sp
II. sp2
III. sp3
(A) I only
(B) II only (C) III only (D) I and III only
31. Which of the following best explains why the normal boiling
point of CCI4(I) (350 K) is higher than the normal boiling
point of CF4(I) (145 K)?
(A) The C-CI bonds in CCI4 are less polar than the C-F
bonds in CF4.
(B) The C-CI bonds in CCI4 are weaker than the C-F
bonds in CF4.
(C) The mass of the CCI4 molecule is greater than that of
the CF4 molecule.
(D) The electron cloud of the CCI4 molecule is more
polarizable than that of the CF4 molecule.
32. Which of the following substances exhibits significant
hydrogen bonding in the liquid state?
(A) CH2F2
(B) N2H4
(C) CH3OCH3
(D) C2H4
Questions 33-34 refer to a various points in time during an
experiment conducted at 1.0 atm. Heat is added at a
constant rate to a sample of a pure substance that is solid
at time to. The graph below shows the temperature of the
sample as a function of time.
39. At which of the following temperatures and pressures
would a real gas be most likely to deviate from ideal
behavior?
Temperature (K) Pressure (atm)
(A)
100
50
(B)
200
5
(C)
300
0.01
(D)
500
0.01
40. In which process are covalent bonds broken?
(A) Solid silver melts.
(B) Solid potassium chloride melts.
(C) Solid carbon (graphite) sublimes.
(D) Solid iodine sublimes.
(A) t1
(B) t2
(C) t3
(D) t5
33. Time when the average distance between particles is
greatest
34. Time when the temperature of the substance is between its
melting point and its boiling point
35. Heat energy is added slowly to a pure solid covalent
compound at its melting point. About half of the solid melts
to become a liquid. Which of the following must be true
about this process?
(A) Covalent bonds are broken as the solid melts.
(B) The temperature of the solid/liquid mixture remains the
same while heat is being added.
(C) The volume of the compound increases as the solid
melts to become a liquid.
(D) The average kinetic energy of the molecules becomes
greater as the molecules leave the solid state and
enter the liquid state.
36. Of the following gases, which has the greatest average
molecular speed at 298 K?
(A) Cl2
(B) NO
(C) H2S
(D) HCN
37. At approximately what temperature will 40. g of argon gas
at 2.0 atm occupy a volume of 22.4 L?
(A) 600 K (B) 550 K (C) 270 K (D) 140 K
38. Three gases in the amounts shown in the table are added
to a previously evacuated rigid tank.
Gas
Ar
CH4
N2
Amount
0.35 mol
0.90 mol
0.25 mol
If the total pressure in the tank is 3.0 atm at 25oC, the
partial pressure of N2(g) in the tank is closest to
(A) 0.75 atm
(B) 0.50 atm
(C) 0.33 atm
(D) 0.25 atm
41. A closed rigid container contains distilled water and N2(g)
at equilibrium. Actions that would increase the
concentration of N2(g) in water include which of the
following?
I. Shaking the container vigorously
II. Raising the temperature of the water
III. Injecting more N2(g) into the container
(A) I only
(B) II only (C) III only (D) I and II only
42. What is the mole fraction of ethanol in a 6 molal aqueous
solution?
(A) 0.006 (B) 0.1
(C) 0.08
(D) 0.2
43. What additional information is needed to determine the
molality of a 1.0-M glucose (C6H12O6) solution?
(A) Volume of the solution
(B) Temperature of the solution
(C) Solubility of glucose in water
(D) Density of the solution
44. A solution of toluene (MM = 90 g) in benzene (MM = 80 g)
is prepared. The mole fraction of toluene in the solution is
0.2. What is the molality of the solution?
(A) 0.2
(B) 0.5
(C) 2
(D) 3
45. Which of the following aqueous solutions has the highest
boiling point at 1.0 atm?
(A) 0.20 M CaCl2
(B) 0.25 M Na2SO4
(C) 0.30 M NaCl
(D) 0.40 M C6H12O6
46.
_CH3OCH3(g) + _O2(g)  _CO2(g) + _ H2O(g)
When the equation above is balanced using the lowest
whole-number coefficients, the coefficient for O2(g) is
(A) 6
(B) 4
(C) 3
(D) 2
47. What mass of KBr (MM = 119 g•mol-1) is required to make
250. mL of a 0.400 M KBr solution?
(A) 0.595 g (B) 1.19 g (C) 2.50 g (D) 11.9 g
48. Na2CO3(s) + 2 HCl(aq)  2 NaCl(aq) + CO2(g) + H2O(l)
In a laboratory, a student wants to quantitatively collect the
CO2(g) generated by adding Na2CO3(s) to 2.5 M HCI(aq).
The student sets up the apparatus to collect the CO2 gas
over water. The volume of collected gas is much less than
the expected volume because CO2 gas
(A) is soluble in water
(B) is produced at a low pressure
(C) is more dense than water vapor
(D) has a larger molar mass than that of N2 gas, the major
component of air
49. Which of the following would produce the LEAST mass of
CO2 if completely burned in excess oxygen gas?
(A) 10.0 g CH4
(B) 10.0 g CH3OH
(C) 10.0 g C2H4
(D) 10.0 g C2H6
50. A solution of RbCl (MM = 121 g•mol-1) contains 11.0 %
RbCl by mass. From the following list, what is needed to
determine the molarity of RbCl in the solution?
I. Mass of the sample
II. Volume of the sample
III. Temperature of the sample
(A) I only
(B) II only (C) III only (D) I and II only
51.
CS2(l) + 3 O2(g)  CO2(g) + 2 SO2(g)
When 0.60 mol of CS2(l) reacts as completely as possible
with 1.5 mol of O2(g) according to the equation above, the
total number of moles of reaction products is
(A) 2.4 mol (B) 2.1 mol (C) 1.8 mol (D) 1.5 mol
52. What is the empirical formula of a hydrocarbon that is 10 %
hydrogen by mass?
(A) CH3
(B) C2H5
(C) C3H4
(D) C4H9
53. By mixing only 0.15 M HCl and 0.25 M HCl, it is possible to
create all of the following solutions EXCEPT
(A) 0.21 M (B) 0.18 M (C) 0.16 M (D) 0.14 M
54.
8 H2(g) + S8(s)  8 H2S(g)
When 25.6 g of S8(s) (MM = 256 g•mol-1) reacts completely
with an excess of H2(g) according to the equation above,
the volume of H2S(g), measured at 0oC and 1.00 atm,
produced is closest to
(A) 30 L
(B) 20 L
(C) 10 L
(D) 5 L
55.
2 N2H4 + N2O4  3 N2 + 4 H2O
What mass of N2 can be produced when 8.0 g of N2H4
(MM = 32 g) and 9.2 g of N2O4 (MM = 92 g) react?
(A) 8.4 g
(B) 12.6 g (C) 7.8 g
(D) 10.5 g
56. A student weighs out 0.0154 mol of pure, dry NaCI in order
to prepare a 0.154 M NaCI solution. Of the following
pieces of laboratory equipment, which would be most
essential for preparing the solution?
(A) 50 mL volumetric pipet
(B) 100 mL Erlenmeyer flask
(C) 100 mL graduated beaker
(D) 100 mL volumetric flask
Questions 57-60 The figures show portions of a buret used in a
titration 0.0464 moles of monoprotic acid with a solution of
Ba(OH)2. Figures I and 2 show the level of the Ba(OH)2
solution at the start and at the endpoint of the titration,
respectively. Phenolphthalein was used as the indicator for
the titration.
Figure 1
Figure 2
57. What is the evidence that the endpoint of the titration has
been reached?
(A) The color of the solution in the buret changes from
pink to colorless.
(B) The color of the solution in the buret changes from
blue to red.
(C) The color of the contents of the flask changes from
colorless to pink.
(D) The color of the contents of the flask changes from
blue to red
58. The volume of Ba(OH), used to neutralize the acid was
closest to
(A) 22.80 mL
(B) 23.02 mL
(C) 23.20 mL
(D) 29.80 mL
59. The concentration of the Ba(OH)2 solution is closest to
(A) 1 M
(B) 2 M
(C) 3 M
(D) 4 M
60. What could explain why the student calculated a
concentration of Ba(OH)2 that was too large?
(A) An extra drop of phenolphthalein was added.
(B) A small amount of the acid was not transferred to the
titration flask.
(C) A drop of Ba(OH)2 remained attached to the buret tip.
(D) Rinsing the buret with distilled water just before filling
it with the Ba(OH)2 to be titrated.
66. The purpose of weighing the cup and its contents again at
CaCl2(s)  Ca2+ + 2 Clthe end of the experiment was to
For the process of solid calcium chloride dissolving in water,
(A) determine the mass of solute that was added.
represented above, the entropy change might be expected to
(B) determine the mass of the thermometer.
be positive. However, S for the process is actually negative.
(C) determine the mass of water that evaporated.
Which best helps to account for the net loss of entropy?
(D) verify the mass of water that was cooled.
(A) Cl- ions are much larger in size than Ca2+ ions.
(B) The particles in solid calcium chloride are more
ordered than are particles in amorphous solids.
(C) Water molecules in the hydrated Ca2+ and Cl- ions are
more ordered than they are in the pure water.
67. Suppose that during the experiment, a significant amount
(D) The Ca2+ and Cl- ions are more free to move around in
of solution spilled from the polystyrene cup before all of the
solution than they are in CaCl2(s)
solute dissolved. How does this affect the calculated value
for the heat of solution of the ionic compound?
(A) The calculated value is too large because less water
was cooled as the remaining solute dissolved.
(B) The calculated value is too large because some solute
62. For which of the processes does entropy decrease (S < 0)?
was lost with the spilled solution.
(A) H2O(s)  H2O(l)
(C) The calculated value is too small because less solute
(B) Br2(l)  Br2(g)
was dissolved than the student assumed.
(C) Crystallization of I2(s) from an ethanol solution
(D) The calculated value is too small because the total
(D) Thermal expansion of a balloon filled with CO2(g)
mass of the calorimeter contents was too small.
61.
63. What mass of Cu(s) would be produced if 0.40 mol of
Cu2O(s) was reduced completely with excess H2(g)?
(A) 13 g
(B) 25 g
(C) 38 g
(D) 51 g
64. A certain reaction is spontaneous at temperatures below
400 K but is not spontaneous at temperatures above 400
K. If Ho for the reaction is -20 kJ•mol-1 and it is assumed
that Ho and So do not change appreciably with
temperature, then the value of So for the reaction is
(A) -50 J•mol-1•K-1
(B) -20.0 J•mol-1•K-1
-1
-1
(C) -0.05 J•mol •K
(D) -20 J•mol-1•K-1
65.
68.
Pb(s)  Pb(l)
Which of the following is true for the process represented
above at 327oC and 1 atm? (The normal melting point for
Pb(s) is 327o(C))
(A) H = 0
(B) TS = 0
(C) S < 0
(D) H = TS
69.
C(diamond)  C(graphite)
For the reaction represented above, the standard Gibbs
free energy change, Go298, has a value of -2.90 kJ•mol-1.
Which of the following best accounts for the observation
that the reaction does NOT occur (i.e. diamond is stable)
at 298 K and 1.00 atm?
(A) So for the reaction is positive.
(B) The activation energy, Ea, for the reaction is very
large.
(C) The reaction is slightly exothermic (Ho < 0).
(D) Diamond has a density greater than that of graphite.
ZX+Y
A pure substance Z decomposes into two products, X and
Y, as shown by the equation. Which of the following graphs
of the concentration of Z versus time is consistent with the
rate of the reaction being first order with respect to Z?
(A)
(B)
(C)
(D)
Questions 66-67 refer to an experiment to determine the heat
of solution of an ionic solid. A student used a calorimeter
consisting of a polystyrene cup and a thermometer. The
cup was weighed, then filled halfway with water, then
weighed again. The temperature of the water was
measured, and some of the ionic solid was added to the
cup. The mixture was gently stirred until all of the solute
dissolved and the lowest temperature reached by the water
in the cup was recorded. The cup and its contents were
weighed again.
70. When a solution is formed by adding some methanol,
CH3OH, to water, processes that are endothermic include
which of the following?
I. Methanol molecules move water molecules apart
as the methanol goes into solution.
II. Water molecules move methanol molecules apart
as the methanol goes into solution.
III. Intermolecular attractions form between
molecules of water and methanol as the methanol
goes into solution.
(A) I only
(B) II only (C) III only (D) I and II only
Answers
#
1
2
3

Explanation
Average = (Mass1 x Abundance1) + (Mass2 x Abundance2)
D
63.5 = (63)(1 – x) + (65)(x) = 63 – 63x + 65 x  x = 0.25 (25 %)
D 3920Ca  3919K + 01
It takes 3 half-lives to reduce the radioactivity to 1/8 (1.25/10.0).
C
195 days/3 = 65 days
4
B
5
C
6
C
7
B
8
A
9
C
10 B
11 C
12 A
13 C
14 C
15 D
16 D
17 C
18 A
19 D
20 B
21 D
22 A
23 C
24 D
25 B
26 C
27 D
28 B
29 D
30 D
31 D
32 B
33 D
34 C
35 B
36 D
37 B
38 B
39 A
40 C
From the diagram: E6 = -0.38 eV and E2 = -3.40 eV
E = E2 – E6 = -3.40 eV – (-0.38 eV) = -3.02 eV
The electron can only absorb energy that will move it to a higher
energy level. 9.4 eV is not enough energy (the minimum needed
is -3.40 eV – (-13.60 eV) = 10.2 eV).
The orbital diagram for C,1s() 2s() 2p()()( ), has two
unpaired electrons (Hund's rule) = paramagnetic.
Pauli states that no orbital can contain electrons with the same
spin. Since two spins, this limits the number to two electrons.
Heisenberg states that the wave nature of matter (DeBroglie)
limits what we can know about position and velocity.
Electron # 31 is located in the 4th row (n = 4), 13th column (p
section, l= 1), which limits ml = 1, 0 or -1 and ms = +½ or -½ 
(4, 1, 1, ½) fits requirement
1s() 2s() 2p()(
): The 2p electron is in an excited
state, otherwise it would go into the 2s sublevel.
1s() 2s() 2p()()(): The 2 2s electrons and 3 2p
electrons are in the valence shell (highest energy level)  five.
1s(): The first ionized electron is from the 1s sublevel. It takes
the most energy to remove electrons that are close to nucleus.
The biggest jump in ionization energy occurs between 4 and 5,
which means 4 valence electrons  Carbon group.
Atomic radius increases going left and down in periodic table.
Anion is larger than atom and cation is smaller than atom.
Elements in the same column in the periodic table have similar
chemical properties.
Lattice energy is a measure of ionic bond strength, which is
proportional to charge and inversely proportional to size.
Single bonds are the weakest (CO, O=O, Cl–Cl, NN)  Cl2.
Most polar bond forms between atoms with the greatest
electronegativity difference (greatest gap on the periodic table).
SeCl4 has 34 valence electrons, which require an expanded octet
system (sp3d) to accommodate all the electrons.
Only SO2 has a single and double bond, which can exchange
places, thus forming resonance forms. (H2S and OF2: 2 single
bonds, CO2: 2 double bonds)
::S=Se:=S:: has zero formal charge. There are 4 bonds and 5
pairs of nonbonding electrons.
CO: The triple bond between C and O is composed of one
sigma bond and two pi bonds.
H-F: The electronegativity difference is greatest between H and F
 the most polar, which produces the highest dipole moment.
PH3: The three H are pushed away from the pair of non-bonding
electrons around phosphorus resulting in a pyramidal structure.
A bond order of 1.5 means 1 sigma bond and 50% share of a pi
bond, which is the case for O3 (O=O–O).
CCl4 (tetrahedron), CO2 (linear), PCl3 (trigonal pyramid), PCl5
(trigonal bipyramid), SF6 (octahedron)
The non-bonding electron pairs take up more space than bonding
pairs  H2O (2 pairs) < NH3 (1 pair) < CH4 (0 pair).
Acids contain the COOH functional group  (B) (a is a ketone, c
is an alcohol, and d is an ether)
4-methylpentane is the same as 2-methylpentane because you
number from the closest end  4 becomes 2.
C1H3–C2C3H: C1 is sp3, C2 is sp, C3 is sp
Both molecules are non-polar, but the larger molar mass of CCl4
means that there are more electrons, which are more polarizable
 generating a stronger dispersion force.
Hydrogen bonding occurs when H is bonded to N, O or F. Only
N2H4 has that arrangement. (CH2F2: H is not bonding to the F)
Farthest from each other in the gaseous phase, which is at t 5.
Melting occurs along 1st plateau (t2) and boiling along 2nd plateau
(t4)  the time between these two temperatures is t3.
(a) Covalent bonds aren't broken if its molecular.
(c) Volume only increases if liquid state is less dense.
(d) Temperature doesn't increase  KE is not greater.
At the same temperature, lighter molecules have greater speed.
40 g of Ar = 1 mol. One mole at STP = 22.4 L. Since P is 2 x
standard, then T has to be 2 x standard (V = nRT/P)  2 x 273 K
PN2 = XN2Ptot = (0.25/(0.35 + 0.90 + 0.25))(3.0 atm) = 0.50 atm
Real gases deviated from ideal behavior at low temperatures
(near their boiling point) and high pressure.
(A) metallic bond, (B) ionic bond, (C) covalent bond, and (D)
molecular bond.
41 C
42 B
43 D
44 D
45 B
46 C
47 D
48 A
49 B
50 D
51 D
52 C
53 D
54 B
55 A
56 D
57 C
58 C
59 A
60 B
61 C
62 C
63 D
64 A
65 D
66 A
67 A
68 D
69 B
70 D
Gas solubility increases with greater partial pressure of the gas in
the container or lower temperature.
6 molal = 6 mol ethanol in 1000 g H2O
1000 g H2O x 1 mol/18 g = 55 mol H2O
 mole fraction = 6/(6 + 55) = 0.1
Molality is moles solute/kg solvent. Molarity is moles solute/L
solution. To determine kg solvent from L solution, you need to
know density; then subtract g solute.
0.2 mole fraction = 0.2 mol toluene in 0.8 benzene
0.8 mole x 80 g/1 mol = 60 g benzene
molality = 0.2 mol/0.060 kg = 3 m
The highest boiling point = highest concentration of ions.
(A) .2 x 3 = .6, (B) .25 x 3 = .75, (C) .3 x 2 = .6, (D) .4 x 1 = .4
1 CH3OCH3(g) + 3 O2(g)  2 CO2(g) + 3 H2O(g)
0.250 L x 0.400 mol/L = 0.1 mol KBr x 119 g/mol = 11.9 g
If CO2 is soluble in water, then some gas would remain in the
water and not bubble into the gas collecting bottle.
CH3OH has the lowest proportion of C (12/32)  given equal
masses; CH3OH would generate the least amount of CO2.
molarity = mol solute/volume solution (L). The mass is needed to
determine the number of moles of solute. The volume is needed
to determine the volume of solution.
0.6 mol CS2 x 3 mol Products/1 mol CS2 = 1.8 mol Products
1.5 mol O2 x 3 mol Products/3 mol O2 = 1.5 mol Products
1.5 mol Products is the lesser number.
Assume 100 g of compound
10 g H x 1 mol/1 g = 10 mol/7.5 = 1.33 x 3 = 4
90 g C x 1 mol/12 g = 7.5 mol/7.5 = 1 x 3 = 3
The resulting solution must have a concentration between the two
solutions added together. 0.14 M is less then both.
25.6 g S8 x 256 g/mol x 8 mol H2S/1 mol S8 = 0.8 mol
0.8 mol x 22.4 L/mol = 20 L
0.25 mol N2H4 x 3 mol N2 x 28 g N2 = 10.5 g
2 mol N2H4 1 mol N2
0.1 mol N2O4 x 3 mol N2/1 mol N2O4 x 28 g N2 = 8.4 g
0.0154 mol NaCl x 1 L/0.154 mol = 0.1 L (100 mL)
The most accurate way to measure 0.1 L of solution is to use a
volumetric flask.
Phenolphthalein changes from clear (acid) to pink (base)
Final volume – Initial volume = Change in volume
35.75 mL – 12.55 mL = 23.20 mL
0.0464 mol H+ x 1 mol OH- x 1 mol B(A).. = 0.0232 mol B(A)..
1 mol H+ 2 mol OH0.0232 mol B(A)../0.0232 L = 1 M
MBa(OH)2 = (½ mol H+)/VBa(OH)2
 too large MBa(OH)2 = too small VBa(OH)2
Titrating less acid would result in smaller VNaOH
Dissolving involves two process; (1) separation into ions, which
increases disorder (+S) and (2) ions combining with water
(solvation), which decreases disorder (-S).
Disorder decreases when I2(aq)  I2(s). Disorder increases
when s  l  g (A) and (B). Disorder also increases when gas
molecules spread out (D).
Cu2O + H2  2 Cu + H2O (balancing wasn't necessary because
mole Cu in reactants and products are equal)
0.80 mol Cu x 63.5 g/mol = 51 g
Tthreshold = H/S
S = H/Tthreshold = -20 kJ•mol-1/400 = -0.05 kJ•mol-1•K-1
-0.05 kJ•mol-1•K-1 x 1000 J/1 kJ = -50 J•mol-1•K-1
For a first order reaction, the straight line graph is ln[Z] vs. t.
(zero order is [Z] vs. t and second order is 1/[Z] vs. t)
At the end of the experiment, the cup contained solute and water.
If this value is subtracted from the mass of the cup and water,
then the difference is the mass of the solute.
The heat needed to dissolve the remaining solute had to come
from less water, which would make T greater than it should have
been (H = -mcT) H would be too large.
At the normal melting point: G = 0 = H – TS H = TS
The reaction is spontaneous, but it must not occur at a fast rate.
This could be because the activation energy is so high, that it
takes too much energy to start the process.
Breaking solute-solute bonds in methanol and water is
endothermic, but forming solute-solute bonds between methanol
and water is exothermic.