Chemistry 4: States of Matter
A.
Gas State (10.2 to 10.9)
1. tend to be small, non-polar molecules
2. form homogenous mixtures
3. distributed throughout the entire container
(molecules typically occupy 0.1 % of the volume)
4. kinetic theory for gases (ideal gas)
a. molecules in continuous, chaotic motion, which is
proportional to temperature
1. kinetic energy, Kmole = 1/2(MM)urms2 = 3/2RT
a. R = 8.31 J/mol•K
b. MM (molar mass) in kg
c. T in kelvin (TK = ToC + 273)
d. average speed, urms = (3RT/MM)½
2. Graham’s law
a. rate r of effusion (leaking) or diffusion
(spreading out) is proportional to speed
b. rA/rB = (MMB/MMA)½
c. time is inversely proportional to rate 
TB/TA = rA/rB
b. molecular volume is insignificant compared to
container volume (approximation—see real gas)
c. collisions produce pressure w/o loss of total
kinetic energy
d. Bonding between molecules is insignificant
(approximation—see real gas)
5. gas laws
a. Ideal gas Law equation: PV = nRT
1. molecules generate pressure via collisions
a. pressure = force/area
b. 1 atm = 101 kPa = 760 mm Hg (torr)
c. measuring tools
1. barometer: atmospheric pressure
2. manometer: enclosed gas pressure
Pgas = Patm ± h (in mm Hg)
2. gas pressure is affected by:
a. n: each molecule exerts pressure  more
molecules exert more pressure: P  n
b. T: hotter molecules move faster and
collide with greater force  generate
more pressure: P  T
c. V: molecules spread out which reduces
collision per surface area  generate
less pressure: P  1/V
3. ideal gas law constant R (V in L, T in K)
a. 8.31 J/mol•K (P in kPa)
b. 0.0821 atm•L/mol•K (P in atm)
4. molar volume at STP = 22.4 L/mol
(standard T = 0oC, standard P = 1 atm)
5. derived equations
a. P1V1/T1 = P2V2/T2
b. MM = mRT/PV = dRT/P
b. Dalton’s law (P  n)
1. Ptot = PA + PB
2. PA = XAPtot , where XA = molA/(molA + molB)
6. real gases
a. Van der Waals:(Preal + n2a/V2)(Vreal – nb) = nRT
b. "a" corrects for molecular bonding
1. "low" temperature (close to boiling point)
molecules clump and collide less often, which
generates less pressure  Preal < Pideal
2. a is proportional to molecular polarity
c. "b" corrects for molecular volume
1. high pressure is generated by crowded
molecules where the volume of empty space
(Videal) is significantly less than 100 % of the
total volume (Vreal) Vreal > Videal
2. b is proportional to molar mass
Name __________________________
B.
Phase Change (11.1 to 11.2, 11.4 to 11.6)
1. molecular-level comparison of gas, liquid, and solid
states of a substance
Characteristic
Gas
Liquid
Solid
Energy
Highest
Middle
Lowest
Disorder
Greatest
Middle
Least
Occupied space
Whole
Bottom
Own
Compressibility
Yes
No
No
2. cohesive forces (van der Waals forces)
a. attraction between molecules
(covalent bonds hold atoms together in molecule)
b. dipole-dipole forces
1. polar molecules
2.  of one molecule attracts  of a neighbor
3. strength  to polarity, if all else is even
c. London dispersion forces
1. attraction between nuclei of one molecule's
atoms for the electrons in a neighboring
molecule causes temporary polarization
throughout the liquid or solid (polarizability)
2. generalization
a. operates between all molecules (stronger
than dipole-dipole for large molecules, i.e.
large nonpolar > small polar)
b. only force for nonpolar (strength  to
mass: Xe > He, I2 > F2, C3H8 > CH4)
d. hydrogen bonding
1. super strong dipole-dipole force
2. H bonded to N, O or F
a. H  +1 charge and N, O or F  –1 charge
because of extreme electronegative
difference and small radius
b. bonding is ionic like (E  Q1Q2/d)
3. explains unusual properties of water
a. each water molecule bonds to 4, which
makes a 3-d structure with open cavities
b. high melting and boiling temperatures
c. low vapor pressure (low volatility)
3. cooling profile for water from 110oC to -10oC
A
gas
B
C
o
100 C
condensing
boiling
liquid
0 oC
a.
b.
c.
D
E
solid
freezing
F
melting
Q (Heat Removed in J)
slope C-D < slope E-F  more heat is
removed when one mole H2O(l) is cooled 1oC
compared to one mole H2O(s)
length B-C > length D-E  more heat is
removed when one mole H2O(g)  H2O(l)
compared to one mole H2O(l)  H2O(s)
calculations C  D
1. Q = nClT = mclT
2. molar heat capacity, Cl = 75.3 J/mol•K(
3. specific heat, cl = 4.18 J/g•K
4.
phase diagram
a.
5.
point A: triple point (three phases at equilibrium)
1. above triple point: melting and vaporization
2. below triple point: sublimation (deposition)
b. line A-C: equilibrium vapor pressure curve for liquid
(nbp: normal boiling point occurs at 1 atm pressure)
c. C: critical point, where there is no distinction
between liquid and vapor (no liquid-vapor surface)
d. line A-D: equilibrium vapor-pressure curve for solid
(sp: sublimation point depends on pressure)
e. line A-B: melting point of solid at various pressures
(mp: normal melting point occurs at 1 atm)
1. negative slope when liquid is the densest
phase (melting point decrease with pressure)
2. positive slope (more common) when solid is
the densest phase (melting point increases
with pressure)
vapor
a. some surface molecules in the condensed phase
have enough kinetic energy (speed) to escape
surface (evaporate) below boiling point
b. as temperature increases more molecules have
sufficient kinetic energy  more vapor molecules
c.
d.
e.
f.
cooling process (hottest evaporate first, leaving
cooler molecules behind)
equilibrium between liquid and vapor
1. evaporation rate = condensation rate in a
closed container
2. concentration of vapor measured as Pvap
3. independent of container size until no liquid
4. Pvap increases at higher temperature because
a. more molecules are in vapor phase
b. vapor molecules exert greater pressure
boiling occurs when Pvap = Patm  boiling point
decreases with elevation (lower air pressure, Patm)
high Pvap indicates volatility—tendency to evaporate
C.
Crystalline Solids (11.8)
1. ions, atoms or molecules fit into a regular geometric
pattern (crystal lattice)
2. minimum energy state—maximum bond energy
3. intermolecular forces (attraction between + and -) or
bonds (covalent, ionic or metallic) hold particles together
4. 4 types of solids
a. metallic—metals only
1. attraction between cations and delocalized
valence electrons (electron sea model)
b.
c.
d.
2. melting point: variable ( bond strength)
3. conductivity: free electrons  yes
4. malleable: non-directional bond  yes
5. water solubility: no molecular interactions  no
6. examples: Cu, Ag, etc.
covalent network—nonmetals w/o H or halogen
1. atoms covalently bond throughout w/o size
limit (different than large molecule)
2. melting point: strong bonds  high
3. conductivity: no free electrons  no
4. malleability: bond highly directional  brittle
5. water solubility: no molecular interactions  no
6. three examples
a. diamond and graphite—C
1. allotropes (2 forms in the same state)
2. diamond: covalently bonded 3-d
structure—good abrasive
3. graphite: covalent bonded planar
sheets linked by dispersion forces
a. separate easily—good lubricant
b. electron flow—good conductivity
b. quartz—SiO2
1. 3-d structure similar to diamond
2. softens when heated until liquid
3. fast cooling = non-crystalline glass
molecular—nonmetals often with H and/or halogen
1. attraction between + of one molecule with
– of another
2. melting point: weak bonds  low
3. conductivity: no free electrons  no
4. malleability: non-directional  yes
(H-bonding in H2O(s) is somewhat directional)
5. water solubility ("like dissolves like")  yes/no
6. examples: H2O, C6H12O6, etc.
Ionic—metal plus nonmetals
1. attraction between cations and anions
2. melting point: strong bonds 
3. conductivity: no free electrons  no
(fused or dissolved state is conducting)
4. malleability: bond highly directional  brittle
5. water solubility: ion-dipole interaction  yes
6. examples: NaCl, CaCO3, etc.
D.
Solubility (13.1 to 13.4)
1. dissolution
a. one substance disperses uniformly throughout other
1. solvent: dissolving medium (usually majority)
2. solute: dissolved in a solvent (usually minority)
a. ionic or acid = electrolyte (forms ions)
b. number of free ions = i (van't Hoff factor)
c. polar molecules = nonelectrolyte
3. solvation: attraction between solute-solvent
(hydration if solvent is water)
a. cation with  side of H2O (O side)
b. anion with  side of H2O (H side)
b. saturated solution
1. undissolved solid  dissolved solid
2. dissolution rate = crystallization rate
3. maximum amount that dissolves = solubility
(liquids that mix in all proportions = miscible)
4. solubility graphs (g solute/100 g H2O)
c. effect of temperature on solubility
1. solvent kinetic energy is used to break solutesolute bonds  solute gains energy; solvent
lose kinetic energy (cools)
2. energy is released when solute-solvent bonds
form and turns into kinetic energy of solution
particles (warms up)
3. H = Esolute-solute – Esolute-solvent
a. when Esolute-solute > Esolute-solvent
1. +H (solution cools = endothermic)
2. raising T increases solubility
b. when Esolute-solute < Esolute-solvent
1. –H (solution warms = exothermic)
2. raising T decreases solubility
5. gas solubility generally decreases with
increased temperature because solution
depends on solute-solvent bonds, which
weaken as temperature increases
d. effect of pressure upon solubility (gas only)
1. solubility increases proportionally to partial
pressure above solution Mg = kPg (Mg = mol/L)
2. gas in solution  gas in air space  more
gas in air space force more into solution
3. only Pg, not Ptot, will increase solubility
2. separation solute and solvent
a. filtration: separate solvent from insoluble solute
b. distillation
1. simple: separate solvent from soluble solid
2. fractional: separate solvent from soluble liquid
3. expressing concentration
a. chemists use molarity: M = molsolute/Vsolution-L
b. making a molar solution from stock

moles needed: molestandard = MstandardVstandard
o mass of stock powder, m = (molestandard)MM
o volume of stock solution, V = (molestandard)/(Mstock)
(Mstock)(Vstock) = (Mstandard)(Vstandard)

add to volumetric flask filled ¾ full with distilled water

dissolve

add sufficient distilled water to bring volume to total
c. other concentration units
1. mole fraction: Xsolute = molsolute/moltotal
2. molality: m = molsolute/msolvent(kg)
d. conversion of concentration units

assume 1 unit of denominator amount of solution
(X: 1 mol total, M: 1 L solution, m: 1 kg solvent)

determine how numerator and denominator change

convert numerator and/or denominator
o mass  volume: d = m/V
o mass  moles: n = m/MM
E.
Colligative Properties (13.5)
1. lower vapor pressure
a. solute particles reduce vapor pressure by
reducing the number of solvent particles on the
surface that can evaporate
nonvolatile-nonelectrolyte: Pvap = XsolventPosolvent
1. known as Raoult's law
2. "ideal" solution
c. two volatile liquids: Pvap = XAPoA + XBPoB
higher boiling point and lower freezing point
a. lowered vapor pressure changes melting and
boiling temperatures (extends liquid phase)
b.
2.
b.



3.
Tb = Kbmi, Tf = Kfmi
1. m = molality
2. i  number of ions (van't Hoff factor)
3. Kb/Kf = molal boiling/freezing point constant
c. determination of molar mass of non-electrolyte
solute by freezing pt. depression
calculate molality: molality = Tf/Kf
calculate molsolute: molsolute = (molality)(msolvent/1000)
calculate molar mass: MM = msolute/molsolute
osmotic pressure
a. semi-permeable membrane blocks solute
b. solvent flows from high [ ] to low [ ]  osmosis
c. osmotic pressure () = pressure to stop flow
d.  = MRTi
(R = 8.31 when  in kPa or 0.0821 when  in atm)
3.
Experiments
1.
Molar Mass of a Gas Lab—Measure the mass and volume
of butane released from a lighter, determine the molar
mass and compare it to the known molar mass.
Mass a butane lighter (m1). Fill a 50 mL graduated cylinder
with water and place it upside down in a filled trough.
Release butane into the graduated cylinder until the water
levels in the cylinder and trough are the same. Record volume
(V). Dry the lighter thoroughly and mass (m2). Record
temperature (T), pressure (Plab) and vapor pressure (PH2O).
a. Record the collected data.
m1
m2
V
T
Plab
PH2O
b.
Complete the following calculations to determine the
molar mass and percent difference from known value.
P = Plab – PH2O
T
V
m = m1 – m2
(atm)
(K)
(L)
(g)
MM = mRT/PV
c.
MM (C4H10)
%
How would the following affect the molar mass value?
(1) The butane lighter was not thoroughly dried.
Molar Mass of Solute Lab—Graph the data to determine
freezing point and use the data to calculate molar mass.
Part 1: 5.00 g of BHT is cooled while recording temperature.
Time (s) 0
20
40
60
80
100 120 140
74.0 72.2 69.8 68.0 66.8 67.8 69.0 69.0
ToC
Time (s) 160 180 200 220 240 260 280 300
68.8 68.8 69.0 68.8 69.0 67.0 65.2 63.8
ToC
Part 2: 0.500 g of naphthalene (C10H8) is added to 5.00 g
BHT. The mixture is cooled while recording temperature.
Time (s) 0
20
40
60
80
100 120 140
70.0 68.0 65.8 64.2 61.8 60.0 58.8 60.2
ToC
Time (s) 160 180 200 220 240 260 280 300
59.8 59.2 58.8 58.2 57.6 57.0 55.2 53.4
ToC
Part 3: 0.500 g of unknown is added to 5.00 g BHT. The
mixture is cooled while recording temperature.
Time (s) 0
20
40
60
80
100 120 140
72.0 70.0 68.2 65.8 64.0 62.8 64.0 63.4
ToC
Time (s) 160 180 200 220 240 260 280 300
63.0 62.8 62.2 61.8 61.4 61.0 59.0 57.2
ToC
a. Graph the data from parts 1, 2 and 3 using three different
colors. Draw a straight line following the cooling process
and a second straight line following the freezing process
for each graph. Record the intersection, Tf (freezing pt.).
Temperature (oC)
74
72
(2) The vapor pressure of water was not subtracted
from the room pressure.
70
68
(3) The water level in the inverted graduated cylinder
was higher than the level in the trough.
66
64
2.
Solute Concentration Lab (Wear Goggles)—Mass a
measured volume of solution and the solute that remains
after all the water is boiled away and use the mass values
to determine the solution concentration in different units.
Mass an empty, clean 125 mL flask (m1). Add 25.0 mL
solution to the flask and mass (m2). Place the flask on the
hot plate until all the water has boiled away and the rim of
the flask is dry. Mass the cooled flask plus solute (m3).
a. Record the masses in the chart below.
m1
m2
m3
62
60
58
56
54
20
b.
b.
Determine the following.
60
100
140 180 220
Time (s)
Complete the following chart.
Naphthalene/BHT Mixture
Tf (BHT)
Tf (Solution)
mNaCl = m3 – m1
mH2O = m2 – m3
n = m/MM
Volume of solution
Molality m
c.
Kf = Tf/m
Calculate the following.
Mole NaCl
Tf (BHT)
Mole H2O
mass %
m = Tf/Kf
mole fraction
n = mBHTm
molarity
MM = m/n
molality
c.
Unknown/BHT Mixture
Tf (Solution)
260
300
T f
T f
The unknown's molecular formula is CH3(CH2)14CH2OH.
Determine the percent error for this experiment.
Practice Problems
1.
A. Gas State
What features of the kinetic theory of gases
a. describe all gas molecules?
b.
2.
3.
4.
7.
A gas is confined inside a container with a movable piston
held down by a fixed pressure.
describe ideal gas molecules only?
a.
What affect would doubling the number of gas
molecules at the same temperature have on the
system? Explain.
b.
What affect would doubling the Kelvin temperature of
the gas have on the system? Explain.
c.
What affect would doubling pressure by the piston at
the same temperature have on the system? Explain.
Consider one mole of Ne gas at 274 K. Determine
a. the total kinetic energy.
b.
the average speed.
a.
What alkane effuses at 1/5 the rate of He?
b.
How many times faster does C2H2 diffuse compared to
the alkane?
Consider the graph below.
8.
Complete the following table for an ideal gas:
P
V
n
2.00 atm
1.00 L
30.3 kPa
1.250 L
650 torr
9.
b.
5.
6.
A and B are He and O2, at 25oC, which is which?
Explain
A and B are at 100 K and 200K, which is which?
Explain
Determine the pressure of 1.22 atm in the following units.
mm Hg
kPa
torr
If the atmospheric pressure is 749 mm Hg, what is the
pressure of the enclosed gas in each case below?
1.500 mol
27oC
0.333 mol
350 K
585 mL
0.250 mol
295 K
Oxygen gas in a 10.0-L container has a pressure of 94.6 kPa
and temperature of 25oC.
a. How many moles of oxygen gas are in the container?
b.
a.
T
How many grams of oxygen gas are in the container?
10. A sample of gas occupies 350 mL at 15oC and 750 torr.
What temperature will the gas have at the same pressure if
its volume increases to 450 mL?
11. Determine the molar mass of an unknown gas given the data.
Mass
Volume
Temperature
Pressure
4.93 g
1.00-L
400. K
1.05 atm
12. Calculate the density of ammonia, NH3, at STP.
13. Consider the samples of gases.
I
II
III
The samples are at the same temperature. Rank them
with respect to the following (1 is highest).
I
II
III
Total Pressure
Partial Pressure of He
Density
Average Kinetic energy per molecule
Total Kinetic energy
14. Each bulb contains a gas at the pressure and volume
shown and temperature of 25oC. Determine
20. For each pair, highlight the molecule with the higher boiling
point and then justify your choice.
Pair
Justification
H2O & H2S
Ne & Kr
Cl2 & SO2
21. Explain the boiling points for the two isomers.
a.
the number of moles of each gas.
N2
Ne
H2
b.
the total pressure after all stopcocks are opened.
c.
the partial pressure of each gas.
22. In addition to dispersion, what type of force would you
expect between the following molecules?
H2
H2S
CHF3
NH3
23. Consider the heating profile for water in your notes.
a. What can you conclude about the value of Cl compared
to Cs based on the slope of line C-D compared to E-F?
N2
Ne
b.
H2
15. 2.00 L of Hydrogen gas is collected over water at 30.0oC.
The total pressure is 740 torr (PH2O = 32 torr).
a. What is the partial pressure of the hydrogen gas?
b.
How many moles of hydrogen gas are collected?
16. A 20-L flask holds 0.20 mol O2 and 0.40 mol NO2 at 27oC.
a. What is the pressure of the mixture in kPa?
b.
What can you conclude about Hfus compared to Hvap
based on the length of line B-C compared to D-E?
24. Consider the phase diagram in your notes.
a. How does melting point change when pressure increases?
b.
How would the diagram differ for most substances?
25. Answer the questions based on the phase diagram.
What is the partial pressure of oxygen in kPa?
17. Which gas, SO2 or CO2, should be least ideal at STP?
Explain
18. a.
Why do gases under high pressure deviate from ideal
behavior?
Temperature (oC)
Can liquid CO2 exist at room pressure?
What happens to CO2(s) at -78.5oC?
b.
Why do gases at temperatures near their boiling point
deviate from ideal behavior?
Which is the most dense phase for CO2?
What is the triple point pressure for CO2?
B. Phase Change
19. Which letter illustrates the types of molecular forces?
What is the critical temperature for CO2?
26. Explain how pure water can boil at room temperature when
placed in an evacuated bell jar.
27. Explain why baking takes longer at high elevations.
Dipole-Dipole
Dispersion
H-bond
28. Answer the questions based on the vapor pressure curves.
a.
Which liquid has the strongest intermolecular bonding?
b.
Which liquid is the most volatile?
c.
What would happen if you slowly heated a beaker filled
with the three liquids from 20oC to 100oC?
29. 0.010 moles of water is added to a 5.0-L container filled
with dry air at 20oC (vapor pressure = 20 torr). The
container is then sealed and equilibrium is established.
a. How many moles of water evaporate?
b.
b.
Cl2 boils at 238 K and HCl boils at 188 K.
c.
KCl melts at is 776oC and NaCl melts at 801oC.
d.
Si melts at 1,410oC and Cl2 melts at -101oC.
D. Solubility
37. What is the approximate van't Hoff factor for the following?
Na2O
CaCl2
AlF3
C6H12O6
38. Indicate whether the solute is likely to dissolve in water?
NaCl
CH3OH
HC2H3O2
C20H42
What percentage of the water evaporates?
30. Explain why water droplets form on a cold water bottle.
C. Crystalline Solids
31. Use the electron-sea model of metals to explain
a. malleability.
b.
36. Explain the following observations. You must discuss both
of the substances in your explanation.
a. SO2 melts at 201 K and SiO2 melts at 1,883 K.
39. When KNO3 is added to water, the temperature of the
solution decreases. Highlight the correct option.
a. lattice energy is (greater/less) than hydration energy.
b. KNO3 is more soluble in (warm/cold) water.
40. Determine the missing value for the following solutes using
the solubility graphs below.
conductivity.
32. What are allotropes?
33. In what way is SiO2 like diamond; unlike diamond?
34. What two factors affect ionic bond strength (lattice energy)?
35. Complete the chart for each type of solid.
Covalent
Metallic
Molecular
Network
Structural Unit
Bond name
Bond strength
Melting point
Solubility
Conductivity
Malleability
Example
Ionic
Solute
mass solute
90 g
K2Cr2O7
NaCl
70 g
KClO3
15 g
mass water
100 g
Temperature
50oC
100 g
90oC
30oC
50 g
41. What is the concentration of CO2 (k = 3.1 x 10-2 mol/L•atm)
that is bottled with a partial pressure of 4.0 atm?
42. What is the concentration of N2 (k = 6.8 x 10-4 mol/L-atm) in a
diver's blood if he breaths air at 2.5 atm that is 78 % N2.
43. A solution is made up of 123 g NaOH and 289 g water.
The total volume is 300. mL. Determine
c.
What volume contains 157 g of Na2SO4?
50. Name the separation technique for the following.
Separate salt from water
Separate sand from water
Separate alcohol from water
E. Colligative Properties
51. 0.25 mol solute is added to 1.0 mol benzene (VP = 450 torr).
a. What is the mole fraction of benzene in the solution?
mole NaOH
mole H2O
b.
What is the vapor pressure of the solution?
mass %
mole fraction
52. What is the vapor pressure when PH2O = 2.4 kPa, when
a. 0.50 mol C6H12O6 is in 5.5 mol H2O.
molarity
b.
0.50 mol C2H5OH (PC2H5OH = 9 kPa) is in 5.5 mol H2O.
molality
44. Determine the density of 12.0 M HCl is 37.0 % HCl by mass.
a. What is the mass of HCl in one liter of solution?
53. 7.90 g of dichlorobenzene (C6H4Cl2) is added to 50.0 g of
benzene. (benzene: Kf = 5.12oC/m, Tf = 5.50oC)
a. How many moles of dichlorobenzene are in the solution?
b.
What is the mass of one liter of solution?
b.
What is the molality of the solution?
c.
What is the density of 12 M HCl in g/mL?
c.
What is the change in freezing point of the solution?
45. How would you prepare 250. mL of a 0.127 M Ca(OH)2
a. from powder Ca(OH)2?
d.
What is the freezing point of the solution?
b.
from 1.00 M Ca(OH)2?
46. You are asked to make 100. mL of a 0.125 M NaHCO3.
a. What mass of powder NaHCO3 would you need?
b.
b.
the moles of ethylene glycol are in the solution.
c.
the molecular mass of ethylene glycol.
What volume of 3.00 M NaHCO3 would you need?
47. a.
How many liters of 0.487 M NaOH is needed to make
0.100 L of a 0.200 M solution?
b.
What is the molarity of a solution when water is added
to 25.0 mL of 0.400 M HNO3 to make 75.0 mL?
48. What is the molarity of a solution that contains 73.2 g of
NH4NO3 in 0.835 L of solution?
49. Consider a 0.250 M solution of Na2SO4.
a. What volume contains 0.700 moles Na2SO4?
b.
54. 5.00 g of ethylene glycol in 100 mL of water (Kf = 1.86 oC/m)
freezes at –1.50oC. Determine
a. the molality of the solution.
How many grams of Na2SO4 are in 0.800 L of solution?
55. What is the freezing point of a solution made from 5.00 g of
glucose (C6H12O6) in 25 mL of water (Kf = 1.86 oC/m)?
56. 100 mL of solution contains 0.0020 mol solute at 25 oC.
a. What is the molarity of the solution?
b.
What is the osmotic pressure in kPa of the solution?
57. What is the concentration of solute particles in a solution with
an osmotic pressure of 73.4 atm and temperature of 25 oC?
58. How do the colligative properties change (, ) when nonvolatile solute is added to solvent?
Vapor P.
Freezing pt.
Boiling pt.
Osmotic P.
Practice Multiple Choice
Briefly explain why the answer is correct in the space provided.
Questions 1-2 The molecules have the normal boiling points.
Molecule
HF
HCl
HBr
HI
Boiling Point, oC
+19
-85
-67
-35
1. The relatively high boiling point of HF can be correctly
explained by which of the following?
(A) HF gas is more ideal.
(B) HF molecules have a smaller dipole moment.
(C) HF is much less soluble in water.
(D) HF molecules tend to form hydrogen bonds.
Questions 9-10 The graph shows the temperature of a pure
solid substance as it is heated at a constant rate to a gas.
9.
2.
3.
The increasing boiling points for HCl, HBr and HI can be
best explained because of the increase in
(A) dispersion force
(B) dipole moment
(C) valence electrons
(D) hydrogen bonding
A sample of an ideal gas is cooled from 50oC to 25oC in a
sealed container of constant volume. Which of the
following values for the gas will decrease?
I. The average kinetic energy of the molecules
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
Questions 4-7 refer to the phase diagram of a pure substance.
Pure liquid exists at time
(A) t1
(B) t2
(C) t3
(D) t4
10. Which of the following best describes what happens to the
substance between t4 and t5?
(A) The molecules are leaving the liquid phase.
(B) The solid and liquid phases coexist in equilibrium.
(C) The vapor pressure of the substance is decreasing.
(D) The average intermolecular distance is decreasing.
11. Gas in a closed rigid container is heated until its absolute
temperature is doubled, which is also doubled?
(A) The density of the gas
(B) The pressure of the gas
(C) The average speed of the gas molecules
(D) The number of molecules per liter
12. A 2.00-L of gas at 27oC is heated until its volume is 5.00 L.
If the pressure is constant, the final temperature is
(A) 68oC
(B) 120oC (C) 477oC (D) 677oC
4.
Which phase is most dense?
(A) solid
(B) liquid
(C) gas
(D) can't determine
5.
Which occurs when the temperature increases from 0°C to
40°C at a constant pressure of 0.5 atm?
(A) Sublimation
(B) Condensation
(C) Freezing
(D) Fusion
6.
Which occurs when the pressure increases from 0.5 to 1.5
atm at a constant temperature of 60°C?
(A) Sublimation
(B) Condensation
(C) Freezing
(D) Fusion
7.
The normal boiling point of the substance is closest to
(A) 20oC
(B) 40oC
(C) 70oC
(D) 100oC
8.
Which actions would be likely to change the boiling point of
a sample of a pure liquid in an open container?
(A) Placing it in a smaller container
(B) Increasing moles of the liquid in the container
(C) Moving the container to a higher altitude
(D) Increase the setting on the hot plate
13. Under the same conditions, which of the following gases
effuse at approximately half the rate of NH3?
(A) O2
(B) He
(C) CO2
(D) Cl2
14. What is the partial pressure (in atm) of N2 in a gaseous
mixture, which contains 7.0 moles N2, 2.5 moles O2, and
0.50 mole He at a total pressure of 0.90 atm.
(A) 0.13
(B) 0.27
(C) 0.63
(D) 0.90
Questions 15-16 refer to the following gases at 0°C and 1 atm.
(A) Ne
(B) Xe
(C) O2
(D) CO
15. Has an average atomic or molecular speed closest to that
of N2 molecules at 0°C and 1 atm
16. Has the greatest density
17. A 2-L container will hold about 4 g of which of the following
gases at 0oC and 1 atm?
(A) SO2
(B) N2
(C) CO2
(D) C4H8
18. As the temperature is raised from 20oC to 40oC, the
average kinetic energy of Ne atoms changes by a factor of
(A) ½
(B) (313/293)½
(C) 313/293
(D) 2
19. Which is the same for the structural isomers C2H5OH and
CH3OCH3? (Assume ideal behavior.)
(A) Gaseous densities at STP
(B) Vapor pressures at the same temperature
(C) Boiling points
(D) Melting points
20. The partial pressures of toluene is 22 torr and benzene is
75 torr in a mixture of these two gases. What is the mole
fraction of benzene in the gas mixture?
(A) 0.23
(B) 0.29
(C) 0.50
(D) 0.77
21. The system shown above is at equilibrium at 28°C. At this
temperature, the vapor pressure of water is 28 mm Hg.
The partial pressure (in mm Hg) of O2(g) in the system is
(A) 28
(B) 56
(C) 133
(D) 161
31. The mole fraction of ethanol in a 6 molal aqueous solution is
(A) 0.006
(B) 0.1
(C) 0.08
(D) 0.2
32. What additional information is needed to determine the
molality of a 1.0-M glucose (C6H12O6) solution?
(A) Volume
(B) Temperature
(C) Solubility of glucose (D) Density of the solution
33. The mole fraction of toluene (MM = 90) in a benzene (MM= 80)
solution is 0.2. What is the molality of the solution?
(A) 0.2
(B) 0.5
(C) 2
(D) 3
34. The volume of distilled water that is added to 10 mL of 6.0 M
HCI in order to prepare a 0.50 M HCI solution is
(A) 50 mL (B) 60 mL (C) 100 mL (D) 110 mL
35. A student wishes to prepare 2.00 L of 0.100 M KIO3
(MM = 214 g). The proper procedure is to weigh out
(A) 42.8 g of KIO3 and add 2.00 kg of H2O
(B) 42.8 g of KIO3 and add H2O to a final volume of 2.00 L
(C) 21.4 g of KIO3 and add H2O to a final volume of 2.00 L
(D) 42.8 g of KIO3 and add 2.00 L of H2O
22. In which of the 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)
36. What volume of 12 M HCl is diluted to obtain 1.0 L of 3.0-M?
(A) 4.0 mL
(B) 40 mL (C) 250 mL (D) 1,000 mL
23. Of the following compounds, which is the most ionic?
(A) SiCl4
(B) BrCl
(C) PCl3
(D) CaCl2
37. 400 mL of distilled water is added to 200 mL of 0.6 M MgCI2,
what is the resulting concentration of Mg2+?
(A) 0.2 M (B) 0.3 M (C) 0.4 M (D) 0.6 M
24. Which of the following oxides is a gas at 25°C and 1 atm?
(A) Rb2O
(B) N2O
(C) Na2O2 (D) SiO2
25. Which of the following has the highest melting point?
(A) S8
(B) I2
(C) SiO2
(D) SO2
26. Under which conditions is O2(g) the most soluble in H2O?
(A) 5.0 atm, 80oC
(B) 5.0 atm, 20oC
(C) 1.0 atm, 80oC
(D) 1.0 atm, 20oC
27. Which is lower for a solution of a volatile solute compared
to the pure solvent?
(A) normal boiling point (B) vapor pressure
(C) normal freezing point (D) osmotic pressure
38. When 70. mL of 3.0 M Na2CO3 is added to 30. mL of 1.0 M
NaHCO3 the resulting concentration of Na+ is
(A) 2.0 M
(B) 2.4 M (C) 4.0 M d. 4.5 M
39. The mass of H2SO4 (MM = 98 g) in 50 mL of 6.0-M solution
(A) 3.10 g
(B) 29.4 g (C) 300. g (D) 12.0 g
40. What mass of CuSO4• 5 H2O (MM = 250 g) is required to
prepare 250 mL of a 0.10 M solution?
(A) 4.0 g
(B) 6.3 g
(C) 34 g
(D) 85 g
41. What is the concentration of OH- in a mixture that contains
40. mL of 0.25 M KOH and 60. mL of 0.15 M Ba(OH)2?
(A) 0.10 M (B) 0.19 M (C) 0.28 M (D) 0.40 M
Questions 28-30 Refer to 0.20 M solutions of the following salts.
(A) NaBr
(B) KI
(C) MgCl2 (D) C6H12O6
28. Has the lowest freezing point
Practice Free Response
1.
a.
3.327 g of an unknown gas occupies 1.00-L at 25oC and
103 kPa. What is the molar mass of the gas?
b.
Which noble gas would have twice the effusion rate?
29. Has the lowest conductivity
30. Has the lowest boiling point
2.
3.
4.
5.
N2 with V = 200. mL, P = 99.7 kPa, and T = 27.0oC is mixed
with O2 and transferred to a 750.-mL container at 27.0oC.
The total pressure of the mixture is 90.4 kPa, at 27.0oC.
a. Calculate the moles of N2.
b.
Calculate the total moles of gas.
c.
Calculate the partial pressure of each gas.
Water is added to a 10.0-L container filled with dry air at
20oC. The container is sealed and equilibrium is established.
Would the amount of water vapor increase (), remain the
same (=) or decrease () for the following changes?
=


Use a 5.0 L container
Use humid air
Raise the temperature to 25oC
Add 20.0 g of water
Consider the following solids.
a. Rank the solids from highest melting point (1) to lowest.
CH4
H2O
MgO
Na
NaCl
SiO2
b.
Justify your relative ranking of CH4 and H2O.
c.
Justify your relative ranking of MgO and NaCl.
Explain the following observations.
a. NH3 boils at 240 K, whereas NF3 boils at 144 K.
b.
6.
At 25°C and 1 atm, F2 is a gas, whereas I2 is a solid.
Hydrogen gas is produced when aluminum foil is added to
a solution of hydrochloric acid.
a. The hydrogen is collected over water at 25oC and a
total pressure of 756 torr. What is the mole fraction of
H2(g) in the wet gas? (PH2O) at 25oC = 23.8 torr.
b.
If 255 mL of wet gas is collected, what is the yield of
hydrogen in grams?
c.
What is the density of the wet gas?
8.
What is the solubility of CO2 in an opened soft drink at
25oC where the partial pressure of CO2 is 3.0 x 10-4 atm?
(kCO2 = 3.1 x 10-2 mol/L•atm).
9.
Explain the following observations. Your responses must
include specific information about all substances.
a. When table salt (NaCI) and sugar (C12H22O11) are
dissolved in water, it is observed that
(1) both solutions have higher boiling points than
pure water
(2) the boiling point of 0.10 M NaCI(aq) is higher than
that of 0.10 M C12H22O11(aq).
b.
10. A student is instructed to prepare 100.0 mL of 1.250 M
NaOH from a stock solution of 5.000 M NaOH. The student
follows the proper safety guidelines.
a. Calculate the volume of 5.000 M NaOH needed to
accurately prepare 100.0 mL of 1.250 M NaOH.
b.
When CaCl2 is added to water, the temperature of the
solution decreases.
a. Justify which bond is stronger; hydration bonds
between ions and water or ionic bonds between ions?
b.
Justify why you expect CaCl2 to be more or less
soluble in warm water compared to cold water?
Describe the steps in a procedure to prepare 100.0
mL of 1.250 M NaOH solution using 5.000 M NaOH.
11. Consider camphor, C10H16O, a substance obtained from the
Formosa camphor tree. It has considerable use in the
polymer and drug industry. A solution of camphor is prepared
by mixing 30.0 g of camphor with 1.25 L of ethanol, C2H5OH
(d = 0.789 g/mL). Assume no change in volume when the
solution is prepared.
a. What is the mass percent of camphor in the solution?
b.
What is the molarity of the solution?
c.
What is the molality of the solution?
d.
The vapor pressure of pure ethanol at 25oC is 59.0
mm Hg. What is the vapor pressure of ethanol in the
solution at this temperature?
e.
What is the osmotic pressure of the solution at 25oC?

f.
7.
Ammonia, NH3, is very soluble in water, whereas
phosphine, PH3, is only moderately soluble in water.
What is the boiling point of the solution? The normal
boiling point of ethanol is 78.26oC (Kb = 1.22oC/m).

g.


The molar mass of cortisone acetate is determined by
dissolving 2.50 g in 50.0 g camphor (Kf = 40.0oC/m).
The freezing point of the mixture is 173.44oC; that of
pure camphor is 178.40oC. What is the molar mass of
cortisone acetate?