Chemistry 4: States of Matter Name __________________________A. Gas State (10.2 to 10.9)

1. tend to be small, non-polar molecules2. form homogenous mixtures3. 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)urms

2 = 3/2RT a. R = 8.31 J/mol•K b. MM (molar mass) in kgc. T in kelvin (TK = ToC + 273)d. average speed, urms = (3RT/MM)½

2. Graham’s lawa. rate r of effusion (leaking) or diffusion

(spreading out) is proportional to speedb. 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 lawsa. Ideal gas Law equation: PV = nRT

1. molecules generate pressure via collisionsa. pressure = force/areab. 1 atm = 101 kPa = 760 mm Hg (torr)c. measuring tools

1. barometer: atmospheric pressure2. 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 equationsa. P1V1/T1 = P2V2/T2

b. MM = mRT/PV = dRT/Pb. 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) = nRTb. "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 polarityc. "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

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 SolidEnergy Highest Middle Lowest

Disorder Greatest Middle LeastOccupied space Whole Bottom OwnCompressibility Yes No No2. cohesive forces (van der Waals forces)

a. attraction between molecules (covalent bonds hold atoms together in molecule)

b. dipole-dipole forces1. polar molecules 2. of one molecule attracts of a neighbor3. strength to polarity, if all else is even

c. London dispersion forces1. 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 bonding1. 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 temperaturesc. low vapor pressure (low volatility)

3. cooling profile for water from 110oC to -10oCA

100oC

0oC

gasB Ccondensingboiling

liquid

D E freezingmelting

solid F

Q (Heat Removed in J)a. slope C-D < slope E-F more heat is

removed when one mole H2O(l) is cooled 1oC compared to one mole H2O(s)

b. 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)

c. calculations C D

1. Q = nClT = mclT2. molar heat capacity, Cl = 75.3 J/mol•K(3. specific heat, cl = 4.18 J/g•K

4. phase diagram

a. point A: triple point (three phases at equilibrium) 1. above triple point: melting and vaporization2. 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)

5. vapora. 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. cooling process (hottest evaporate first, leaving cooler molecules behind)

d. equilibrium between liquid and vapor1. evaporation rate = condensation rate in a

closed container2. concentration of vapor measured as Pvap

3. independent of container size until no liquid4. Pvap increases at higher temperature because

a. more molecules are in vapor phase b. vapor molecules exert greater pressure

e. boiling occurs when Pvap = Patm boiling point decreases with elevation (lower air pressure, Patm)

f. 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 energy3. intermolecular forces (attraction between + and -) or

bonds (covalent, ionic or metallic) hold particles together 4. 4 types of solids

a. metallic—metals only1. attraction between cations and delocalized

valence electrons (electron sea model)

2. melting point: variable ( bond strength)3. conductivity: free electrons yes4. malleable: non-directional bond yes5. water solubility: no molecular interactions no6. examples: Cu, Ag, etc.

b. covalent network—nonmetals w/o H or halogen1. atoms covalently bond throughout w/o size

limit (different than large molecule) 2. melting point: strong bonds high3. conductivity: no free electrons no4. malleability: bond highly directional brittle5. water solubility: no molecular interactions no6. three examples

a. diamond and graphite—C1. allotropes (2 forms in the same state)2. diamond: covalently bonded 3-d

structure—good abrasive3. graphite: covalent bonded planar

sheets linked by dispersion forcesa. separate easily—good lubricantb. electron flow—good conductivity

b. quartz—SiO2 1. 3-d structure similar to diamond2. softens when heated until liquid3. fast cooling = non-crystalline glass

c. molecular—nonmetals often with H and/or halogen1. attraction between + of one molecule with

– of another2. melting point: weak bonds low3. conductivity: no free electrons no4. malleability: non-directional yes

(H-bonding in H2O(s) is somewhat directional)5. water solubility ("like dissolves like") yes/no6. examples: H2O, C6H12O6, etc.

d. Ionic—metal plus nonmetals1. attraction between cations and anions2. melting point: strong bonds high3. conductivity: no free electrons no

(fused or dissolved state is conducting)4. malleability: bond highly directional brittle5. water solubility: ion-dipole interaction yes6. 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 solution1. undissolved solid dissolved solid2. dissolution rate = crystallization rate3. 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 solubility1. solvent kinetic energy is used to break solute-

solute 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 solution3. only Pg, not Ptot, will increase solubility

2. separation solute and solventa. filtration: separate solvent from insoluble solute b. distillation

1. simple: separate solvent from soluble solid 2. fractional: separate solvent from soluble liquid

3. expressing concentrationa. 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)MMo 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 units1. 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/Vo 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

b. nonvolatile-nonelectrolyte: Pvap = XsolventPosolvent

1. known as Raoult's law2. "ideal" solution

c. two volatile liquids: Pvap = XAPoA + XBPo

B

2. higher boiling point and lower freezing pointa. lowered vapor pressure changes melting and

boiling temperatures (extends liquid phase)

b. Tb = Kbmi, Tf = Kfmi 1. m = molality2. 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

3. osmotic pressurea. semi-permeable membrane blocks soluteb. solvent flows from high [ ] to low [ ] osmosisc. osmotic pressure () = pressure to stop flow d. = MRTi

(R = 8.31 when in kPa or 0.0821 when in atm)

Experiments1. 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

(atm)T

(K)V

(L)m = m1 – m2

(g)

MM = mRT/PV MM (C4H10) %

c. How would the following affect the molar mass value?(1) The butane lighter was not thoroughly dried.

(2) The vapor pressure of water was not subtracted from the room pressure.

(3) The water level in the inverted graduated cylinder was higher than the level in the trough.

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

b. Determine the following.

mNaCl = m3 – m1

mH2O = m2 – m3

Volume of solution

c. Calculate the following.

Mole NaCl

Mole H2O

mass %

mole fraction

molarity

molality

3. 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

ToC 74.0 72.2 69.8 68.0 66.8 67.8 69.0 69.0Time (s) 160 180 200 220 240 260 280 300

ToC 68.8 68.8 69.0 68.8 69.0 67.0 65.2 63.8Part 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

ToC 70.0 68.0 65.8 64.2 61.8 60.0 58.8 60.2Time (s) 160 180 200 220 240 260 280 300

ToC 59.8 59.2 58.8 58.2 57.6 57.0 55.2 53.4Part 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

ToC 72.0 70.0 68.2 65.8 64.0 62.8 64.0 63.4Time (s) 160 180 200 220 240 260 280 300

ToC 63.0 62.8 62.2 61.8 61.4 61.0 59.0 57.2a. 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

70

68

66

64

62

60

58

56

5420 60 100 140 180 220 260 300

Time (s)b. Complete the following chart.

Naphthalene/BHT MixtureTf (BHT) Tf (Solution) Tf

n = m/MM

Molality m

Kf = Tf/m

Unknown/BHT MixtureTf (BHT) Tf (Solution) Tf

m = Tf/Kf

n = mBHTm

MM = m/n

c. The unknown's molecular formula is CH3(CH2)14CH2OH. Determine the percent error for this experiment.

Practice ProblemsA. Gas State

1. What features of the kinetic theory of gasesa. describe all gas molecules?

b. describe ideal gas molecules only?

2. Consider one mole of Ne gas at 274 K. Determinea. the total kinetic energy.

b. the average speed.

3. a. What alkane effuses at 1/5 the rate of He?

b. How many times faster does C2H2 diffuse compared to the alkane?

4. Consider the graph below.

a. A and B are He and O2, at 25oC, which is which? Explain

b. A and B are at 100 K and 200K, which is which? Explain

5. Determine the pressure of 1.22 atm in the following units.mm Hg kPa torr

6. If the atmospheric pressure is 749 mm Hg, what is the pressure of the enclosed gas in each case below?

7. A gas is confined inside a container with a movable piston held down by a fixed pressure.

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.

8. Complete the following table for an ideal gas:P V n T

2.00 atm 1.00 L 1.500 mol30.3 kPa 1.250 L 27oC650 torr 0.333 mol 350 K

585 mL 0.250 mol 295 K9. 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. 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 Pressure4.93 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 IIIThe samples are at the same temperature. Rank them with respect to the following (1 is highest).

I II IIITotal PressurePartial Pressure of HeDensityAverage Kinetic energy per moleculeTotal Kinetic energy

14. Each bulb contains a gas at the pressure and volume shown and temperature of 25oC. Determine

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.

N2

Ne

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 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?

b. Why do gases at temperatures near their boiling point deviate from ideal behavior?

B. Phase Change19. Which letter illustrates the types of molecular forces?

Dipole-Dipole Dispersion H-bond

20. For each pair, highlight the molecule with the higher boiling point and then justify your choice.

Pair JustificationH2O & H2S

Ne & KrCl2 & SO2

21. Explain the boiling points for the two isomers.

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?

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.

Temperature (oC)Can liquid CO2 exist at room pressure?What happens to CO2(s) at -78.5oC?Which is the most dense phase for CO2?What is the triple point pressure for CO2?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.

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. What percentage of the water evaporates?

30. Explain why water droplets form on a cold water bottle.

C. Crystalline Solids31. Use the electron-sea model of metals to explain

a. malleability.

b. 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.

Metallic Covalent Network Molecular Ionic

Structural UnitBond name

Bond strengthMelting point

SolubilityConductivityMalleabilityExample

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.

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. Solubility37. 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

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.

Solute mass solute mass water Temperature90 g 100 g 50oC

K2Cr2O7 100 g 90oCNaCl 70 g 30oCKClO3 15 g 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

mole NaOH

mole H2O

mass %

mole fraction

molarity

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?

b. What is the mass of one liter of solution?

c. What is the density of 12 M HCl in g/mL?

45. How would you prepare 250. mL of a 0.127 M Ca(OH)2

a. from powder Ca(OH)2?

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. 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. How many grams of Na2SO4 are in 0.800 L of solution?

c. What volume contains 157 g of Na2SO4?

50. Name the separation technique for the following.Separate salt from waterSeparate sand from waterSeparate alcohol from water

E. Colligative Properties51. 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?

b. What is the vapor pressure of the solution?

52. What is the vapor pressure when PH2O = 2.4 kPa, whena. 0.50 mol C6H12O6 is in 5.5 mol H2O.

b. 0.50 mol C2H5OH (PC2H5OH = 9 kPa) is in 5.5 mol H2O.

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 molality of the solution?

c. What is the change in freezing point of the solution?

d. What is the freezing point of the solution?

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.

b. the moles of ethylene glycol are in the solution.

c. the molecular mass of ethylene glycol.

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 25oC.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 25oC?

58. How do the colligative properties change (, ) when non-volatile solute is added to solvent?Vapor P. Freezing pt. Boiling pt. Osmotic P.

Practice Multiple ChoiceBriefly explain why the answer is correct in the space provided.Questions 1-2 The molecules have the normal boiling points.

Molecule HF HCl HBr HIBoiling 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.

2. 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

3. 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.

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

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. 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

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

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)

23. Of the following compounds, which is the most ionic? (A) SiCl4 (B) BrCl (C) PCl3 (D) CaCl2

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

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

29. Has the lowest conductivity

30. Has the lowest boiling point

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

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

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

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

Practice Free Response1. 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?

2. 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.

3. 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 airRaise the temperature to 25oCAdd 20.0 g of water

4. 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.

5. Explain the following observations.a. NH3 boils at 240 K, whereas NF3 boils at 144 K.

b. At 25°C and 1 atm, F2 is a gas, whereas I2 is a solid.

6. 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?

7. 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?

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. Ammonia, NH3, is very soluble in water, whereas phosphine, PH3, is only moderately soluble in water.

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. 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. 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?