AP Chemistry Unit 5 States of Matter Part A: Phase Changes - - PDF document

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Slide 1 / 43 New Jersey Center for Teaching and Learning Progressive Science Initiative This material is made freely available at www.njctl.org and is intended for the non-commercial use of students and teachers. These materials may not be

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Unit 5 States of Matter Part A: Phase Changes

AP Chemistry Slide 2 / 43 States of Matter

Many scientists now believe that the temperature becomes warm enough in Mars equatorial regions to melt subsurface ice and cause small gullies to form in the soil.

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SLIDE 2

Phase Changes

Phase changes involve a change in the potential energy of a material brought about by either the strengthening or weakening

f the interactions between representative particles.

The table below lists the six phase changes of water and the potential energy change of each. Phase Change States Involved Change in PE Change in Tsurr Endo or Exothermic Melting s --> l +6.0 kJ/mol

Endo

Evaporation l --> g +41 kJ/mol

Endo

Condensation g --> l

41 kJ/mol

+ Exo Freezing l --> s

6.0 kJ/mol

+ Exo Sublimation s --> g +47 kJ/mol

Endo

Deposition g --> s

47 kJ/mol

+ Exo

Slide 4 / 43

Phase Changes

Endothermic phase changes occur when the molecules have enough kinetic energy to overcome the coulombic attractions between the particles. Examining the melting of water (ice)

Potential Energy solid liquid

Since the kinetic energy of the molecules was converted to potential energy, the T

surr will decrease.

When a substance melts, the bonds are only stretched thus requiring less energy than boiling where the bonds are broken.

Slide 5 / 43 Phase Changes

Exothermic phase changes occur when the molecules have lost enough kinetic energy to permit coulombic attractions to form between the particles. Examining the condensing of water

Potential Energy liquid gas

Since the potential energy of the molecules was converted to kinetic energy, the T

surr will increase.

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SLIDE 3

Phase Changes

At a certain temperature and pressure, two phase changes will

ccur at equal rates resulting in equilibrium.

Standard Melting Point of H2O(0 C at 1 atm) H2O(s) <--> H2O(l) = equilibrium Temperature above 0 C at 1 atm H2O(s) --> H2O(l) = not at equilibrium Temperature below 0 C at 1 atm H2O(s) <-- H2O(l) = not at equilibrium

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Phase Changes

The standard (1 atm) melting and boiling point temperatures are dependent on the types of interactions found within the substance. Substance Example Particle Interactions Influenced by Standard MP Ionic MgO Ionic (intra) charge and ionic radii 2852 C Covalent Network C(diamond) Covalent (intra) atomic radii 3550 C Metallic Au Covalent (intra) ionic radii 1064 C Molecular I2 LDF's (inter) polarizability (# of e-) 114 C Note the big difference in melting point between substances with intra vs. inter molecular interactions.

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1 Which of the following phase changes would be exothermic? A Na(s) --> Na(l) B Na(l) --> Na(g) C Na(s) --> Na(g) D Na(g) --> Na(l)

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SLIDE 4

1 Which of the following phase changes would be exothermic? A Na(s) --> Na(l) B Na(l) --> Na(g) C Na(s) --> Na(g) D Na(g) --> Na(l)

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Answer D

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2 Which of the following would be true when ethyl alcohol evaporates? A I only B II only C III only D I and III only E I, II, and III

I. The potential energy of the

substance increases.

II. The kinetic energy of the

surroundings increases.

III. Coulombic attractions are

broken.

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2 Which of the following would be true when ethyl alcohol evaporates? A I only B II only C III only D I and III only E I, II, and III

I. The potential energy of the

substance increases.

II. The kinetic energy of the

surroundings increases.

III. Coulombic attractions are

broken.

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Answer D

Slide 10 (Answer) / 43

SLIDE 5

3 Which of the following changes would be endothermic? A sublimation B deposition C condensation D freezing E All of these

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3 Which of the following changes would be endothermic? A sublimation B deposition C condensation D freezing E All of these

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Answer A

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4 For which of the following are covalent bonds breaking? A CO2(s) --> CO2(l) B CO2(l) --> CO2(g) C C(s) --> C(g) D C(s) --> C(l) E NaCl(l) --> NaCl(g)

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SLIDE 6

4 For which of the following are covalent bonds breaking? A CO2(s) --> CO2(l) B CO2(l) --> CO2(g) C C(s) --> C(g) D C(s) --> C(l) E NaCl(l) --> NaCl(g)

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Answer C

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5 Which of the following BEST characterize the standard boiling point of a material? A Liquid and vapor phases are in equilibrium at any pressure B Liquid and vapor phases are in equilibrium at1 atm pressure C The rate of vapor formation exceeds that of liquid formation at any pressure D The rate of liquid formation exceeds that of vapor formation at 1 atm E The potential energy of the material is decreasing

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5 Which of the following BEST characterize the standard boiling point of a material? A Liquid and vapor phases are in equilibrium at any pressure B Liquid and vapor phases are in equilibrium at1 atm pressure C The rate of vapor formation exceeds that of liquid formation at any pressure D The rate of liquid formation exceeds that of vapor formation at 1 atm E The potential energy of the material is decreasing

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Answer B

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SLIDE 7

6 Which of the following correctly ranks the materials from lowest to highest melting point? A Na < Li < K B CaO < MgO < NaF C CH4 < NH3 < H2O D I2 < Br2 < F2 E Na < H2O < I2

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6 Which of the following correctly ranks the materials from lowest to highest melting point? A Na < Li < K B CaO < MgO < NaF C CH4 < NH3 < H2O D I2 < Br2 < F2 E Na < H2O < I2

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Answer C

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Phase Changes

The heat of fusion is the amount of energy that is required to melt a given quantity of a material. The same amount of energy is released when that substance freezes. The heat of fusion depends on the strength of the particle interactions within the solid. Substance Substance Type Particle Interactions Heat of Fusion NaCl Ionic Ionic Bonds (Intra) 28 kJ/mol NaF Ionic Ionic Bonds (Intra) 32 kJ/mol H2O Molecular H-Bonds, LDF's (Inter) 6 kJ/mol CH4 Molecular LDF's (Inter) 1 kJ/mol Ag Metallic Covalent Bonds (Intra) 11.3 kJ/mol

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SLIDE 8

Phase Changes

The heat of vaporization is the amount of energy that is required to vaporize a given quantity of a material. The same amount of energy is released when that substance condenses. The heat of vaporization depends on the strength of the particle interactions within the solid. Substance Substance Type Particle Interactions Heat of Fusion CH3OH Molecular LDF's, H- Bonds 35 kJ/mol CH3CH2CH3 Molecular LDF's 16 kJ/mol H2O Molecular H-Bonds, LDF's 41 kJ/mol CH3COCH3 Molecular LDF's, DDF's 31 kJ/mol I2 Molecular LDF's 42 kJ/mol

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Phase Changes

Calculating the heats of fusion or vaporization requires measuring the energy absorbed or released as the material undergoes a phase change. Example: What is the heat of fusion for water (kJ/mole) assuming that the temperature of a 100 gram sample of water (Cp = 4.2 J/gC) decreased by 14.2 degrees when 18.0 grams of ice is melted in a container surrounded by the water.

1. Find the heat absorbed from water:

Heat absorbed (q) = m* T*cp = 100 g x 14.2 C x 4.2 J = 5,960 J gC

2. Find kJ/mol:

18 g H2O x 1 mol = 1 mole H2O 18 g 5.96 kJ/ 1mole = 5.96 kJ/mol

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7 Which of the following would have the highest heat of fusion? A Na B Li C K D Mg E Be

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SLIDE 9

7 Which of the following would have the highest heat of fusion? A Na B Li C K D Mg E Be

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Answer E

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8 Which of the following would have the lowest heat of vaporization? A C4H10 B CH3CH2CH2CHO C I2 D H2O E C2H4

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8 Which of the following would have the lowest heat of vaporization? A C4H10 B CH3CH2CH2CHO C I2 D H2O E C2H4

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Answer E

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SLIDE 10

9 What is the heat of vaporization (J/mol) of ethanol assuming that 419 Joules of energy are needed to vaporize 0.5 grams of the material?

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9 What is the heat of vaporization (J/mol) of ethanol assuming that 419 Joules of energy are needed to vaporize 0.5 grams of the material?

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Answer 38,090 J/mol

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10 How many grams of water must have been present if 3210 J of energy were needed to vaporize the sample? (The heat of vaporization of water is 41 kJ/mol)

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SLIDE 11

10 How many grams of water must have been present if 3210 J of energy were needed to vaporize the sample? (The heat of vaporization of water is 41 kJ/mol)

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Answer 14.09 grams

Slide 21 (Answer) / 43

11 Calculate the heat of fusion (kJ/mol) for benzene (C6H6) based on the following data. When a 12 gram sample of benzene was melted in in a 500 gram sample of water, the temperature of the water dropped by 0.73 degrees.

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11 Calculate the heat of fusion (kJ/mol) for benzene (C6H6) based on the following data. When a 12 gram sample of benzene was melted in in a 500 gram sample of water, the temperature of the water dropped by 0.73 degrees.

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Answer 9.9 kJ/mol

Slide 22 (Answer) / 43

SLIDE 12

Heating/Cooling Curves

A heating curve plots the change in temperature over time as a material is heated through it's phase changes.

time s s <--> l l l <--> g g

Melting/Freezing Point Boiling Point

More time/heat is needed to vaporize the material than melt

it. This is due to the heat of

vaporization being much higher than the heat of fusion.

Note how the temperature (average kinetic energy) remains stable during a phase change. This is because kinetic energy is being converted to potential energy when a substance changes state.

Slide 23 / 43 Heating/Cooling Curves

A heating curve provides information about the strength of particle interactions.

time s s <--> l l l <--> g g

Melting/Freezing Point Boiling Point

Assuming a constant amount of added heat, the shallower the slope, the higher the specific heat, indicating stronger particle interactions. H2O Cp = 4.2 J/gC

time s s <--> l l l <--> g g

Melting/Freezing Point Boiling Point

CH3OH Cp = 1.7 J/gC

Slide 24 / 43 Heating/Cooling Curves

A cooling curve plots the change in temperature over time as a material is cooled through it's phase changes. . time cooled Temperature (C) boiling point 100 freezing point KE KE KE PE

PE

super-cooled liquid

H2O Cooling Curve

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SLIDE 13

12 At which point(s) would the solid and liquid phase be in equilibrium? A A B A & B C C D C & D E E

time

A B C D E

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12 At which point(s) would the solid and liquid phase be in equilibrium? A A B A & B C C D C & D E E

time

A B C D E

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Answer B

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13 At which point would kinetic energy be transferred to potential energy? A A only B A & C & E C B & D D A & B E E only

time

A B C D E

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SLIDE 14

13 At which point would kinetic energy be transferred to potential energy? A A only B A & C & E C B & D D A & B E E only

time

A B C D E

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Answer C

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14 What best characterizes what is occurring at point D

n the heating curve?

A The distance between the molecules is decreasing B The kinetic energy of the molecules is increasing C The potential energy of the molecules is decreasing D Coulombic attractions are being broken E None of these

time

A B C D E

Slide 28 / 43

14 What best characterizes what is occurring at point D

n the heating curve?

A The distance between the molecules is decreasing B The kinetic energy of the molecules is increasing C The potential energy of the molecules is decreasing D Coulombic attractions are being broken E None of these

time

A B C D E

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Answer D

Slide 28 (Answer) / 43

SLIDE 15

15 The cooling curve for an unknown substance is found below. What is its freezing point? A 94 C B 100 C C 80 C D 78 C E 70 C

T time 100 90 80 70

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15 The cooling curve for an unknown substance is found below. What is its freezing point? A 94 C B 100 C C 80 C D 78 C E 70 C

T time 100 90 80 70

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Answer C

Slide 29 (Answer) / 43

16 How much energy (in Joules) would be needed to heat of 120 gram chunk of ice from -7 C to 112 C assuming the following:

Heat of Fusion = 6 kJ/mol Heat of vaporization = 41 kJ/mol Cp (ice) = 2.1 J/gC Cp(water) = 4.2 J/gC Cp(steam) = 2.0 J/gC

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SLIDE 16

16 How much energy (in Joules) would be needed to heat of 120 gram chunk of ice from -7 C to 112 C assuming the following:

Heat of Fusion = 6 kJ/mol Heat of vaporization = 41 kJ/mol Cp (ice) = 2.1 J/gC Cp(water) = 4.2 J/gC Cp(steam) = 2.0 J/gC

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Answer 368,000 J or 368 kJ

Slide 30 (Answer) / 43 Phase Diagrams

Phase Diagrams are not officially a part of the AP Chemistry curriculum as they expect the student to have mastered the content in a first year chemistry course. For those students who did not master the material yet for whatever reason, the slides from the first year course covering phase diagrams have been included directly following this one.

Slide 31 / 43 Phase Diagrams

Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases. In general, increasing the pressure strengthens bonds and tends to produce the solid state. In contrast, increasing the temperature will weaken bonds and will promote the gaseous state.

Slide 32 / 43

SLIDE 17

Phase Diagrams

The liquid-vapor interface ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

Slide 33 / 43 Phase Diagrams

The circled line in the diagram below is the interface between liquid and solid. The melting point at each pressure can be found along this line.

Slide 34 / 43 Phase Diagrams

Below the triple point the substance cannot exist in the liquid state. Along the circled line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line. Sublimation and deposition only occur below the triple point.

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SLIDE 18

The circled line is the liquid-vapor interface. It starts at the triple point (T), the point at which all three states are in equilibrium.

Phase Diagrams Slide 36 / 43 Phase Diagram of Water

Note the solid-liquid equilibrium line has a negative slope. This means that an increase in pressure causes the liquid state to form. Due to the hydrogen bonding between water molecules, the solid state of water is less dense than the liquid so an increase in pressure will disrupt these hydrogen bonds and produce the liquid.

Slide 37 / 43 Comparison of Two Phase Diagrams

For carbon dioxide, the slope of the solid-liquid line is positive, as it is for most

ther substances. This

means that an increase in pressure can cause substances to freeze. For water, the slope of the solid-liquid line is negative. This means that an increase in pressure can cause this substance to melt.

Slide 38 / 43

SLIDE 19

Phase Diagram of Carbon Dioxide

Notice that CO

2 has a triple point at a pressure 5x normal

atmospheric pressure. As a result, when solid CO

2 is heated,

it will sublimate directly to a gas. This produces a nice foggy effect in theaters!

Slide 39 / 43

17 Which point represents a freezing point of the substance?

A B C D E P (atm) T (C)

20 0 40 60

1 2 3 A B C D E

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17 Which point represents a freezing point of the substance?

A B C D E P (atm) T (C)

20 0 40 60

1 2 3 A B C D E

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Answer C

Slide 40 (Answer) / 43

SLIDE 20

18 If the substance has its pressure increased isothermally at 40 C, from 1 atm to 3 atm, what phase change will it undergo?

A condensation B evaporation C sublimation D deposition E melting P (atm) T (C)

20 0 40 60

1 2 3 A B C D E

Slide 41 / 43

18 If the substance has its pressure increased isothermally at 40 C, from 1 atm to 3 atm, what phase change will it undergo?

A condensation B evaporation C sublimation D deposition E melting P (atm) T (C)

20 0 40 60

1 2 3 A B C D E

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Answer A

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19 In the phase diagram below, if the material is heated from -20 C to 60 C at a constant pressure of 1 atm, what phase change will occur?

A Evaporation B Condensation C Melting D Freezing E Sublimation P (atm) T (C)

20 0 40 60

1 2 3 A B C D E

Slide 42 / 43

SLIDE 21